The materials to prepare students for practical lessons of inorganic chemistry

The materials to prepare students for practical lessons of inorganic chemistry

LESSON № 14.

Themes. Chemical elements and their classification. Human and biosphere. s – elements. Alkali metals. Hydrogen and its compounds. Water.

 

 

Chemical elements

A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus. All chemical matter consists of these elements. Common examples of elements are iron, copper, silver, gold, hydrogen, carbon, nitrogen and oxygen.

Save for the hydrogen and helium in the universe, which are thought to have been mostly produced in the Big Bang, most chemical elements are thought to have been produced by later processes.

These processes are divided into:

·          cosmic ray spallation (important for lithium, beryllium and boron, though some of these may have formed in the Big Bang), and

·          stellar nucleosynthesis which produces all elements heavier than boron (with carbon being the first of this series). The very heaviest elements (those beyond element 94, plutonium) decay with half lives too short to allow them to be observed naturally on Earth.

In total, 118 elements have been characterized as of March 2010, and new elements of higher atomic number are “discovered” from time to time, as new synthetic products of artificial nuclear reactions. Of the known elements, the first 92 occur naturally on Earth. Of these, oxygen is the most abundant element in the Earth’s crust. About 80 elements have stable isotopes: namely all elements with atomic numbers 1 to 82, except elements 43 and 61 (technetium and promethium). About half of the 80 stable elements are expected to be radioactive with such long half lives that their decay happens only in theory, and is too slow to have yet been detected by experiment. These elements (such as bismuth, only recently measured as unstable) have half lives at least 100 million to 1000 million times longer than the estimated age of the universe.

Classification of elements

• periodic table compactly shows relationships between elements

• features of the periodic table

• Periods are horizontal rows on the table.

• Groups (or families) are columns on the table.

• elements in the same group are called congeners. They have similar chemical properties.

• Blocks are regions on the table.

• important groups:

• alkali metals (Group IA, first column )

• soft, extremely reactive metals

• react with cold water to form hydrogen gas

• form +1 ions

• alkaline earth metals (Group IIA, second column):

• soft, reactive metals

• compounds are a major component of earth’s crust

• form +2 ions

• halogens (Group VIIA, next-to-last column):

• poisonous and extremely reactive nonmetals

• fluorine and chlorine are yellow-green gases

• bromine is a volatile red-brown liquid

• iodine is a volatile blue black solid

• all form -1 ions

• noble gases (Group 0, last column)

• all are monatomic gases

• a. k. a. inert gases; almost completely unreactive

• Important blocks:

• transition metals are the elements in the region from the third to twelfth columns.

• hard, dense metals

• less reactive than Group IA and IIA

• rare earth metals are the elements in the annex at the bottom of the table.

• lanthanides (annex, top row)

• actinides (annex, bottom row)

• main group elements are all elements except the transition and rare earth metals.

• group numbers end with “A”

• metals, nonmetals, and metalloids (semimetals)

• metallic properties

• luster

• malleability: can be hammered into thin sheets

• ductility: can be drawn into wire

• conduct heat and electricity well

 

Separating mixtures

·        mixture’s components have different properties

·        devise a process that selects components with certain properties

·        density, melting point, boiling point, solubility, reactivity, magnetism, polarity

·        some basic techniques

·        filtration: select components by particle size

·        floatation: select components by density

·        crystallization: select components by solubility

·        extraction: select components by solubility

·        distillation: select components by boiling point

·        chromatography: select components by affinity for a ‘stationary phase’

 

Elements with atomic numbers 83 or higher (bismuth and above) are unstable to the point that their instability has been detected, and they undergo radioactive decay. The elements from atomic number 83 to 94 are composed entirely of radioactive isotopes. However, along with unstable elements 43 and 61, they are nevertheless found on Earth, though sometimes in very small amounts. Some of these elements, notably uranium and thorium, have one or more isotopes with half lives long enough to survive as remnants of the primordial explosive stellar nucleosynthesis that produced the heavy elements before formation of our solar system (see primordial nuclide). Other radioactive elements continue to be produced iatural processes, such as production by cosmic rays, or as shorter-lived daughter nuclides or transmutation products from natural decay of longer-lived radioactives.

The periodic table of the chemical elements

When two distinct elements are chemically combined, with the atoms held together by chemical bonds, the result is termed a compound. Chemical compounds may result in elements combined in exact whole number ratios of atoms (a familiar example is water). The term “compound” does not always imply an exact combination ratio, however, inasmuch as chemical bonding of many types of elements results in crystalline solids and metallic alloys, for which exact formulas do not exist. Most of the solid substance of the Earth is of this latter type: the atoms are present in the substance of the Earth’s crust, mantle, and core are combined into chemical compounds of many compositions, but these do not have precise empirical formulas.

History

Ancient philosophy posited a set of classical elements to explain patterns iature. Elements originally referred to earth, water, air and fire rather than the chemical elements of modern science.

The term ‘elements’ (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE, in his dialogue Timaeus, which includes a discussion of the composition of inorganic and organic bodies and is a speculative treatise on chemistry. Plato believed the elements introduced a century earlier by Empedocles were composed of small polyhedral forms: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).

Aristotle, also used the term stoicheia and added a fifth element called aether, which formed the heavens. Aristotle defined an element as:

Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.

In 1661, Robert Boyle showed that there were more than just four classical elements as the ancients had assumed. The first modern list of chemical elements was given in Antoine Lavoisier’s 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric. By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that cannot be decomposed into any simpler substance. Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. In 1913, Henry Moseley discovered that the physical basis of the atomic number of the atom was its nuclear charge, which eventually led to the current definition. The current definition also avoids some ambiguities due to isotopes and allotropes.

By 1919, there were seventy-two known elements. In 1955, element 101 was discovered and named mendelevium in honor of Mendeleev, the first to arrange the elements in a periodic manner. In October 2006, the synthesis of element 118 was reported; the synthesis of element 117 was reported in April 2010.

Description

The lightest elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe in a ratio of around 3:1 by mass (approximately 12:1 by number of atoms). Almost all other elements found iature, including some further hydrogen and helium created since then, were made by various natural or (at times) artificial methods of nucleosynthesis, including occasionally breakdown activities such as nuclear fission, alpha decay, cluster decay, and cosmic ray spallation.

As of 2010, there are 118 known elements (in this context, “known” means observed well enough, even from just a few decay products, to have been differentiated from any other element). Of these 118 elements, 94 occur naturally on Earth. Six of these occur in extreme trace quantities: technetium, atomic number 43; promethium, number 61; astatine, number 85; francium, number 87; neptunium, number 93; and plutonium, number 94. These 94 elements, and also possibly element 98 californium, have been detected in the universe at large, in the spectra of stars and also supernovae, where short-lived radioactive elements are newly being made. The first 94 elements have been detected directly on Earth as naturally-occurring fission or transmutation products of uranium and thorium. Some californium may be present on Earth, but at present, natural californium is only known from supernovae spectra.

The remaining 24 elements, not found on Earth or in astronomical spectra, have been derived artificially. All of the elements that are derived solely through artificial means are radioactive with very short half-lives; if any atoms of these elements were present at the formation of Earth, they are extremely likely to have already decayed, and if present iovae, have been in quantities too small to have beeoted. Technetium was the first purportedly non-naturally occurring element to be synthesized, in 1937, although trace amounts of technetium have since been found iature, and the element may have been discovered naturally in 1925. This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally occurring trace elements.

Lists of the elements are available by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as Ionization energies of the elements. The nuclides of stable and radioactive elements are also available as a list of nuclides, sorted by length of half life for those that are unstable. One of the most convenient, and certainly the most traditional presentation of the elements, is in form of periodic table, which groups elements with similar chemical properties (and usually also similar electronic structures) together.

Atomic number

The atomic number of an element, Z, is equal to the number of protons that defines the element. For example, all carbon atoms contain 6 protons in their nucleus; so the atomic number “Z” of carbon is 6. Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known as isotopes of the element.

The number of protons in the atomic nucleus also determines its electric charge, which in turn determines the electrons of the atom in its non-ionized state. This in turn (by means of the Pauli exclusion principle) determines the atom’s various chemical properties. So all carbon atoms, for example, ultimately have identical chemical properties because they all have the same number of protons in their nucleus, and therefore have the same atomic number. It is for this reason that atomic number rather than mass number (or atomic weight) is considered the identifying characteristic of an element.

Atomic nuclei consist of protons and neutrons (general title – nucleons) (Fig.1). The number of protons (Z) in atomic nuclei is strictly defined and equal to the serial number of element in the Periodic system. The number of neutrons in the atomic nuclei of one and the same element can be different – A-Z (where A ‑ relative atomic mass of element; Z ‑serial number) (Fig. 2)

Fig. 1. Structure of atomic nuclei

 

Fig. 2. Mass number of helium atomic nuclei

The number of protons defines nucleus charge of atom.

Nucleus mass is defined by the sum of protons and neutrons.

 

Atomic mass

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a super-index on the left hand side of the atomic symbol (e.g., ²³⁸U).

The relative atomic mass of an element is the average of the atomic masses of all the chemical element’s isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit (u). This number may be a fraction that is not close to a whole number, due to the averaging process. On the other hand, the atomic mass of a pure isotope is quite close to its mass number. Whereas the mass number is a natural (or whole) number, the atomic mass of a single isotope is a real number that is close to a natural number. In general, it differs slightly from the mass number as the mass of the protons and neutrons is not exactly 1 u, the electrons also contribute slightly to the atomic mass, and because of the nuclear binding energy. For example, the mass of 19F is 18.9984032 u. The only exception to the atomic mass of an isotope not being a natural number is 12C, which has a mass of exactly 12, because u is defined as 1/12th of the mass of a free carbon-12 atom.

Isotopes

Isotopes are atoms of the same element (that is, with the same number of protons in their atomic nucleus), but having different numbers of neutrons. Most (66 of 94) naturally occurring elements have more than one stable isotope. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons in the nucleus, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, the three isotopes of carbon are known as carbon-12, carbon-13, and carbon-14, often abbreviated to ¹²C, ¹³C, and ¹⁴C. Carbon in everyday life and in chemistry is a mixture of ¹²C, ¹³C, and ¹⁴C atoms.

Except in the case of the isotopes of hydrogen (which differ greatly from each other in relative mass—enough to cause chemical effects), the isotopes of the various elements are typically chemically nearly indistinguishable from each other. For example, the three naturally occurring isotopes of carbon have essentially the same chemical properties, but different nuclear properties. In this example, carbon-12 and carbon-13 are stable atoms, but carbon-14 is unstable; it is radioactive, undergoing beta decay into nitrogen-14.

As illustrated by carbon, all of the elements have some isotopes that are radioactive (radioisotopes), which decay into other elements upon radiating an alpha or beta particle. Certain elements only have radioactive isotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one stable isotope, and the meaumber of stable isotopes for the 80 stable elements is 3.1 stable isotopes per element. The largest number of stable isotopes that occur for an element is 10 (for tin, element 50).

Radioactive, primordial, and stable isotopes

Some isotopes are radioactive, and are therefore described as radioisotopes or radionuclides, while others have never been observed to undergo radioactive decay and are described as stable isotopes or stable nuclides. For example, 14C is a radioactive form of carbon while 12C and 13C are stable isotopes. There are about 339 naturally occurring nuclides on Earth, of which 288 are primordial nuclides, meaning that they have existed since the solar system’s formation.

Primordial nuclides include 35 nuclides with very long half-lives (over 80 million years) and 254 that are formally considered as “stable nuclides”, since they have not been observed to decay. In most cases, for obvious reasons, if an element has stable isotopes, those isotopes predominate in the elemental abundance found on Earth and in the solar system. However, in the cases of three elements (tellurium, indium, and rhenium) the most abundant isotope found iature is actually one (or two) extremely long lived radioisotope(s) of the element, despite these elements having one or more stable isotopes.

Many apparently “stable” isotopes/nuclides are predicted by theory to be radioactive, with extremely long half-lives (this does not count the possibility of proton decay, which would make all nuclides ultimately unstable). Of the 254 nuclides never observed to decay, only 90 of these (all from the first 40 elements) are stable in theory to all known forms of decay. Element 41 (niobium) is theoretically unstable via spontaneous fission, but this has never been detected. Many other stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay products have yet been observed, and so these isotopes are described as “observationally stable”. The predicted half-lives for these nuclides often greatly exceed the estimated age of the universe, and in fact there are also 27 known radionuclides (see primordial nuclide) with half-lives longer than the age of the universe.

Adding in the radioactive nuclides that have been created artificially, there are more than 3100 currently knowuclides. These include 905 nuclides that are either stable or have half-lives longer than 60 minutes. See list of nuclides for details.

History

Radioactive isotopes

The existence of isotopes was first suggested in 1913 by the radiochemist Frederick Soddy, based on studies of radioactive decay chains that indicated about 40 different species described as radioelements (i.e. radioactive elements) between uranium and lead, although the periodic table only allowed for 11 elements from uranium to lead.

Several attempts to separate these new radioelements chemically had failed. For example, Soddy had shown in 1910 that mesothorium (later shown to be 228Ra), radium (226Ra, the longest-lived isotope), and thorium X (224Ra) are impossible to separate. Attempts to place the radioelements in the periodic table led Soddy and Kazimierz Fajans independently to propose their radioactive displacement law in 1913, to the effect that alpha decay produced an element two places to the left in the periodic table, while beta decay emission produced an element one place to the right. Soddy recognized that emission of an alpha particle followed by two beta particles led to the formation of an element chemically identical to the initial element but with a mass four units lighter and with different radioactive properties.

Soddy proposed that several types of atoms (differing in radioactive properties) could occupy the same place in the table. For example, the alpha-decay of uranium-235 forms thorium-231, while the beta decay of actinium-230 forms thorium-230 The term “isotope”, Greek for “at the same place”, was suggested to Soddy by Margaret Todd, a Scottish physician and family friend, during a conversation in which he explained his ideas to her.

In 1914 T. W. Richards found variations between the atomic weight of lead from different mineral sources, attributable to variations in isotopic composition due to different radioactive origins.

Stable isotopes

The first evidence for multiple isotopes of a stable (non-radioactive) element was found by J. J. Thomson in 1913 as part of his exploration into the composition of canal rays (positive ions). Thomson channeled streams of neon ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path. Each stream created a glowing patch on the plate at the point it struck. Thomson observed two separate patches of light on the photographic plate (see image), which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest.

F. W. Aston subsequently discovered multiple stable isotopes for numerous elements using a mass spectrograph. In 1919 Aston studied neon with sufficient resolution to show that the two isotopic masses are very close to the integers 20 and 22, and that neither is equal to the known molar mass (20.2) of neon gas. This is an example of Aston’s whole number rule for isotopic masses, which states that large deviations of elemental molar masses from integers are primarily due to the fact that the element is a mixture of isotopes. Aston similarly showed that the molar mass of chlorine (35.45) is a weighted average of the almost integral masses for the two isotopes Cl-35 and Cl-37.

Chemical and molecular properties

A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and share a similar electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium (1H) and deuterium (2H), because deuterium has twice the mass of protium. The mass effect between deuterium and the relatively light protium also affects the behavior of their respective chemical bonds, by means of changing the center of gravity (reduced mass) of the atomic systems. However, for heavier elements, which have more neutrons than lighter elements, the ratio of the nuclear mass to the collective electronic mass is far greater, and the relative mass difference between isotopes is much less. For these two reasons, the mass-difference effects on chemistry are usually negligible.

In similar manner, two molecules that differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. As a consequence, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.

Nuclear properties and stability

Atomic nuclei consist of protons and neutrons bound together by the residual strong force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their copresence pushes protons slightly apart, reducing the electrostatic repulsion between the protons, and they exert the attractive nuclear force on each other and on protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, so does the ratio of neutrons to protons necessary to ensure a stable nucleus (see graph at right). For example, although the neutron:proton ratio of 3/2He is 1:2, the neutron:proton ratio of 238/92U is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 (Z = N). The nuclide 40/20Ca (calcium-40) is observationally the heaviest stable nuclide with the same number of neutrons and protons; (theoretically, the heaviest stable one is sulfur-32). All stable nuclides heavier than calcium-40 contain more neutrons than protons.

Numbers of isotopes per element

Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin). No element has nine stable isotopes. Xenon is the only element with eight stable isotopes. Four elements have seven stable isotopes, eight have six stable isotopes, ten have five stable isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have two stable isotopes (counting 180/73Ta as stable), and 26 elements have only a single stable isotope (of these, 19 are so-called mononuclidic elements, having a single primordial stable isotope that dominates and fixes the atomic weight of the natural element to high precision; 3 radioactive mononuclidic elements occur as well). In total, there are 254 nuclides that have not been observed to decay. For the 80 elements that have one or more stable isotopes, the average number of stable isotopes is 254/80 = 3.2 isotopes per element.

The proton:neutron ratio is not the only factor affecting nuclear stability. It depends also on evenness or oddness of its atomic number Z, neutroumber N and, consequently, of their sum, the mass number A. Oddness of both Z and N tends to lower the nuclear binding energy, making odd nuclei, generally, less stable. This remarkable difference of nuclear binding energy between neighbouring nuclei, especially of odd-A isobars, has important consequences: unstable isotopes with a nonoptimal number of neutrons or protons decay by beta decay (including positron decay), electron capture or other exotic means, such as spontaneous fission and cluster decay.

The majority of stable nuclides are even-proton-even-neutron, where all numbers Z, N, and A are even. The odd-A stable nuclides are divided (roughly evenly) into odd-proton-even-neutron, and even-proton-odd-neutrouclides. Odd-proton-odd-neutrouclei are the least common.

Even atomic number

The 148 even-proton, even-neutron (EE) nuclides comprise ~ 58% of all stable nuclides and all have spin 0 because of pairing. There are also 22 primordial long-lived even-even nuclides. As a result, each of the 41 even-numbered elements from 2 to 82 has at least one stable isotope, and most of these elements have several primordial isotopes. Half of these even-numbered elements have six or more stable isotopes. The extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five or eight nucleons from existing for long enough to serve as platforms for the buildup of heavier elements via nuclear fusion in stars (see triple alpha process).

These 53 stable nuclides have an eveumber of protons and an odd number of neutrons. They are a minority in comparison to the even-even isotopes, which are about 3 times as numerous. Among the 41 even-Z elements that have a stable nuclide, only three elements (argon, cerium, and lead) have no even-odd stable nuclides. One element (tin) has three. There are 24 elements that have one even-odd nuclide and 13 that have two odd-eveuclides. Of 35 primordial radionuclides there exist four even-odd nuclides (see table at right), including the fissile 235/92U. Because of their odd neutroumbers, the even-odd nuclides tend to have large neutron capture cross sections, due to the energy that results from neutron-pairing effects. These stable even-proton odd-neutrouclides tend to be uncommon by abundance iature, generally because in order to form and be enter into primordial abundance, they must have escaped capturing neutrons to form yet other stable even-even isotopes, during both the s-process and r-process of neutron capture, during nucleosynthesis in stars. For this reason, only 19578Pt and 94Be are the most naturally abundant isotopes of their element.

Only five stable nuclides contain both an odd number of protons and an odd number of neutrons. The first four “odd-odd” nuclides occur in low mass nuclides, for which changing a proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio (2/1H, 6/3Li, 10/5B, and 14/7N; spins 1, 1, 3, 1). The only other entirely “stable” odd-odd nuclide is 180/73Ta (spin 9), the only primordial nuclear isomer, which has not yet been observed to decay despite experimental attempts. Hence, all observationally stable odd-odd nuclides have nonzero integer spin. This is because the single unpaired neutron and unpaired proton have a larger nuclear force attraction to each other if their spins are aligned (producing a total spin of at least 1 unit), instead of anti-aligned. See deuterium for the simplest case of this nuclear behavior.

Many odd-odd radionuclides (like tantalum-180) with comparatively short half lives are known. Usually, they beta-decay to their nearby even-even isobars that have paired protons and paired neutrons. Of the nine primordial odd-odd nuclides (five stable and four radioactive with long half lives), only 14/7N is the most common isotope of a common element. This is the case because it is a part of the CNO cycle. The nuclides 63Li and 10/5B are minority isotopes of elements that are themselves rare compared to other light elements, while the other six isotopes make up only a tiny percentage of the natural abundance of their elements.  For example, 180/73Ta is thought to be the rarest of the 254 stable isotopes. Actinides with odd neutroumber are generally fissile (with thermal neutrons), while those with eveeutroumber are generally not, though they are fissionable with fast neutrons.

Occurrence in nature

Elements are composed of one or more naturally occurring isotopes. The unstable (radioactive) isotopes are either primordial or postprimordial. Primordial isotopes were a product of stellar nucleosynthesis or another type of nucleosynthesis such as cosmic ray spallation, and have persisted down to the present because their rate of decay is so slow (e.g., uranium-238 and potassium-40). Postprimordial isotopes were created by cosmic ray bombardment as cosmogenic nuclides (e.g., tritium, carbon-14), or by the decay of a radioactive primordial isotope to a radioactive radiogenic nuclide daughter (e.g., uranium to radium). A few isotopes also continue to be naturally synthesized as nucleogenic nuclides, by some other natural nuclear reaction, such as when neutrons from natural nuclear fission are absorbed by another atom.

As discussed above, only 80 elements have any stable isotopes, and 26 of these have only one stable isotope. Thus, about two thirds of stable elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (50Sn). There are about 94 elements found naturally on Earth (up to plutonium inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists estimate that the elements that occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total. Only 254 of these naturally occurring isotopes are stable in the sense of never having been observed to decay as of the present time. An additional 35 primordial nuclides (to a total of 289 primordial nuclides), are radioactive with known half-lives, but have half-lives longer than 80 million years, allowing them to exist from the beginning of the solar system. See list of nuclides for details.

All the known stable isotopes occur naturally on Earth; the other naturally occurring-isotopes are radioactive but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. These include the afore-mentioned cosmogenic nuclides, the nucleogenic nuclides, and any radiogenic radioisotopes formed by ongoing decay of a primordial radioactive isotope, such as radon and radium from uranium.

An additional ~3000 radioactive isotopes not found iature have been created iuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminium-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale.

The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. Before the discovery of isotopes, empirically determined noninteger values of atomic mass confounded scientists. For example, a sample of chlorine contains 75.8% chlorine-35 and 24.2% chlorine-37, giving an average atomic mass of 35.5 atomic mass units.

According to generally accepted cosmology theory, only isotopes of hydrogen and helium, traces of some isotopes of lithium and beryllium, and perhaps some boron, were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously produced isotopes. (See nucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.

Atomic mass of isotopes

The atomic mass (mr) of an isotope is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus (see mass defect), the slight difference in mass between proton and neutron, and the mass of the electrons associated with the atom, the latter because the electron:nucleon ratio differs among isotopes.

The mass number is a dimensionless quantity. The atomic mass, on the other hand, is measured using the atomic mass unit based on the mass of the carbon-12 atom. It is denoted with symbols “u” (for unified atomic mass unit) or “Da” (for dalton).

The atomic masses of naturally occurring isotopes of an element determine the atomic mass of the element. When the element contains N isotopes, the expression below is applied for the average atomic mass m_a:

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element. Isotope separation is a significant technological challenge, particularly with heavy elements such as uranium or plutonium. Lighter elements such as lithium, carbon, nitrogen, and oxygen are commonly separated by gas diffusion of their compounds such as CO and NO. The separation of hydrogen and deuterium is unusual since it is based on chemical rather than physical properties, for example in the Girdler sulfide process. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectrometry.

Use of chemical and biological properties

Isotope analysis is the determination of isotopic signature, the relative abundances of isotopes of a given element in a particular sample. For biogenic substances in particular, significant variations of isotopes of C, N and O can occur. Analysis of such variations has a wide range of applications, such as the detection of adulteration in food products or the geographic origins of products using isoscapes. The identification of certain meteorites as having originated on Mars is based in part upon the isotopic signature of trace gases contained in them.

Isotopic substitution can be used to determine the mechanism of a chemical reaction via the kinetic isotope effect.

Another common application is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, even different nonradioactive stable isotopes can be distinguished by mass spectrometry or infrared spectroscopy. For example, in ‘stable isotope labeling with amino acids in cell culture (SILAC)’ stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling).

Use of nuclear properties

A technique similar to radioisotopic labeling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials.

Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes, both radioactive and stable. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are 1H, 2D,15N, 13C, and 31P.

Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as 57Fe.

Radionuclides also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes. Nuclear medicine and radiation oncology utilize radioisotopes respectively for medical diagnosis and treatment.

Allotropes

Atoms of pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple structures (spacial arrangements of atoms), known as allotropes, which differ in their properties. For example, carbon can be found as diamond, which has a tetrahedral structure around each carbon atom; graphite, which has layers of carbon atoms with a hexagonal structure stacked on top of each other; graphene, which is a single layer of graphite that is incredibly strong; fullerenes, which have nearly spherical shapes; and carboanotubes, which are tubes with a hexagonal structure (even these may differ from each other in electrical properties). The ability for an element to exist in one of many structural forms is known as ‘allotropy’.

History

The concept of allotropy was originally proposed in 1841 by the Swedish scientist Baron Jöns Jakob Berzelius (1779–1848). The term is derived from the Greek άλλοτροπἱα (allotropia; variability, changeableness). After the acceptance of Avogadro’s hypothesis in 1860 it was understood that elements could exist as polyatomic molecules, and the two allotropes of oxygen were recognized as O₂ and O₃. In the early 20th century it was recognized that other cases such as carbon were due to differences in crystal structure.

By 1912, Ostwald noted that the allotropy of elements is just a special case of the phenomenon of polymorphism known for compounds, and proposed that the terms allotrope and allotropy be abandoned and replaced by polymorph and polymorphism.Although many other chemists have repeated this advice, IUPAC and most chemistry texts still favour the usage of allotrope and allotropy for elements only.

Differences in properties of an element’s allotropes

Allotropes are different structural forms of the same element and can exhibit quite different physical properties and chemical behaviours. The change between allotropic forms is triggered by the same forces that affect other structures, i.e. pressure, light, and temperature. Therefore the stability of the particular allotropes depends on particular conditions. For instance, iron changes from a body-centered cubic structure (ferrite) to a face-centered cubic structure (austenite) above 906 °C, and tin undergoes a transformation known as tin pest from a metallic form to a semiconductor form below 13.2 °C (55.8 °F). As an example of allotropes having different chemical behaviour, ozone (O₃) is a much stronger oxidizing agent than dioxygen (O₂).

Non-metals

 

Metalloids

Metals

Among the metallic elements that occur iature in significant quantities (up to U, without Tc and Pm), 27 are allotropic at ambient pressure: Li, Be, Na, Ca, Ti, Mn, Fe, Co, Sr, Y, Zr, Sn, La, Ce, Pr, Nd, Sm, Gd, Tb, Dy, Yb, Hf, Tl, Th, Pa and U. Some phase transitions between allotropic forms of technologically-relevant metals are those of Ti at 882˚C, Fe at 912˚C and 1394˚C, Co at 422˚C, Zr at 863˚C, Sn at 13˚C and U at 668˚C and 776˚C.

Lanthanides and actinides

Cerium, samarium, terbium, dysprosium and ytterbium have three allotropes.

Praseodymium, neodymium, gadolinium and terbium have two allotropes.

Plutonium has six distinct solid allotropes under “normal” pressures. Their densities vary within a ratio of some 4:3, which vastly complicates all kinds of work with the metal (particularly casting, machining, and storage). A seventh plutonium allotrope exists at very high pressures. The transuranium metals Np, Am, and Cm are also allotropic.

Promethium, americium, berkelium and californium have 3 allotropes each.

 

Standard state

The standard state, or reference state, of an element is defined as its thermodynamically most stable state at 1 bar at a given temperature (typically at 298.15 K). In thermochemistry, an element is defined to have an enthalpy of formation of zero in its standard state. For example, the reference state for carbon is graphite, because it is more stable than the other allotropes.

The naming of elements precedes the atomic theory of matter, although at the time it was not known which chemicals were elements and which compounds. When these facts were learned, the existing names for anciently-known elements (e.g., gold, mercury, iron) were kept in most countries. National differences emerged over the names of elements either for convenience, linguistic niceties, or nationalism. For a few illustrative examples: German speakers use “Wasserstoff” (water substance) for “hydrogen”, “Sauerstoff” (acid substance) for “oxygen” and “Stickstoff” (smothering substance) for “nitrogen”, while English and some romance languages use “sodium” for “natrium” and “potassium” for “kalium”, and the French, Italians, Greeks, Portuguese and Poles prefer “azote/azot/azoto” (from roots meaning “no life”) for “nitrogen”.

For purposes of international trade, the official names of the chemical elements both ancient and recent are decided by the International Union of Pure and Applied Chemistry, which has decided on a sort of international English language. That organization has recently prescribed that “aluminium” and “caesium” take the place of the U.S. spellings “aluminum” and “cesium”, while the U.S. “sulfur” takes the place of the British “sulphur”. Chemicals that are practical to sell in bulk in many countries, however, still have national names. Also, those countries that do not use the Latin alphabet to write their national language, cannot be expected to use the IUPAC name for elements.

Symbols of chemical elements, however, are capitalized: thus the symbols for the elements just discussed are Cf and Es; C-12 and U-235. IUPAC prefers that isotope symbols be written in superscript notation, however: ¹²C and ²³⁵U.

In the second half of the twentieth century physics laboratories became able to produce nuclei of chemical elements that have a half life too short for them to remain in any appreciable amounts. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This can lead to the controversial question of which research group actually discovered an element, a question that delayed naming of elements with atomic number of 104 and higher for a considerable time.

Precursors of such controversies involved the nationalistic namings of elements in the late nineteenth century. For example, lutetium was named in reference to Paris, France. The Germans were reluctant to relinquish naming rights to the French, often calling it cassiopeium. Similarly, the British discoverer of niobium originally named it columbium, in reference to the New World. It was used extensively as such by American publications prior to international standardization.

Isotope symbols

The three main isotopes of the element hydrogen are often written as H for protium, D for deuterium and T for tritium. This is in order to make it easier to use them in chemical equations, as it replaces the need to write out the mass number for each atom. E.g. the formula for heavy water may be written D₂O instead of ²H₂O.

The periodic table

The periodic table of the chemical elements is a tabular method of displaying the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869. Mendeleev intended the table to illustrate recurring (“periodic”) trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.

The periodic table is now ubiquitous within the academic discipline of chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behavior. The table has also found wide application in physics, biology, engineering, and industry. The current standard table contains 118 confirmed elements as of April 10, 2010.

Periodic table

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

Al

Si

P

S

Cl

Ar

K

Ca

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Zn

Ga

Ge

As

Se

Br

Kr

Rb

Sr

Y

Zr

Nb

Mo

Tc

Ru

Rh

Pd

Ag

Cd

In

Sn

Sb

Te

I

Xe

Cs

Ba

La

Ce

Pr

Nd

Pm

Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

Lu

Hf

Ta

W

Re

Os

Ir

Pt

Au

Hg

Tl

Pb

Bi

Po

At

Rn

Fr

Ra

Ac

Th

Pa

U

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

Lr

Rf

Db

Sg

Bh

Hs

Mt

Ds

Rg

Cn

Uut

Uuq

Uup

Uuh

Uus

Uuo

Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Other metals Metalloids Other nonmetals Halogens Noble gases

 

 

Abundance

During the early phases of the Big Bang, nucleosynthesis of hydrogen nuclei resulted in the production of hydrogen and helium isotopes, as well as very minuscule amounts (on the order of 10⁻¹⁰) of lithium and beryllium. There is evidence that some boron was produced in the Big Bang, since it has been observed in some very young stars, even though carbon has not. There is agreement that no heavier elements than boron were produced in the Big Bang. As a result, the primordial abundance of atoms consisted of roughly 75% ¹H, 25% ⁴He, and 0.01% deuterium. Subsequent enrichment of galactic halos occurred due to stellar nucleosynthesis and supernova nucleosynthesis. However intergalactic space can still closely resemble the primordial abundance, unless it has been enriched by some means.

The following graph (note log scale) shows abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass. Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. As physical laws and processes appear common throughout the visible universe, however, it is expected that these galaxies will likewise have evolved similar abundances of elements.

Abundances of the chemical elements in the Solar system. Hydrogen and helium are most common, from the Big Bang. The next three elements (Li, Be, B) are rare because they are poorly synthesized in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers, and (2) a general decrease in abundance, as elements become heavier.

The Biogenic Elements: The Chemical Elements Essential To Earth’s Living Systems

Almost every one of the chemical elements plays some role in Earth’s living systems, however, ~20 elements account for the vast majority of material in living systems. These biogenic elements are divided into …

→ six major biogenic elements (elements found in almost all of Earth’s living systems, often in relatively large quantities),

→ five minor biogenic elements (elements found in many of Earth’s living systems, and/or in relatively small quantities),

→ trace elements (essential elements necessary only in very small quantities to maintain the chemical reactions on which life depends.

The biogenic  elements can be classified as:

 MACROELEMENTS (or macrominerals) – the content in the organism is more than 10^-2%

 MICROELEMENTS (or trace elements) – the content in the organism is 10^-3 – 10^-5 %

Macroelements are elements found in almost all of Earth`s living systems. There are 11 of them. Six are colled organogens ore major biogenic elements. The content of them is 97% in the organism.

Major Biogenic Elements or Organogens

Carbon, Hydrogen, Oxygen, Nitrogen, Sulfur, Phosphorous

Minor Biogenic Elements

Sodium, Potassium, Magnesium, Calcium, Chlorine

Microelements are essential elements necessary only in very small quantities to maintain the chemical reaction on which life depends. These are:

Biogenic Trace Elements

Manganese, Iron, Cobalt, Copper, Zinc, Boron, Aluminum, Vanadium, Molybdenum, Iodine, Silicon, Nickel, Bromine

According to their abundance in the organism , biogenic elements can be classsified into macroelements, microelements and contaminating elements.

Macroelements (12 elements in total) form up to 99 % of any organism, and can be further subdivided into:

a) a group of stable primary elements (1-60 % of  total organism weight). These are: O,C, H,N,

b) a group of stable secondary elements (0.05/1 % of total organism weight). These are Ca, S, Mg, Cl, Na, K, Fe

Microelements can be divided into three categories:

a) a subgroup of 8 stable elements (less than 0.05%). These are the elements: Cu, Zn, Mn, Co, B, Si, F, I

b) a subgroup of approximately 20 elements that are present at conc. of 0.001% and lower.

c) a subgroup of contaminating elements: Their constant excess in the organism leads to disease: Mn, He, Ar, Hg, Tl, Bi, Al, Cr, Cd.

According to their physiological importance, biogenic elements are essential and nonessential.

– Essential elements can be divided into two groups:

a) main essential elements, such as Ca, Mg, Na, K, P, S, Cl, they are present in food. They usually play multiple roles.

b) Trace essential elements.

By mass, human cells consist of 65–90% water (H₂O), and a significant portion of the remainder is composed of carbon-containing organic molecules. Oxygen therefore contributes a majority of a human body’s mass, followed by carbon. Almost 99% of the mass of the human body is made up of six elements: oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus. The next 0.75% is made up of the next five elements: potassium, sulfur, chlorine, sodium, and magnesium. Only 17 elements are known for certain to be necessary to human life, with one additional element (fluorine) thought to be helpful for tooth enamel strength. A few more trace elements appear to be necessary to mammals in carefully dust-free conditions. Boron and silicon are notably necessary for plants but have uncertain roles in animals. The elements aluminium and silicon, although very common in the earth’s crust, are conspicuously rare in the human body.

CALCIUM: (Ca)

A soft grey metal from Group II of the periodic table. It is moderately reactive. Calcium is the fifth most common element in the Earth System by weight. Calcium compounds are found in many common rocks, e.g., limestone and chalk (calcium carbonate, CaCO₃).

Calcium is essential to living things:

Plants require calcium to build cell walls. Animals require calcium to build bones and teeth.

Calcium is present in bones in the form of hydroxyapatite (Ca₃(PO₄)₂)₂ , Ca(OH)₂.

Calcium present in ionic form, serving as an important regulator of processes in cell cytoplasm.

Biological importance of calcium:

•it affects neuromuscular excitability of muscle (together with ions K, Na and Mg and takes part in muscle contraction. Hypocalcemia leads to cramps-tetany

•it takes part in regulation of glycogenolysis in muscle and of  gluconeogenesis in kidney and liver

•it decreases cell membrane and capillary wall permeability, what results in its antiinflammatory, antiexudative and antiallergic effects.

•it is necessary for blood coagulation.

•calcium ions are important intracellular messengers, influencing secretion of insulin into the circulation and secretion of digestion enzymes into small intestine. Calcium resorption is affected by the mutual ratio of  calcium to phosphates in the intestinal content, and bythe presence of cholecalciferol, which regulates active resorption of calcium and phosphorus.

Exchange of calcium and phosphates is regulated  hormonally by parathormone and calcitonin.  Parathormone releases calcium from bones into blood. Calcitonin promotes deposition of calcium into bones, which decreases its blood concentration.

Lime has been used as a material for building since 7000 to 14,000 BCE, and kilns used for lime have been dated to 2,500 BCE in Khafaja, Mesopotamia. Calcium as a material has been known since at least the first century, as the ancient Romans were known to have used calcium oxide by preparing it from lime. Calcium sulfate has been known to be able to set broken bones since the tenth century. Calcium itself, however, was not isolated until 1808, when Humphry Davy, in England, used electrolysis on a mixture of lime and mercuric oxide,  after hearing that Jöns Jakob Berzelius had prepared a calcium amalgam from the electrolysis of lime in mercury.

In chemical terms, calcium is reactive and soft for a metal (though harder than lead, it can be cut with a knife with difficulty). It is a silvery metallic element that must be extracted by electrolysis from a fused salt like calcium chloride. Once produced, it rapidly forms a gray-white oxide and nitride coating when exposed to air. In bulk form (typically as chips or “turnings”), the metal is somewhat difficult to ignite, more so even than magnesium chips; but, when lit, the metal burns in air with a brilliant high-intensity orange-red light. Calcium metal reacts with water, generating hydrogen gas at a rate rapid enough to be noticeable, but not fast enough at room temperature to generate much heat, making it useful for generating hydrogen. In powdered form, however, the reaction with water is extremely rapid, as the increased surface area of the powder accelerates the reaction with the water. Part of the slowness of the calcium–water reaction results from the metal being partly protected by insoluble white calcium hydroxide. In water solutions of acids, where this salt is soluble, calcium reacts vigorously.

Calcium, with a density of 1.55 g/cm3, is the lightest of the alkaline earth metals; magnesium (specific gravity 1.74) and beryllium (1.84) are more dense, although lighter in atomic mass. From strontium onward, the alkali earth metals become more dense with increasing atomic mass.

It has two allotropes

Calcium has a higher electrical resistivity than copper or aluminium, yet weight-for-weight, due to its much lower density, it is a rather better conductor than either. However, its use in terrestrial applications is usually limited by its high reactivity with air.

Calcium salts are colorless from any contribution of the calcium, and ionic solutions of calcium (Ca2+) are colorless as well. As with magnesium salts and other alkaline earth metal salts, calcium salts are often quite soluble in water. Notable exceptions include the hydroxide, the sulfate (unusual for sulfate salts), the carbonate and the phosphates. With the exception of the sulfate, even the insoluble ones listed are in general more soluble than its transition metal counterparts. When in solution, the calcium ion to the human taste varies remarkably, being reported as mildly salty, sour, “mineral like” or even “soothing.” It is apparent that many animals can taste, or develop a taste, for calcium, and use this sense to detect the mineral in salt licks or other sources. In humautrition, soluble calcium salts may be added to tart juices without much effect to the average palate.

Calcium is the fifth-most-abundant element by mass in the human body, where it is a common cellular ionic messenger with many functions, and serves also as a structural element in bone. It is the relatively high-atomic-number calcium in the skeleton that causes bone to be radio-opaque. Of the human body’s solid components after drying and burning of organics (as for example, after cremation), about a third of the total “mineral” mass remaining, is the approximately one kilogram of calcium that composes the average skeleton (the remainder being mostly phosphorus and oxygen).

MAGNESIUM (Mg):

Magnesium is the seventh most common element in the Earth System by weight. Magnesium is an essential component of chlorophyll.

• Is a typical intracellular cation.

• Is an essential part of tissues and body fluids.

Is present in skeleton (70 %) and muscles of animals .

Elemental magnesium is a rather strong, silvery-white, light-weight metal (two thirds the density of aluminium). It tarnishes slightly when exposed to air, although unlike the alkali metals, an oxygen-free environment is unnecessary for storage because magnesium is protected by a thin layer of oxide that is fairly impermeable and difficult to remove. Like its lower periodic table group neighbor calcium, magnesium reacts with water at room temperature, though it reacts much more slowly than calcium. When submerged in water, hydrogen bubbles almost unnoticeably begin to form on the surface of the metal—though if powdered, it reacts much more rapidly. The reaction occurs faster with higher temperatures (see precautions). Magnesium’s ability to react with water can be harnessed to produce energy and run a magnesium-based engine. Magnesium also reacts exothermically with most acids, such as hydrochloric acid (HCl). As with aluminium, zinc and many other metals, the reaction with hydrochloric acid produces the chloride of the metal and releases hydrogen gas.

Magnesium is a highly flammable metal, but while it is easy to ignite when powdered or shaved into thin strips, it is difficult to ignite in mass or bulk. Once ignited, it is difficult to extinguish, being able to burn iitrogen (forming magnesium nitride), carbon dioxide (forming magnesium oxide and carbon) and water (forming magnesium oxide and hydrogen). This property was used in incendiary weapons used in the firebombing of cities in World War II, the only practical civil defense being to smother a burning flare under dry sand to exclude the atmosphere. On burning in air, magnesium produces a brilliant white light that includes strong ultraviolet. Thus magnesium powder (flash powder) was used as a source of illumination in the early days of photography. Later, magnesium ribbon was used in electrically ignited flash bulbs. Magnesium powder is used in the manufacture of fireworks and marine flares where a brilliant white light is required. Flame temperatures of magnesium and magnesium alloys can reach 3,100 °C (3,370 K; 5,610 °F), although flame height above the burning metal is usually less than 300 mm (12 in). Magnesium may be used as an ignition source for thermite, a mixture of aluminium and iron oxide powder that is otherwise difficult to ignite. Those properties are due to magnesium’s high specific heat, the fourth highest specific heat among the metals.

Magnesium compounds are typically white crystals. Most are soluble in water, providing the sour-tasting magnesium ion Mg2+. Small amounts of dissolved magnesium ion contribute to the tartness and taste of natural waters. Magnesium ion in large amounts is an ionic laxative, and magnesium sulfate (commoame: Epsom salt) is sometimes used for this purpose. So-called “milk of magnesia” is a water suspension of one of the few insoluble magnesium compounds, magnesium hydroxide. The undissolved particles give rise to its appearance and name. Milk of magnesia is a mild base commonly used as an antacid, which has some laxative side effect.

Alloy

As of 2013 magnesium alloy consumption is less than a million tons per year, compared with 50 million tons of aluminum alloys. Its use has been historically limited by its tendency to corrode, high-temperature creep and flammability.

Research and development eliminated magnesium’s tendency toward high-temperature creep by inclusion of scandium and gadolinium. Flammability was greatly reduced by introducing a small amount of calcium into the mix.

The presence of iron, nickel, copper and cobalt strongly activates corrosion. This is due to their low solid solubility limits (above a very small percentage they precipitate out as intermetallic compounds) and because they behave as active cathodic sites that reduce water and cause the loss of magnesium.

Reducing the quantity of these metals improves corrosion resistance. Sufficient manganese overcomes the corrosive effects of iron. This requires precise control over composition, increasing costs.

Adding a cathodic poison captures atomic hydrogen within the structure of a metal. This prevents the formation of free hydrogen gas which is required for corrosive chemical processes. The addition of about one-third of a percent of arsenic reduces its corrosion rate in a salt solution by a factor of nearly ten.

 

SODIUM (Na):

Sodium is the eighth most common element in the Earth System by weight.

• Important extracellular cation, take part in homeostasis of  the organism.

• protects organism form excessive water losses.

• takes part in spreading of nerve excitation.

Sodium is a chemical element with the symbol Na (from Latin: natrium) and atomic number 11. It is a soft, silver-white, highly reactive metal and is a member of the alkali metals; its only stable isotope is 23Na. The free metal does not occur iature, but instead must be prepared from its compounds; it was first isolated by Humphry Davy in 1807 by the electrolysis of sodium hydroxide. Sodium is the sixth most abundant element in the Earth’s crust, and exists iumerous minerals such as feldspars, sodalite and rock salt. Many salts of sodium are highly water-soluble, and their sodium has been leached by the action of water so that chloride and sodium (NaCl) are the most common dissolved elements by weight in the Earth’s bodies of oceanic water.

Many sodium compounds are useful, such as sodium hydroxide (lye) for soap-making, and sodium chloride for use as a deicing agent and a nutrient (edible salt). Sodium is an essential element for all animals and some plants. In animals, sodium ions are used against potassium ions to build up charges on cell membranes, allowing transmission of nerve impulses when the charge is dissipated. The consequent need of animals for sodium causes it to be classified as a dietary inorganic macro-mineral.

Physical

Sodium at standard temperature and pressure is a soft metal that can be readily cut with a knife and is a good conductor of electricity. Freshly exposed, sodium has a bright, silvery luster that rapidly tarnishes, forming a white coating of sodium hydroxide and sodium carbonate. These properties change at elevated pressures: at 1.5 Mbar, the color changes to black, then to red transparent at 1.9 Mbar, and finally clear transparent at 3 Mbar. All of these allotropes are insulators and electrides.

When sodium or its compounds are introduced into a flame, they turn it yellow, because the excited 3s electrons of sodium emit a photon when they fall from 3p to 3s; the wavelength of this photon corresponds to the D line at 589.3 nm. Spin-orbit interactions involving the electron in the 3p orbital split the D line into two; hyperfine structures involving both orbitals cause many more lines.

Chemical

Sodium is generally less reactive than potassium and more reactive than lithium. Like all the alkali metals, it reacts exothermically with water, to the point that sufficiently large pieces melt to a sphere and may explode; this reaction produces caustic sodium hydroxide and flammable hydrogen gas. When burned in dry air, it mainly forms sodium peroxide as well as some sodium oxide. In moist air, sodium hydroxide results. Sodium metal is highly reducing, with the reduction of sodium ions requiring −2.71 volts. Hence, the extraction of sodium metal from its compounds (such as with sodium chloride) uses a significant amount of energy. However, potassium and lithium have even more negative potentials.

Sodium in biology

In humans, sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day. Sodium chloride is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for pickling and jerky; most of it comes from processed foods. The DRI for sodium is 2.3 grams per day, but on average people in the United States consume 3.4 grams per day, the minimum amount that promotes hypertension;  this in turn causes 7.6 million premature deaths worldwide.

The renin-angiotensin system regulates the amount of fluids and sodium in the body. Reduction of blood pressure and sodium concentration in the kidney result in the production of renin, which in turn produces aldosterone and angiotensin, retaining sodium in the urine. Because of the increase in sodium concentration, the production of renin decreases, and the sodium concentration returns to normal. Sodium is also important ieuron function and osmoregulation between cells and the extracellular fluid, their distribution mediated in all animals by Na+/K+-ATPase; hence, sodium is the most prominent cation in extracellular fluid.

In C4 plants, sodium is a micronutrient that aids in metabolism, specifically in regeneration of phosphoenolpyruvate and synthesis of chlorophyll. In others, it substitutes for potassium in several roles, such as maintaining turgor pressure and aiding in the opening and closing of stomata. Excess sodium in the soil limits the uptake of water due to decreased water potential, which may result in wilting; similar concentrations in the cytoplasm can lead to enzyme inhibition, which in turn causes necrosis and chlorosis. To avoid these problems, plants developed mechanisms that limit sodium uptake by roots, store them in cell vacuoles, and control them over long distances; excess sodium may also be stored in old plant tissue, limiting the damage to new growth.

Precautions

Care is required in handling elemental sodium, as it generates flammable hydrogen and caustic sodium hydroxide upon contact with water; powdered sodium may spontaneously explode in the presence of oxygen. Excess sodium can be safely removed by hydrolysis in a ventilated cabinet; this is typically done by sequential treatment with isopropanol, ethanol and water. Isopropanol reacts very slowly, generating the corresponding alkoxide and hydrogen. Fire extinguishers based on water accelerate sodium fires; those based on carbon dioxide and bromochlorodifluoromethane lose their effectiveness when they dissipate. An effective extinguishing agent is Met-L-X, which comprises approximately 5% Saran in sodium chloride together with flow agents; it is most commonly hand-applied with a scoop. Other materials include Lith+, which has graphite powder and an organophosphate flame retardant, and dry sand.

POTASSIUM: (K)

Potassium is the sixth most common element in the Earth  System by weight important intracellular ion

•take part in homeostasis of the organism.

•takes part in spreading of nerve excitation.

•deficiency of potassium leads to heart arrest.

Potassium is a chemical element with symbol K (from Neo-Latin kalium) and atomic number 19. Elemental potassium is a soft silvery-white alkali metal that oxidizes rapidly in air and is very reactive with water, generating sufficient heat to ignite the hydrogen emitted in the reaction and burning with a lilac flame.

Because potassium and sodium are chemically very similar, their salts were not at first differentiated. The existence of multiple elements in their salts was suspected from 1702, and this was proven in 1807 when potassium and sodium were individually isolated from different salts by electrolysis. Potassium in nature occurs only in ionic salts. As such, it is found dissolved in seawater (which is 0.04% potassium by weight and is part of many minerals.

Most industrial chemical applications of potassium employ the relatively high solubility in water of potassium compounds, such as potassium soaps. Potassium metal has only a few special applications, being replaced in most chemical reactions with sodium metal.

Potassium ions are necessary for the function of all living cells. Potassium ion diffusion is a key mechanism ierve transmission, and potassium depletion in animals, including humans, results in various cardiac dysfunctions. Potassium accumulates in plant cells, and thus fresh fruits and vegetables are a good dietary source of it. Conversely, most plants except specialist halophytes are intolerant of salt, and sodium is present in them only in low concentration. This resulted in potassium first being isolated from potash, the ashes of plants, giving the element its name. For the same reason, heavy crop production rapidly depletes soils of potassium, and agricultural fertilizers consume 95% of global potassium chemical production.

Potassium atoms have 19 electrons, which is one more than the extremely stable configuration of argon. A potassium atom is thus much more likely to lose the “extra” electron than to gain one; however, the alkalide ions, K–, are known. Because of the low first ionization energy (418.8 kJ/mol) the potassium atom easily loses an electron and oxidizes into the monopositive cation, K+. This process requires so little energy that potassium is readily oxidized by atmospheric oxygen. In contrast, the second ionization energy is very high (3052 kJ/mol), because removal of two electrons breaks the stable noble gas electronic configuration. Potassium therefore does not readily form compounds with the oxidation state of +2 (or higher).

Potassium is the second least dense metal after lithium. It is a soft solid that has a low melting point and can easily be cut with a knife. Freshly cut potassium is silvery in appearance, but it begins to tarnish toward gray immediately after being exposed to air. In a flame test, potassium and its compounds emit a lilac color with a peak emission wavelength of 766.5 nm (see movie below).

Chemical

Potassium is an extremely active metal, which reacts violently with oxygen and water in air. With oxygen, it converts to potassium peroxide and with water potassium hydroxide. The reaction of potassium with water is dangerous because of its violent exothermic character and the production of hydrogen gas. Hydrogen reacts again with atmospheric oxygen, producing water, which reacts with the remaining potassium. This reaction requires only traces of water; because of this, potassium and its liquid alloy with sodium — NaK — are potent desiccants that can be used to dry solvents prior to distillation.

Because of the sensitivity of potassium to water and air, the reactions are possible only in inert atmosphere, such as argon gas using air-free techniques. Potassium does not react with most hydrocarbons, such as mineral oil or kerosene. It readily dissolves in liquid ammonia, up to 480 g per 1000 g of ammonia at 0 °C. Depending on the concentration, the ammonia solutions are blue to yellow, and their electrical conductivity is similar to that of liquid metals. In a pure solution, potassium slowly reacts with ammonia to form KNH2, but this reaction is accelerated by minute amounts of transition metal salts. It can reduce the salts to the metal; potassium is often used as the reductant in the preparation of finely divided metals from their salts by the Rieke method. For example, the preparation of Rieke magnesium employs potassium as the reductant:

MgCl₂ + 2 K → Mg + 2 KCl

 

PHOSPHORUS (P):

Phosphorus is found in ATP, nucleic acids, cell membranes, and animal bones.

• its main function is structural, in bones as calcium phosphate.

• in buffer systems as a phosphate anions.

Phosphorus is a chemical element with symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus as a mineral is almost always present in its maximally oxidised state, as inorganic phosphate rocks. Elemental phosphorus exists in two major forms—white phosphorus and red phosphorus—but due to its high reactivity, phosphorus is never found as a free element on Earth.

The first form of elemental phosphorus to be produced (white phosphorus, in 1669) emits a faint glow upon exposure to oxygen – hence its name given from Greek mythology, Φωσφόρος meaning “light-bearer” (Latin Lucifer), referring to the “Morning Star”, the planet Venus. The term “phosphorescence”, meaning glow after illumination, originally derives from this property of phosphorus, although this word has since been used for a different physical process that produces a glow. The glow of phosphorus itself originates from oxidation of the white (but not red) phosphorus— a process now termed chemiluminescence.

The vast majority of phosphorus compounds are consumed as fertilisers. Other applications include the role of organophosphorus compounds in detergents, pesticides and nerve agents, and matches.

Phosphorus is essential for life. As phosphate, it is a component of DNA, RNA, ATP, and also the phospholipids that form all cell membranes. Demonstrating the link between phosphorus and life, elemental phosphorus was historically first isolated from human urine, and bone ash was an important early phosphate source. Phosphate minerals are fossils. Low phosphate levels are an important limit to growth in some aquatic systems. The chief commercial use of phosphorus compounds for production of fertilisers is due to the need to replace the phosphorus that plants remove from the soil.

Biological role

Inorganic phosphorus in the form of the phosphate PO3−4 is required for all known forms of life, playing a major role in biological molecules such as DNA and RNA where it forms part of the structural framework of these molecules. Living cells also use phosphate to transport cellular energy in the form of adenosine triphosphate (ATP). Nearly every cellular process that uses energy obtains it in the form of ATP. ATP is also important for phosphorylation, a key regulatory event in cells. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts assist in stiffening bones.

Living cells are defined by a membrane that separates it from its surroundings. Biological membranes are made from a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol, such that two of the glycerol hydroxyl (OH) protons have been replaced with fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.

An average adult human contains about 0.7 kg of phosphorus, about 85–90% of which is present in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids (~1%). The phosphorus content increases from about 0.5 weight% in infancy to 0.65–1.1 weight% in adults. Average phosphorus concentration in the blood is about 0.4 g/L, about 70% of that is organic and 30% inorganic phosphates. A well-fed adult in the industrialized world consumes and excretes about 1–3 grams of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as nucleic acids and phospholipids; and excretion almost exclusively in the form of phosphate ions such as H₂PO₄−4 and HPO₂−4. Only about 0.1% of body phosphate circulates in the blood, and this amount reflects the amount of phosphate available to soft tissue cells.

Bone and teeth enamel

The main component of bone is hydroxyapatite as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. Water fluoridation enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material called fluoroapatite.

 

Phosphorus deficiency

In medicine, low-phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as re-feeding after malnutrition) or pass too much of it into the urine. All are characterized by hypophosphatemia, which is a condition of low levels of soluble phosphate levels in the blood serum, and therefore inside cells. Symptoms of hypophosphatemia include muscle and neurological dysfunction, and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhoea and calcification (hardening) of organs and soft tissue, and can interfere with the body’s ability to use iron, calcium, magnesium, and zinc.

Phosphorus is an essential macromineral for plants, which is studied extensively in edaphology in order to understand plant uptake from soil systems. In ecological terms, phosphorus is often a limiting factor in many environments; i.e. the availability of phosphorus governs the rate of growth of many organisms. In ecosystems an excess of phosphorus can be problematic, especially in aquatic systems, resulting in eutrophication which sometimes lead to algal blooms.

Food sources

The main food sources for phosphorus are foods containing protein, although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if one holds a meal plan providing sufficient amount of protein and calcium then the amount of phosphorus is also likely sufficient.

Precautions

Organic compounds of phosphorus form a wide class of materials, many are required for life, but some are extremely toxic. Fluorophosphate esters are among the most potent neurotoxins known. A wide range of organophosphorus compounds are used for their toxicity to certain organisms as pesticides (herbicides, insecticides, fungicides, etc.) and weaponised as nerve agents. Most inorganic phosphates are relatively nontoxic and essential nutrients.

The white phosphorus allotrope presents a significant hazard because it ignites in air and produces phosphoric acid residue. Chronic white phosphorus poisoning leads to necrosis of the jaw called “phossy jaw”. Ingestion of white phosphorus may cause a medical condition known as “Smoking Stool Syndrome”.

Phosphorus explosion

Upon exposure to elemental phosphorus, in the past it was suggested to wash the affected area with 2% copper sulfate solution to form harmless compounds that can be washed away. According to the recent US Navy’s Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries, “Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis.

The manual suggests instead “a bicarbonate solution to neutralize phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots. Promptly debride the burn if the patient’s condition will permit removal of bits of WP (white phosphorus) that might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns.” As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.

REACTIVITY WITH PHOSPHORUS

1. Reaction with air

White phosphorus is very reactive. It has an appreciable vapour pressure at room temperature and inflames in dry air at about 320 K or at even lower temperatures if finely divided. In air at room temperature it emits a faint green light called phosphorescence; the reaction occurring is a complex oxidation process, but this happens only at certain partial pressures of oxygen. It is necessary, therefore, to store white phosphorus under water, unlike the less reactive red and black allotropes which do not react with air at room temperature. Both red and black phosphorus burn to form oxides when heated in air, the red form igniting at temperatures exceeding 600 K, the actual temperature depending on purity. Black phosphorus does not ignite until even higher temperatures.

2. Reaction with acids

Hydrochloric and dilute sulphuric acids have no appreciable action at room temperature on the pure Group V elements. Concentrated sulphuric acid and nitric acid—powerful oxidizing agents—attack all the elements except nitrogen, particularly when the acids are warm. The products obtained reflect changes in stability of the oxidation states V and III of the Group V elements. Both white and red phosphorus dissolve in, for example, concentrated nitric acid to form phosphoric(V) acid, the reaction between hot acid and white phosphorus being particularly violent. Arsenic dissolves in concentrated nitric acid forming arsenic(V) acid, H₃AsO₄, but in dilute nitric acid and concentrated sulphuric acid the main product is the arsenic(III) acid, H₃AsO₃. The more metallic element, antimony, dissolves to form the (III) oxide Sb₄O₆ with moderately concentrated nitric acid, but the (V) oxide Sb₂O₅ (structure unknown) with the more concentrated acid. Bismuth, however, forms the salt bismuth(III) nitrate Bi(NO₃)₃ 5H₂O.

3As + 5HNO₃ + 2H₂O = 3H₃AsO₄ + 5NO

3Sb + 5HNO₃ = 3HSbO₃ + 5NO + H₂O

Bi + 4HNO₃ = Bi(NO₃)₃ + NO + 2H₂O

 

3. Reaction with alkalis

The change from non-metallic to metallic properties of the Group V elements as the atomic mass of the element increases is shown in their reactions with alkalis. The head element nitrogen does not react. White phosphorus, however, reacts when warmed with a concentrated solution of a strong alkali to form phosphine, a reaction which can be regarded as a disproportionation reaction of phosphorus:

P₄ + 3KOH + 3H₂O → 3KH₂PO₂ + PH₃↑

potassium phosphine

phosphinate

hypophosphite

The phosphine produced is impure and contains small quantities of diphosphane, P₂H₄.

Arsenic, unlike phosphorus, is only slightly attacked by boiling sodium hydroxide; more rapid attack takes place with the fused alkali; an arsenate(III) is obtained in both cases,

As₄ + 12OH ^– → 4AsO₃ ^3- + 6H₂↑ cf. aluminiu.

Arsine is not formed in this reaction. Antimony and bismuth do not react with sodium hydroxide.

4. Reaction with halogens

Nitrogen does form a number of binary compounds with the halogens but none of these can be prepared by the direct combination of the elements and they are dealt with below. The other Group V elements all form halides by direct combination.

White and red phosphorus combine directly with chlorine, bromine and iodine, the red allotrope reacting in each case at a slightly higher temperature. The reactions are very vigorous and white phosphorus is spontaneously inflammable in chlorine at room temperature. Both chlorine and bromine first form a trihalide:

P₄ + 6X₂ → 4PX₃   (X = Cl or Br)

2P + 3Cl₂ ® 2PCl₃

PCl₃ + 2H₂O = H₂[PO₂H] + 3HCl

PCl₅ + HOH Û POCl₃ + 2HCl

P₂O₅ + 3PCl₅ ® 5POCl₃

 

but this is converted to a pentahalide by excess of the halogen. No pentaiodide is known.

 

SULFUR (S):

→ A yellow non-metallic element in Group VI of the periodic table. Sulfur is reactive with both metals and non-metals.

→ Sulfur has two allotropes: below 96°C, rhombic sulfur is the stable form while above that temperature, monoclinic sulfur is the stable form.

→ Methionine, cysteine.

Although this element is only sixteenth in abundance at the surface of the earth, it is one of the few that has been known and used throughout history. Deposits of elemental sulfur are not uncommon, and, because they were stones that would burn, were originally called brimstone. Burning sulfur produces sulfur dioxide,

S₈(s) + 😯₂(g) → 8SO₂(g)

This colorless gas has a choking odor and is more poisonous than carbon monoxide. It is the anhydride of sulfurous acid, a weak diprotic acid:

SO₂(g) + H₂O(l) → H₂SO₃(aq)

SO₂ is also produced when almost any sulfur-containing substance is burned in air. Coal, for example, usually contains from 1 to 4% sulfur, and so burning coal releases SO₂ to the atmosphere. Many metal ores are sulfides, and when they are heated in air, SO₂ is produced. Copper, for example, may be obtained as the element by heating copper(I) sulfide:

Cu₂S(s) + O₂(g) 2Cu(s) + SO₂(g)

Since SO₂ is so poisonous, its release to the atmosphere is a major pollution problem. Once in the air, SO₂ is slowly oxidized to sulfur trioxide, SO₃:

2SO₂(g) + O₂(g) → 2SO₃(g)

This compound is the anhydride of sulfuric acid, H₂SO₄:

SO₃(g) + H₂O(l) → H₂SO₄(aq)

Thus if air is polluted with SO₂ and SO₃, a fine mist of dilute droplets of can form. All three substances are very irritating to the throat and lungs and are responsible for considerable damage to human health.

The natural mechanism for removal of sulfur oxides from the air is solution in raindrops, followed by precipitation. This makes the rainwater more acidic than it would otherwise be, and acid rain is now common in industrialized areas of the United States and Europe. Acid rain can slowly dissolve limestone and marble, both of which consist of CaCO₃:

CaCO₃(s) + H₃O⁺(aq) → Ca²⁺(aq) + HCO₃^–(aq) + H₂O(l)

Thus statues and buildings made of these materials may be damaged.

Despite the fact that a tremendous amount of sulfur is released to the environment by coal combustion and ore smelting, this element is not usually recovered from such processes. Instead it is obtained commercially from large deposits along the U.S. Gulf Coast and from refining of sour petroleum. Sour petroleum contains numerous sulfur compounds, including H₂S, which smells like rotten eggs. The deposits of elemental sulfur in Texas and Louisiana are mined by the Frasch process. Water at 170°C is pumped down a pipe to melt the sulfur, and the latter is forced to the surface by compressed air. Most of the H₂S or S₈ obtained from these sources is oxidized to SO₂, passed over a vanadium catalyst to make SO₃, and dissolved in water to make H₂SO₄. In 2005 an estimated 190 billion kg of H₂SO₄ was produced in the world, making H₂SO₄ one of the most important industrial chemicals. About half of it is used in phosphate fertilizer production.

Pure H₂SO₄ is a liquid at room temperature and has a great affinity for H₂O. This is apparently due to the reaction

H₂SO₄ + H₂O → H₃O⁺ + HSO₄^–

Formation of H₃O⁺ releases energy, and the reaction is exothermic. Concentrated H₂SO₄ is 93% H₂SO₄ and 7% H₂O by mass, corresponding to more than twice as many H₂SO₄ as H₂O molecules. Since many H₂SO₄ molecules still have protons to donate, concentrated H₂SO₄ also has a great affinity for H₂O. It is often used as a drying agent and can be employed in condensation reactions which give off H₂O.

 

TRACE essential elements :

• necessary in quantities of a miligram or microgram per day only

• play an important role in enzymatic activities, where they can be a part of the cofactor or of the prosthetic group.

Enzymatic activities can be regulated by these elements in the following ways:

1. The elements directly takes part in the reaction in the enzyme center (Cu in superoxiddismutase, Fe in catalase).

2. The element serves a mediator between the substrate and  the enzyme active center by forming a complex with both.

ACCORDING to their chemical properties, biogenic elements are metallic and nonmetallic.

NONMETALLIC ELEMENTS

A halogen halogen ( (Group VII Group VII) that is a (poisonous) green ) that is a (poisonous) green gas at room temperature. Chlorine is a strong oxidizer and so ve gas at room temperature. Chlorine is a strong oxidizer and so very reactive, forming reactive, forming chlorides chlorides with with hydrogen hydrogen and and metals metals.

CHLORINE

Chlorine (IPA: /klɔərin/, Greek: χλωρóς chloros, meaning “pale green”), is the chemical element with atomic number 17 and symbol Cl. It is a halogen, found in the periodic table in group 17 (formerly VIIa or VIIb). As the chloride ion, which is part of common salt and other compounds, it is abundant iature and necessary to most forms of life, including humans. In its common elemental form (Cl2 or “dichlorine”) under standard conditions, it is a pale green gas about 2.5 times as dense as air. It has a disagreeable, suffocating odor that is detectable in concentrations as low as 3.5 ppm and is poisonous. Chlorine is a powerful oxidant and is used in bleaching and disinfectants. As a common disinfectant, chlorine compounds are used in swimming pools to keep them clean and sanitary. In the upper atmosphere, chlorine based molecules have been implicated in the destruction of the ozone layer.

Cl₂ + F₂ = 2ClF

H₂ + Cl₂ = 2HCl

Cl₂ + NaOH ® NaClO + NaCl + H₂O

Characteristics

Chlorine gas in a plastic container. It is not advisable to store chlorine in this manner.

Chlorine gas is diatomic, with the formula Cl₂. It combines readily with all elements except O₂ and N₂ and the noble gases. Compounds with oxygen, nitrogen, and xenon are known but do not form by direct reaction of the elements. Chlorine is not as extremely reactive as fluorine. Pure chlorine gas does, however, support combustion of organic compounds such as hydrocarbons, although the carbon component tends to burn incompletely, with much of it remaining as soot. At 10 °C and atmospheric pressure, one liter of water dissolves 3.10 L of gaseous chlorine, and at 30°C, 1 L of water dissolves only 1.77 liters of chlorine.

Interaction with metals:

2Fe + 3Cl₂ = 2FeCl₃;

Cu + Cl₂ = CuCl₂.

Interaction with non-metals:

H₂ + Cl₂ = 2HCl

2S + Cl₂ = S2Cl₂

Si + 2Cl₂ = SiCl₄

2P + 5Cl = 2PCl₅.

Cl₂ + H₂O Û HCl + HClO

This element is a member of the salt-forming halogen series and is extracted from chlorides through oxidation often by electrolysis. As the chloride ion, Cl−, it is also the most abundant dissolved ion in ocean water.

Isotopes

 Chlorine has isotopes with mass numbers ranging from 32 to 40. There are two principal stable isotopes, ³⁵Cl (75.77%) and ³⁷Cl (24.23%), giving chlorine atoms in bulk an apparent atomic weight of 35.5 g/mol.

³⁶Cl Trace amounts of radioactive ³⁶Cl exist in the environment, in a ratio of about 7×10⁻¹³ to 1 with stable isotopes. ³⁶Cl is produced in the atmosphere by spallation of ³⁶Ar by interactions with cosmic ray protons. In the subsurface environment, ³⁶Cl is generated primarily as a result of neutron capture by ³⁵Cl or muon capture by ⁴⁰Ca. ³⁶Cl decays to ³⁶S and to ³⁶Ar, with a combined half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotope makes it suitable for geologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of ³⁶Cl were produced by irradiation of seawater during atmospheric detonations of nuclear weapons between 1952 and 1958. The residence time of ³⁶Cl in the atmosphere is about 1 week. Thus, as an event marker of 1950s water in soil and ground water, ³⁶Cl is also useful for dating waters less than 50 years before the present. ³⁶Cl has seen use in other areas of the geological sciences, including dating ice and sediments.

Occurrence

Iature, chlorine is found primarily as the chloride ion, a component of the salt that is deposited in the earth or dissolved in the oceans — about 1.9% of the mass of seawater is chloride ions. Even higher concentrations of chloride are found in the Dead Sea and in underground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing minerals are usually only found in abundance in dry climates or deep underground. Common chloride minerals include halite (sodium chloride), sylvite (potassium chloride), and carnallite (potassium magnesium chloride hexahydrate). Over 2000 naturally-occurring organic chlorine compounds are known.

Industrially, elemental chlorine is usually produced by the electrolysis of sodium chloride dissolved in water. Along with chlorine, this chloralkali process yields hydrogen gas and sodium hydroxide, according to the following chemical equation:

2 NaCl + 2 H₂O → Cl₂ + H₂ + 2 NaOH

FLUORINE

A halogen halogen ( (Group VII Group VII)that is a (highly poisonous)) that is a (highly poisonous) gas at room temperature. Flourine is a very strong oxidizer, for gas at room temperature. Flourine is a very strong oxidizer, forming ming fluorides fluorides with with hydrogen hydrogen and and metals metals.

Chemical uses of fluorine:

·          Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS (microelectromechanical systems) fabrication. Xenon difluoride is also used for this last purpose.

•  Hydrofluoric acid (chemical formula HF) is used to etch glass in light bulbs and other products.

• Fluorine is indirectly used in the production of low friction plastics such as Teflon (or polytetrafluoroethylene), and in halons such as freon.

• Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles.

• Sodium hexafluoroaluminate (cryolite), is used in the electrolysis of aluminium.

• In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches.

• Fluorides have been used in the past to help molten metal flow, hence the name.

• Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse. The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive.

• Compounds of fluorine such as fluoropolymers, potassium fluoride and cryolite are utilized in applications such as anti-reflective coatings and dichroic mirrors on account of their unusually low refractive index.

 

IODINE

A halogen halogen ( (Group VII Group VII) that is a solid at room temperature) that is a solid at room temperature. Iodine is a weak oxidizer, forming Iodine is a weak oxidizer, forming iodides iodides with with hydrogen hydrogen and and metals metals.

CARBON

A non-metal from Group IV of the periodic table with atomic number 6. Carbon is found iature in two different forms or allotropes: diamond and graphite. Carbon is the element on which all life on Earth is based – all living tissues contain carbon compounds. Many carbon compounds burn in air, releasing heat and light, and so are attractive as fuels. Coal, coke, and charcoal are impure forms of carbon often used as fuel. Most carbon originates in stars that do not explode. After such stars become red giants, they convert helium into carbon. This seeps into the star’s outer atmosphere and gets ejected when the star casts off  its atmosphere to form a planetary nebula.

1. THE REACTIONS WITH ACIDS

Dilute acids have no effect on any form of carbon, and diamond is resistant to attack by concentrated acids at room temperature, but is oxidised by both concentrated sulphuric and concentrated nitric acid at about 500 K, when an additional oxidising agent is present. Carbon dioxide is produced and the acids are reduced to gaseous oxides:

C + 4HNO₃ → CO₂ + 2H₂O + 4NO₂

C + 2H₂SO₄ → CO₂ + 2H₂O + 2SO²

Graphite reacts rather differently with mixtures of oxidising agents and concentrated oxoacids. A “graphite oxide’ is formed; the graphite swells because oxygen atoms become attached to some of the carbon atoms in the rings and distend the layer structure. ‘Graphite oxide’ is rather indefinite in composition. With concentrated sulphuric acid and an oxidising agent a blue solution called ‘graphite hydrogen sulphate‘ is formed; this has an approximate formula (Cx) ⁺ HSO₄ ^–. 2H₂SO₄.

Amorphous carbon, having a far greater effective surface area than either diamond or graphite, is the most reactive form of carbon. It reacts with both hot concentrated sulphuric and hot concentrated nitric acids in the absence of additional oxidising agents but is not attacked by hydrochloric acid.

2. THE REACTIONS WITH ALKALIS

Carbon does not react, even with molten alkali.

3. THE REACTIONS WITH OXYGEN

All forms of carbon, if heated to a sufficiently high temperature, give carbon dioxide in a plentiful supply of air, and carbon monoxide if the supply is limited:

C + O₂ → CO₂; ∆H= – 394 kJ mol^-1

C + ½O₂ → CO; ∆H = – 111 kJ mol ^-1

 

NITROGEN

The element with atomic number 7. Nitrogen is the chief gas in Earth’s atmosphere. Nitrogen is vital to living things as it is an component of proteins and nucleic acids. Nearly all nitrogen is produced in stars that do not explode;  when these stars grow old, they expand into red giants and cast their atmospheres into space, forming planetary nebulae and so enriching the Universe with nitrogen.

CHEMICAL REACTIVITY OF NITROGEN

1. Reaction with air

The dissociation energy of the N ≡ N bond is very large. 946 kJ mol^-1 and dissociation of nitrogen molecules into atoms is not readily effected until very high temperatures, being only slight even at 3000 K. It is this high bond energy coupled with the absence of bond polarity that explains the low reactivity of nitrogen, in sharp contrast to other triple bond structures such as —C ≡ N, —C ≡ O, —C ≡ C—. Nitrogen does, however, combine with oxygen to a small extent when a mixture of the gases is subjected to high temperature or an electric discharge, the initial product being nitrogen monoxide, NO. The combination caused by an electric discharge can readily be shown in the laboratory.

2. Reaction with acids

Hydrochloric and dilute sulphuric acids have no appreciable action at room temperature on the pure Group V elements. Concentrated sulphuric acid and nitric acid—powerful oxidizing agents—attack all the elements except nitrogen, particularly when the acids are warm. The products obtained reflect changes in stability of the oxidation states V and III of the Group V elements. Both white and red phosphorus dissolve in, for example, concentrated nitric acid to form phosphoric(V) acid, the reaction between hot acid and white phosphorus being particularly violent. Arsenic dissolves in concentrated nitric acid forming arsenic(V) acid, H₃AsO₄, but in dilute nitric acid and concentrated sulphuric acid the main product is the arsenic(III) acid, H₃AsO₃. The more metallic element, antimony, dissolves to form the (III) oxide Sb₄O₆ with moderately concentrated nitric acid, but the (V) oxide Sb₂O₅ (structure unknown) with the more concentrated acid. Bismuth, however, forms the salt bismuth(Ill) nitrate Bi(NO₃)₃. 5H₂O.

3. Reaction with alkalis

The change from non-metallic to metallic properties of the Group V elements as the atomic mass of the element increases is shown in their reactions with alkalis. The head element nitrogen does not react.

4. Reaction with halogens

Nitrogen does form a number of binary compounds with the halogens but none of these can be prepared by the direct combination of the elements and they are dealt with below. The other Group V elements all form halides by direct combination.

 

OXYGEN

A non-metallic element with atomic number 8 (Group VI) of the periodic table. It is a vigorous oxidizing agent. Gaseous molecular oxygen is diatomic (O₂). Oxygen is the second most common element in the Earth System by weight. Oxygen is a component of water (H₂O) and, in gaseous form, of Earth’s atmosphere, and it is an essential element for plant and animal life (see respiration, photosynthesis for details). Oxygen is produced primarily by high-mass stars when they fuse  helium in oxygen (and carbon); this oxygen gets ejected into the Galaxy when the star explodes as supernovae.

Since oxygen has the second largest electronegativity among all the elements, it is found in the –2 oxidation state in most compounds. Important oxides have already been discussed in sections dealing with the elements from which they form, and so we will deal only with unusual oxidation states of oxygen here. One of these is the +2 state found in OF₂, the most common compound in which oxygen is combined with the more electronegative fluorine. We have already mentioned the –½ and –1 states observed in alkali-metal superoxides and peroxides, but one important peroxide, hydrogen peroxide (H₂O₂), has not yet been discussed.

H₂O₂ can be prepared by electrolysis of solutions containing sulfate ions. H₂O₂ is a weak acid, and it can serve as an oxidizing agent (oxygen being reduced to the –2 state) or as a reducing agent (oxygen being oxidized to the 0 state). Like the peroxide ion, the H₂O₂ molecule contains an O—O single bond. This bond is rather weak compared with many other single bonds, and this contributes to the reactivity of H₂O₂. The compound decomposes easily, especially if exposed to light or contaminated with traces of transition metals. The decomposition

2H₂O₂(l) → 2H₂O(l) + O₂(g)

can occur explosively in the case of the pure liquid.

 

HYDROGEN

The lightest and most common element in the Universe. It commonly exists as a diatomic gas, H₂. Hydrogen is reactive and highly inflammable, forming explosive mixtures with oxygen. Hydrogen forms co-valent bonds by sharing electrons. Hydrogen has three isotopes:

→ hydrogen-1, the most common isotope, contains one proton and no neutrons;

→ hydrogen-2 or deuterium, contains one proton and one neutron;

→ hydrogen-3 or tritium, contains one proton and two neutrons and is radioactive.

Main-sequence stars, such as the Sun, generate energy by fusing hydrogen into helium in their cores. Hydrogen readily fuses, even in most brown dwarfs.

Water

Function of Function of water water in the organism: in the organism:

• it is an universal solvent and transport medium it is an universal solvent and transport medium.

•it is a it is a structural component of biological macromolecules structural component of biological macromolecules.

•mediates energy transfer in ordered biological systems mediates energy transfer in ordered biological systems.

•it is an activator of certain chem. reactions it is an activator of certain chemical reactions.

•it takes part essentially in organism thermoregulation it takes part essentially in organism thermoregulation.

•as a as a basic factor it secures stability of internal environment of basic factor it secures stability of internal environment of  cells and organisms.

SELENIUM

-has important role in organism.

-Is in enzyme Glutation peroxidase (GSH) whic protects cells from toxic effects of hydrogen peroxide, which is converted into water:

2 GSH + H₂O₂ → GSSG + 2 H₂O

METALS

IRON

A transition metal with atomic number 26. It can have valency 2 or 3. Iron is the fourth most common element in the Earth System by weight; it makes up much of Earth’s core. Small amounts of iron are essential to human as it is a component of hemoglobin, the oxygen carrier in blood. All these elements are in enzymes.

COPPER

in ceruloplasmin, SOD, azurin, monoamine oxidase.

ZINC

•in oxidoreductases :alkoholdehydrogenase, SOD, transferases

•hydroxylases:alkaline phosphatase

•lyases

•isomerases

•ligases

MANGAN in SOD, arginase, pyruvat carboxylase

COBALT  in vitamin B12, ribonucleotide reductase.

MOLYBDEN  in xantine oxidase, nitrate reductase, formate dehydrogenase.

NICKEL  in Complexes, urease.

TOXICITY of ELEMENTS TOXICITY of ELEMENTS:

Cadmium, lead, Hg, Al, Bi, Tl, /harmfull effect in higher concentrations.

 

S – ELEMENTS. ALKALI METALS

Alkali metals are the chemical elements found in Group 1 of the periodic table. The alkali metals include: Lithium (Li), Sodium (Na),  Potassium (K), Rubidium (RB), Cesium (Cs), and Francium (Fr). Hydrogen, while it appears to be listed within Group 1, is not included in the alkali metals since it rarely exhibits similar behavior. The word “alkali” received its name from the Arabic word “al qali,” meaning “from ashes”. These particular elements were given the name “Alkali” because they react with water to form hydroxide ions, creating a basic solution (pH>7). Solutions that have a pH greater than 7 are called alkaline solutions.

Properties and Facts About Alkali Metals

Alkali metals are known for being some of the most reactive metals. This is due in part to their larger atomic radii and low ionization energies. They tend to donate their electrons in reactions and often have an oxidation state of +1. These metals are characterized as being extremely soft and silvery in color. They also have low boiling and melting points and are less dense than most elements. Li, Na, and K have the ability to float on water because of their low density. All of these characteristics can be attributed to the large atomic radii and weak metallic bonding these elements possess. Group 1 elements have a valence electron configuration is ns¹ and are good reducing agents (meaning they are easily oxidized). All of the alkali metals are found naturally iature, but not in their pure forms. Most combine with oxygen and silica to form minerals in the Earth and are readily mined as they are of relatively low densitys and thus do not sink.

Alkali Metal Reactions

REACTIONS WITH OXYGEN

The alkali metals tend to form ionic solids in which the alkali metal has an oxidatioumber of +1. Therefore, neutral compounds with oxygen can be readily classified according to the nature of the oxygen species involved. Ionic oxygen species include the oxide, O^2-, peroxide, O₂^2-, superoxide, O₂^–, andozonide O₃^–. Compounds that can be prepared that contain an alkali metal, M, and oxygen are therefore the monoxide, M₂O, peroxide, M₂O₂, superoxide, MO₂, and ozonide, MO₃. Rubidium and cesium and, possibly, potassium also form the sesquioxide, M₄O₆, which contains two peroxide anions and one superoxide anion per formula unit. Lithium forms only the monoxide and the peroxide.

All the alkali metals react directly with oxygen; lithium and sodium form monoxides, Li₂O and Na₂O, and the heavier alkali metals form superoxides, MO₂. The rate of reaction with oxygen, or with air, depends upon whether the metals are in the solid or liquid state, as well as upon the degree of mixing of the metals with the oxygen or air. In the liquid state, alkali metals can be ignited in air with ease, generating copious quantities of heat and a dense choking smoke of the oxide.

The free energy of formation (a measure of stability) of the alkali metal oxides at 25 °C (77 °F) varies widely from a high of −133 kcal/mole for lithium oxide to −63 kcal/mole for cesium oxide. The close approach of the small lithium ion to the oxygen atom results in the unusually high free energy of formation of the oxide. The peroxides (Li₂O₂and Na₂O₂) can be made by passing oxygen through a liquid-ammoniasolution of the alkali metal, although sodium peroxide is made commercially by oxidation of sodium monoxide with oxygen. Sodium superoxide (NaO₂) can be prepared with high oxygen pressures, whereas the superoxides of rubidium, potassium, and cesium can be prepared directly by combustion in air. By contrast, no superoxides have been isolated in pure form in the case of lithium or the alkaline-earth metals, although the heavier members of that group can be oxidized to the peroxide state. The cyanides of potassium, rubidium, and cesium, which are less stable than the lower oxides, can be prepared by the reaction of the superoxides with ozone.

REACTIONS WITH WATER

The alkali metals all react violently with water according to M + H₂O → MOH + 1/2 H₂. The rate of the reaction depends on the degree of metal surface presented to the liquid. With small metal droplets or thin films of alkali metal, the reaction can be explosive. The rate of the reaction of water with the alkali metals increases with increasing atomic weight of the metal. With the heavier alkali metals, the hydroxides are highly soluble; thus, they are removed readily from the reacting surface, and the reaction can proceed with unabated vigour. The reaction involves equimolar mixtures (that is, equal numbers of atoms or molecules) of the alkali metal and water to form a mole (an amount equal to that of the reactants) of alkali metal hydroxide and half a mole of hydrogen gas. These reactions are highly exothermic (give off heat), and the hydrogen that is generated can react with oxygen to increase further the heat that is generated.

REACTIONS WITH NONMETALS

Of the alkali metals, only lithium reacts with nitrogen, and it forms a nitride (Li3N). In this respect it is more similar to the alkaline-earth metals than to the Group 1 metals. Lithium also forms a relatively stablehydride, whereas the other alkali metals form hydrides that are more reactive. Lithium forms a carbide(Li₂C₂) similar to that of calcium. The other alkali metals do not form stable carbides, although they do react with the graphite form of carbon to give intercalation compounds (substances in which the metal atoms are inserted between layers of carbon atoms in the graphite structure).

The alkali metals can be burned in atmospheres of the various halogens to form the correspondinghalides. The reactions are highly exothermic, producing up to 235 kcal/mole for lithium fluoride. The alkali metals react with nonmetals in Groups 15 and 16 (Va and VIa) of the periodic table. Sulfides can be formed by the direct reaction of the alkali metals with elemental sulfur, furnishing a variety of sulfides.Phosphorus combines with the alkali metals to form phosphides with the general formula M₃P.

FORMATION OF ALLOYS

The characteristics of alloy behaviour in alkali metals can be evaluated in terms of the similarity of the elements participating in the alloy. Elements with similar atomic volumes form solid solutions (that is, mix completely in all proportions); some dissimilarity in atomic volumes results in eutectic-type systems (solutions formed over limited concentration ranges), and further dissimilarity results in totally immiscible systems. The high-pressure transition in potassium, rubidium, and cesium that converts these s-type metals to more transition metal-like d-type metals yields atomic volumes that are similar to those of many transition metals at the same pressure. This permits alloys or compounds to form between these alkali metals and such transition metals as nickel or iron.

The elements potassium, rubidium, and cesium, which have rather similar atomic volumes and ionization energies, form complete solid solutions and mixed crystals. Sodium, which is a significantly smaller atom than potassium and has a higher ionization energy, tends to form eutectic systems with potassium, rubidium, and cesium. Even greater dissimilarity exists in the atomic volumes of sodium and lithium, resulting in insolubilities of the liquid phases. The consolute temperature (the temperature at which the two liquids become completely miscible) increases on going from the lithium-sodium alloy system to the lithium-cesium system. Lithium and cesium can coexist as two separate liquid phases at temperatures up to at least 1,100 °C (2,000 °F).

There is only one example of solid miscibility in alkali–alkaline-earth-metal binaries—the lithium-magnesium system, in which the two elements are very similar. Sodium forms compounds only with barium in the alkaline-earth-metal series. The heavier alkali metals all tend to form immiscible liquid phases with the alkaline earth metals.

Several elements in Group 12 (IIb) of the periodic table (zinc, cadmium, and mercury) react with the alkali metals to form compounds. Mercury forms at least six compounds, commonly termed amalgams, with each of the five alkali metals, and with the exception of the amalgam with lithium, the highest melting point compound in each series has the formula MHg2. Lithium and sodium also form compounds with cadmium and zinc.

FORMATION OF COMPLEXES

Until the late 1960s there were few complexes of the alkali metal cations with organic molecules. Specialized biological molecules such as valinomycin were known to complex selectively the potassium cation K+ for transport across cell membranes, but synthetic ionophores (molecules that can form complexes with ions) were rare. All the alkali cations have a charge of +1 and, except for lithium, are chemically similar and rather inert. The only significant difference between one alkali cation and another is the size.

The synthesis of crown ethers by American chemist Charles J. Pedersen in 1967 provided size-selective cyclic molecules consisting of ether oxygens forming a ring or “crown” that could complex a cation of the right size to fit into the hole in the centre of the molecule. In some cases two crown ether molecules can encapsulate a cation in a “sandwich” fashion. For example, K⁺ just fits into the centre of an 18-crown-6 ring (18 atoms in the ring, 12 of which are carbon atoms and 6 are ether oxygen atoms) to form a 1:1 complex (that is, 1 cation:1 crown ether), K⁺(18C6). Cs⁺ is too large to fit into the ring but can be complexed on one side to form the Cs⁺(18C6) complex or can be sandwiched between two 18-crown-6 molecules to form the 1:2 complex, Cs⁺(18C6)₂. Thus, the selectivity of a crown ether for a particular cation depends on the ring size. Common crown ethers are 12-crown-4, 15-crown-5, and 18-crown-6. These molecules are selective for Li⁺, Na⁺, and K⁺, respectively.

Even greater affinity for alkali cations was achieved by the synthesis of cryptands by French chemist Jean-Marie Lehn in 1968 and spherands by American chemist Donald Cram in 1979. These are three-dimensional molecules with an internal cavity or crypt that can completely encapsulate the alkali cation. By synthesizing molecules with different cavity sizes, the selectivity for particular cations over those of the “wrong” size to fit in the cavity can be controlled. It should be noted, however, that these molecules are not rigid and that flexibility of the framework can alter the cavity size to accommodate alkali cations of different sizes, although with differences in the strength of complexation.

Since the initial syntheses of crown ethers and cryptands, thousands of complexants for cations of various sizes, charges, and geometries have been synthesized. This has led to an entirely new branch of chemistrycalled supramolecular chemistry.

Analytical chemistry of the alkali metals

Classical methods of separation and analysis of alkali metals are rather difficult and time consuming. Forlithium they include such procedures as selective extraction of lithium chloride into organic solvents and the detection of lithium with azo dyes that give highly sensitive colour reactions in alkaline solutions. A modification of the uranyl acetate test (the precipitation of an insoluble sodium salt with uranyl acetate) has been used as a standard test for the presence of sodium. The use of a cobaltinitrite solution permits separation of potassium from sodium by precipitation of the insoluble potassium salt. There are essentially no satisfactory analytical methods for rubidium and cesium based on the use of reagents in solution.

Classical methods of separation of the alkali metals have been largely supplanted by chromatographic elution. Strongly acidic cation-exchange resins and aqueous acidic solutions are used. Generally the affinity increases with atomic weight so that the ions are eluted in the order Fr+ > Cs+ > Rb+ > K+ > Na+ > Li+, which is the order of decreasing size of the hydrated ions. Ion-exchange resins that are specific for lithium have been developed. Macrocyclic compounds such as crown ethers and cryptands that are selective for particular alkali metal ions have been synthesized. They form cationic complexes that can be dissolved in organic solvents such as chloroform (CHCl₃) with counterions such as picrate (C₆H₂[NO₂]₃O-).

The characteristic flame colours of the alkali metals (red, yellow, violet, red, and blue for Li, Na, K, Rb, and Cs, respectively) are qualitative indicators of the modern analytical methods used to determine the concentrations of alkali-metal salts in aqueous solution. The intensities of the characteristic spectral lines in emission after excitation by a flame or ICP (inductively coupled plasma) give quantitative measures of the individual alkali metal concentration in the parts per million range or lower. Determination of one alkali metal in the presence of another, however, can result in interference, which can be reduced by using specially prepared standard solutions that contain known amounts of the interfering metals.

The analysis of the alkali-metal samples for the presence of nonmetallic elements, such as oxygen, carbon, hydrogen, and nitrogen, requires specialized techniques. The oxygen content of sodium and potassium samples can be determined by extraction of the free alkali metal with mercury, leaving behind mercury-insoluble oxides and carbonates, which can subsequently be analyzed by means of solution methods. The oxygen content of rubidium and cesium can be accurately determined by precise measurement of the freezing point of these two elements.

The carbon content of alkali metals can be analyzed by oxidation of the alkali metal in pure oxygen, followed by infrared measurement of the carbon dioxide generated during combustion. For the analysis of nitride in lithium, the nitride commonly is converted to ammonia, and the ammonia is measured by colorimetric analysis.

 

1) With Hydrogen: all alkali metals react with hydrogen to form hydrides

2K(l) + H₂(g) → 2KH(s)

2) With Water: Alkali metals and water react violently to form strong bases and hydrogen gas.

General Reaction: 2M(s) + 2H₂O → MOH(aq) + H₂(g)

where M=alkali metal

example: 2Na(s) + 2H₂O → 2NaOH(aq) + H₂(g)

Reactivity with water increases as you go down the group.

The explosive reaction of sodium with water. In this case, the exothermic reaction is enough to ignite the hydrogen gas that

3) With Halogens: Alkali metals and halogens combine to form ionic salts

General Reaction: M(s) + X(g)→ MX(s)

where M=alkali metal and X=halogen

example: Na⁺(s) + Cl^–(g) → NaCl (s)

4) With Nitrogen: only Lithium reacts with Nitrogen at room temperature

6Li(s) + N₂(g) → 2Li₃N(s)

5) With Oxygen: Alkali metals form multiple types of oxides, peroxides and superoxides when combined with oxygen:

·                    Oxide ion= O^2-

o         compounds generally look like M₂O

§          ex. Li₂O

·                    Sodium forms Peroxides

o        Peroxide Ion= O₂^2-

§          compounds generally look like M₂O₂

§                ex. Na₂O₂

·                    Potassium, Cesium, and Rubidium form superoxides

o        Superoxide ion=O₂^–

§           compounds generally look like MO₂

§                ex. KO₂

Trends

·                    Electronegativity and Ionization energy increase from LEFT TO RIGHT and BOTTOM TO TOP

o        Alkali metals have the lowest electronegativity and ionization energy

o        Francium is the least electronegative element.

·                    Atomic radius increases from RIGHT TO LEFT and TOP TO BOTTOM

o        Francium is the largest element

·                    Boiling points and melting points increase going BOTTOM TO TOP

o        Lithium has the highest boiling point and Francium has the lowest boiling point in Group 1.

Uses

·                    Sodium Vapor Lamps

·                    Atomic Clocks

·                    Table Salt

Flame Colors

All alkali metals have their own specific flame color. The colors are caused by the difference in energy among the valence shell of s and p orbitals, which corresponds to wavelengths of visible light. When the element is introduced into the flame, its outer electrons are excited and jump to a higher electron orbital. The electrons then fall and emit energy in the form of light. The different colors of light depend on how much energy or how far the electron falls back to a lower energy level. For this reason, they are often used in fireworks. Each alkali metal has a unique color and is easily identifiable.

                                 

 Lithium                       Sodium                 Potassium            

Elements of the Alkali Metal Group

Lithium

·                    named after the Greek word for stone (lithos)

·                    discovered in Sweden in 1817

·                    Atomic number: 3

·                    Atomic weight: 6.941

·                    the lightest and least dense of all alkali metals

·                    highly reactive

·                    a soft metal

·                    has a low ionization energy

·                    Electron configuration: [He]2s¹

·                    Often used in rechargeable batteries.

o        include those used in cell phones, camcorders, laptop computers, and cardiac pacemakers.

Sodium

·                    named after the Latin word for soda, Natria

·                    discovered in 1807

·                    Atomic number: 11

·                    Atomic weight: 22.9897

·                    soft silvery metal.

·                    extremely reactive metal

·                    Electron configuration: [Ne]3s¹

·                    used iuclear reactors because of its low boiling point.

·                    Sodium is reacted with chlorine to produce the ionic halide, NaCl

o        Sodium chloride is an important part of human diet

§          It is used during winter months to control the ice on the road.

Potassium

·                    named after the word Potash

o        Potash: means that Potassium is an element contained in the compound

·                    discovered in 1807

·                    Atomic number: 19

·                    Atomic Weight: 39.0983

·                    one of the most abundant elements in the earth’s crust

·                    oxidizes easily

·                    lavender flame color

·                    Electron configuration: [Ar]4s¹

·                    used mostly to produce chemicals, such as fertilizers for use in agriculture.

o        Potassium is an important nutrient needed for plant growth.

Rubidium

·                    named after the latin word for red, rubidius

·                    Atomic number: 37

·                    soft metal

·                    reddish flame color

·                    Electron configuration: [Kr] 5s¹.

·                    discovered in 1861

·                    known to have about 26 isotopes

·                    very large half life at an estimated 49 billion years

Cesium

·                    Atomic number: 55

·                    forms a strong base with water

·                    Atomic Weight: 132.91

·                    discovered in 1860

·                    often used as a catalyst in various hydrogenation organic reactions

·                    low melting point

·                    Electron configuration: [Xe]6s¹.

Francium

·                    discovered in 1939

·                    very radioactive

·                    hardly any Francium occurring naturally in the earth’s crust

·                    Atomic number: 87

·                    Electron configuration: [Rn]6s¹

·                    heaviest and most electropositive metal

·                    has the lowest boiling point

o        melts at low temperatures.

·                    most reactive of the alkali metals group

Problems

1.                Which alkali metal has a higher melting point, Sodium (Na) or Francium (Fr)? Explain.

2.                True or False. NH₃ is an ionic hydride.

3.                What is the electron configuration of Rubidium?

4.                Which alkali metals form superoxides?

5.                Complete and balance the following equation:   Li₂O₂ + H₂O → ?

6.                Which element is the most electronegative: Francium, Potassium, or Lithium?

7.                True or False: Rubidium has a very short half-life and decays quickly.

8.                True or False: All alkali metals react with Nitrogen.

9.                Balance the following equation: Li(s)+N₂(g)→ ?

10.           Compounds that generally look like M₂O₂ are formed with a metal and what kind of oxygen ion?.

Answers

Highlight to see the answers.

1.                Sodium has a higher boiling point because it has a larger atomic radius that Francium. Greater atomic radius means a bigger molecule thus having a higher boiling point.

2.                False. Group 1 and 2 form ionic hydrides. P block forms molecular hydride and Nitrogen is in P block.

3.                [Kr] 5s¹

4.                K, Rb and Cs.

5.                Li₂O₂ + H₂O

6.                Lithium

7.                False

8.                False, only Lithium reacts with Nitrogen

9.                6Li(s)+N₂(g)→2Li₃N(s)

10.           Peroxide Ion= O₂^2-

 

Hydrogen

Hydrogen is the chemical element with atomic number 1. It is represented by the symbol H. With an average atomic weight of 1.00794 u (1.007825 u for Hydrogen-1), hydrogen is the lightest and most abundant chemical element, constituting roughly 75 % of the Universe’s chemical elemental mass. Stars in the main sequence are mainly composed of hydrogen in its plasma state. Naturally occurring elemental hydrogen is relatively rare on Earth.

The most common isotope of hydrogen is protium (name rarely used, symbol 1H) with a single proton and no neutrons. In ionic compounds it can take a negative charge (an anion known as a hydride and written as H−), or as a positively charged species H+. The latter cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds always occur as more complex species. Hydrogen forms compounds with most elements and is present in water and most organic compounds. It plays a particularly important role in acid-base chemistry with many reactions exchanging protons between soluble molecules. As the simplest atom known, the hydrogen atom has been of theoretical use. For example, as the only neutral atom with an analytic solution to the Schrödinger equation, the study of the energetics and bonding of the hydrogen atom played a key role in the development of quantum mechanics.

Hydrogen gas (now known to be H₂) was first artificially produced in the early 16th century, via the mixing of metals with strong acids. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, and that it produces water when burned, a property which later gave it its name, which in Greek means “water-former.” At standard temperature and pressure, hydrogen is a colorless, odorless, nonmetallic, tasteless, highly combustible diatomic gas with the molecular formula H₂.

Industrial production is mainly from the steam reforming of natural gas, and less often from more energy-intensive hydrogen production methods like the electrolysis of water. Most hydrogen is employed near its production site, with the two largest uses being fossil fuel processing (e.g., hydrocracking) and ammonia production, mostly for the fertilizer market.

Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.

Properties

Combustion

Colorless gas with purple glow in its plasma state. Spectral lines of Hydrogen

 

The Space Shuttle Main Engine burns hydrogen with oxygen, producing a nearly invisible flame at full thrust.

Hydrogen gas (dihydrogen or molecular hydrogen) is highly flammable and will burn in air at a very wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion for hydrogen is −286 kJ/mol:

2 H₂(g) + O₂(g) → 2 H₂O(l) + 572 kJ (286 kJ/mol)

Hydrogen gas forms explosive mixtures with air in the concentration range 4–74% (volume per cent of hydrogen in air) and with chlorine in the range 5–95%. The mixtures spontaneously detonate by spark, heat or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C (932 °F). Pure hydrogen-oxygen flames emit ultraviolet light and are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle main engine compared to the highly visible plume of a Space Shuttle Solid Rocket Booster. The detection of a burning hydrogen leak may require a flame detector; such leaks can be very dangerous. The destruction of the Hindenburg airship was an infamous example of hydrogen combustion; the cause is debated, but the visible flames were the result of combustible materials in the ship’s skin. Because hydrogen is buoyant in air, hydrogen flames tend to ascend rapidly and cause less damage than hydrocarbon fires. Two-thirds of the Hindenburg passengers survived the fire, and many deaths were instead the result of falls or burning diesel fuel.

H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are also potentially dangerous acids.

Electron energy levels

Depiction of a hydrogen atom with size of central proton shown, and the atomic diameter shown as about twice the Bohr model radius (image not to scale).

The ground state energy level of the electron in a hydrogen atom is −13.6 eV, which is equivalent to an ultraviolet photon of roughly 92 nm wavelength.

The energy levels of hydrogen can be calculated fairly accurately using the Bohr model of the atom, which conceptualizes the electron as “orbiting” the proton in analogy to the Earth’s orbit of the Sun. However, the electromagnetic force attracts electrons and protons to one another, while planets and celestial objects are attracted to each other by gravity. Because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, and therefore only certain allowed energies.

A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton. The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all— an illustration of how different the “planetary orbit” conception of electron motion differs from reality.

Elemental molecular forms

First tracks observed in liquid hydrogen bubble chamber at the Bevatron

There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1 (½+½); in the parahydrogen form the spins are antiparallel and form a singlet with a molecular spin quantum number of 0 (½–½). At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form, also known as the “normal form”. The equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy than the para form, it is unstable and cannot be purified. At very low temperatures, the equilibrium state is composed almost exclusively of the para form. The liquid and gas phase thermal properties of pure parahydrogen differ significantly from those of the normal form because of differences in rotational heat capacities, as discussed more fully in Spin isomers of hydrogen. The ortho/para distinction also occurs in other hydrogen-containing molecules or functional groups, such as water and methylene, but is of little significance for their thermal properties.

The uncatalyzed interconversion between para and ortho H2 increases with increasing temperature; thus rapidly condensed H2 contains large quantities of the high-energy ortho form that converts to the para form very slowly. The ortho/para ratio in condensed H2 is an important consideration in the preparation and storage of liquid hydrogen: the conversion from ortho to para is exothermic and produces enough heat to evaporate some of the hydrogen liquid, leading to loss of liquefied material. Catalysts for the ortho-para interconversion, such as ferric oxide, activated carbon, platinized asbestos, rare earth metals, uranium compounds, chromic oxide, or some nickel compounds, are used during hydrogen cooling.

A molecular form called protonated molecular hydrogen, or H₃⁺, is found in the interstellar medium (ISM), where it is generated by ionization of molecular hydrogen from cosmic rays. It has also been observed in the upper atmosphere of the planet Jupiter. This molecule is relatively stable in the environment of outer space due to the low temperature and density. H₃⁺ is one of the most abundant ions in the Universe, and it plays a notable role in the chemistry of the interstellar medium. Neutral triatomic hydrogen H₃ can only exist in an excited form and is unstable.

Compounds

Covalent and organic compounds

While H₂ is not very reactive under standard conditions, it does form compounds with most elements. Millions of hydrocarbons are known, but they are not formed by the direct reaction of elementary hydrogen and carbon. Hydrogen can form compounds with elements that are more electronegative, such as halogens (e.g., F, Cl, Br, I); in these compounds hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or nitrogen, hydrogen can participate in a form of strong noncovalent bonding called hydrogen bonding, which is critical to the stability of many biological molecules. Hydrogen also forms compounds with less electronegative elements, such as the metals and metalloids, in which it takes on a partial negative charge. These compounds are often known as hydrides.

Hydrogen forms a vast array of compounds with carbon. Because of their general association with living things, these compounds came to be called organic compounds; the study of their properties is known as organic chemistry and their study in the context of living organisms is known as biochemistry. By some definitions, “organic” compounds are only required to contain carbon. However, most of them also contain hydrogen, and because it is the carbon-hydrogen bond which gives this class of compounds most of its particular chemical characteristics, carbon-hydrogen bonds are required in some definitions of the word “organic” in chemistry.

In inorganic chemistry, hydrides can also serve as bridging ligands that link two metal centers in a coordination complex. This function is particularly common in group 13 elements, especially in boranes (boron hydrides) and aluminium complexes, as well as in clustered carboranes.

Hydrides

Compounds of hydrogen are often called hydrides, a term that is used fairly loosely. The term “hydride” suggests that the H atom has acquired a negative or anionic character, denoted H⁻, and is used when hydrogen forms a compound with a more electropositive element. The existence of the hydride anion, suggested by Gilbert N. Lewis in 1916 for group I and II salt-like hydrides, was demonstrated by Moers in 1920 with the electrolysis of molten lithium hydride (LiH), that produced a stoichiometric quantity of hydrogen at the anode. For hydrides other than group I and II metals, the term is quite misleading, considering the low electronegativity of hydrogen. An exception in group II hydrides is BeH2, which is polymeric. In lithium aluminium hydride, the AlH−4 anion carries hydridic centers firmly attached to the Al(III). Although hydrides can be formed with almost all main-group elements, the number and combination of possible compounds varies widely; for example, there are over 100 binary borane hydrides known, but only one binary aluminium hydride. Binary indium hydride has not yet been identified, although larger complexes exist.

 

Protons and acids

Oxidation of hydrogen, in the sense of removing its electron, formally gives H⁺, containing no electrons and a nucleus which is usually composed of one proton. That is why H⁺ is often called a proton. This species is central to discussion of acids. Under the Bronsted-Lowry theory, acids are proton donors, while bases are proton acceptors.

A bare proton, H⁺, cannot exist in solution or in ionic crystals, because of its unstoppable attraction to other atoms or molecules with electrons. Except at the high temperatures associated with plasmas, such protons cannot be removed from the electron clouds of atoms and molecules, and will remain attached to them. However, the term ‘proton’ is sometimes used loosely and metaphorically to refer to positively charged or cationic hydrogen attached to other species in this fashion, and as such is denoted “H⁺” without any implication that any single protons exist freely as a species.

To avoid the implication of the naked “solvated proton” in solution, acidic aqueous solutions are sometimes considered to contain a less unlikely fictitious species, termed the “hydronium ion” (H₃O⁺). However, even in this case, such solvated hydrogen cations are thought more realistically physically to be organized into clusters that form species closer to H₉O+ 4. Other oxonium ions are found when water is in solution with other solvents.

Although exotic on earth, one of the most common ions in the universe is the H+3 ion, known as protonated molecular hydrogen or the triatomic hydrogen cation.

Isotopes

Hydrogen discharge (spectrum) tube

Deuterium discharge (spectrum) tube

Protium, the most common isotope of hydrogen, has one proton and one electron. Unique among all stable isotopes, it has no neutrons.

Hydrogen has three naturally occurring isotopes, denoted ¹H, ²H and ³H. Other, highly unstable nuclei (⁴H to ⁷H) have been synthesized in the laboratory but not observed iature.

·                    ¹H is the most common hydrogen isotope with an abundance of more than 99.98%. Because the nucleus of this isotope consists of only a single proton, it is given the descriptive but rarely used formal name protium.

·                    ²H, the other stable hydrogen isotope, is known as deuterium and contains one proton and one neutron in its nucleus. Essentially all deuterium in the universe is thought to have been produced at the time of the Big Bang, and has endured since that time. Deuterium is not radioactive, and does not represent a significant toxicity hazard. Water enriched in molecules that include deuterium instead of normal hydrogen is called heavy water. Deuterium and its compounds are used as a non-radioactive label in chemical experiments and in solvents for ¹H-NMR spectroscopy. Heavy water is used as a neutron moderator and coolant for nuclear reactors. Deuterium is also a potential fuel for commercial nuclear fusion.

·                    ³H is known as tritium and contains one proton and two neutrons in its nucleus. It is radioactive, decaying into helium-3 through beta decay with a half-life of 12.32 years. Small amounts of tritium occur naturally because of the interaction of cosmic rays with atmospheric gases; tritium has also been released during nuclear weapons tests. It is used iuclear fusion reactions,  as a tracer in isotope geochemistry, and specialized in self-powered lighting devices.  Tritium has also been used in chemical and biological labeling experiments as a radiolabel.

Hydrogen is the only element that has different names for its isotopes in common use today. (During the early study of radioactivity, various heavy radioactive isotopes were giveames, but such names are no longer used). The symbols D and T (instead of ²H and ³H) are sometimes used for deuterium and tritium, but the corresponding symbol P is already in use for phosphorus and thus is not available for protium. In its nomenclatural guidelines, the International Union of Pure and Applied Chemistry allows any of D, T, ²H, and ³H to be used, although ²H and ³H are preferred.

Production

H₂ is produced in chemistry and biology laboratories, often as a by-product of other reactions; in industry for the hydrogenation of unsaturated substrates; and iature as a means of expelling reducing equivalents in biochemical reactions.

Laboratory

In the laboratory, H₂ is usually prepared by the reaction of acids on metals such as zinc with Kipp’s apparatus.

Zn + 2 H⁺ → Zn²⁺ + H₂

Aluminium can also produce H₂ upon treatment with bases:

2 Al + 6 H₂O + 2 OH⁻ → 2 Al(OH)₄⁻ + 3 H₂

The electrolysis of water is a simple method of producing hydrogen. A low voltage current is run through the water, and gaseous oxygen forms at the anode while gaseous hydrogen forms at the cathode. Typically the cathode is made from platinum or another inert metal when producing hydrogen for storage. If, however, the gas is to be burnt on site, oxygen is desirable to assist the combustion, and so both electrodes would be made from inert metals. (Iron, for instance, would oxidize, and thus decrease the amount of oxygen given off.) The theoretical maximum efficiency (electricity used vs. energetic value of hydrogen produced) is between 80–94%.

2 H₂O(aq) → 2 H₂(g) + O₂(g)

In 2007, it was discovered that an alloy of aluminium and gallium in pellet form added to water could be used to generate hydrogen. The process also creates alumina, but the expensive gallium, which prevents the formation of an oxide skin on the pellets, can be re-used. This has important potential implications for a hydrogen economy, as hydrogen can be produced on-site and does not need to be transported.

Industrial

Hydrogen can be prepared in several different ways, but economically the most important processes involve removal of hydrogen from hydrocarbons. Commercial bulk hydrogen is usually produced by the steam reforming of natural gas. At high temperatures (1000–1400 K, 700–1100 °C or 1300–2000 °F), steam (water vapor) reacts with methane to yield carbon monoxide and H₂.

CH₄ + H₂O → CO + 3 H₂

This reaction is favored at low pressures but is nonetheless conducted at high pressures (2.0  MPa, 20 atm or 600 inHg). This is because high-pressure H2 is the most marketable product and Pressure Swing Adsorption (PSA) purification systems work better at higher pressures. The product mixture is known as “synthesis gas” because it is often used directly for the production of methanol and related compounds. Hydrocarbons other than methane can be used to produce synthesis gas with varying product ratios. One of the many complications to this highly optimized technology is the formation of coke or carbon:

CH₄ → C + 2 H₂

Consequently, steam reforming typically employs an excess of H2O. Additional hydrogen can be recovered from the steam by use of carbon monoxide through the water gas shift reaction, especially with an iron oxide catalyst. This reaction is also a common industrial source of carbon dioxide:

CO + H₂O → CO₂ + H₂

Other important methods for H₂ production include partial oxidation of hydrocarbons:

2 CH₄ + O₂ → 2 CO + 4 H₂

and the coal reaction, which can serve as a prelude to the shift reaction above:

C + H₂O → CO + H₂

Hydrogen is sometimes produced and consumed in the same industrial process, without being separated. In the Haber process for the production of ammonia, hydrogen is generated from natural gas. Electrolysis of brine to yield chlorine also produces hydrogen as a co-product.

Simple stuff

With him one electron, hydrogen forms a diatomic molecule only with electron configuration of the ground state (1S²).

Hydrogen molecules differ great durability and low polarizability, small size and low weight, respectively, and great mobility. Therefore, low melting point of  hydrogen (- 259,1 ° C) and boiling (- 252,6 ° C) in that it yields only helium. Therefore he is slightly soluble in water and organic solvents.

The disintegration of molecules into atoms is noticeable extent only at temperatures higher 2000 ° C:              H₂ ® 2H

Hydrogen and restorative and shows oxidative properties. Iormal conditions because of the strength of  the molecule it was relatively active and interacts directly only with fluorine. When heated, it interacts with many non-metals – chlorine, bromine, oxygen. Reducing properties of hydrogen are used for some simple substances with oxides and gallides:

CuO + H₂ = Cu + H₂O

CuCl₂ + H₂ = Cu + 2HCl

As oxidants hydrogen interacts with active metals:

2Na + H₂ = 2NaH

 

Compounds of  Hydrogen with oxidation state -1.

Depending on the nature of the element associated with it the hydrogen atoms in crystals can be polarized positively (the oxidation state +1) or negative (Os -1):

In addition the group of compounds, which link E – H is close to non-polar. These compounds in appropriate circumstances can be seen as derivative as H ⁺ and H^–derivatives.

In cases where hydrogen acts as an oxidant it is similar to the halogens: gallides hidrydy. But, forms similar H^–ion formation of H₂ molecules – a process endothermic (enthalpy of H^–∆Н₂₉₈=150kJ/mol). Therefore, by oxidizing hydrogen activity substantially inferior to halogens.

Ionic hydrides – white crystalline substance with high melting point, that is salts. Their fusion – electrolytes, electrolysis of molten hydrides hydrogen released at the anode.

By Ionic hydrides, chemical nature – basic compounds:

                       _{-1           +1}

            КН + НОН = КОН + Н₂

In covalent hydrides are less electronegative than the hydrogen itself, non-metallic elements (for example, hydrides of VH₃ and SiH₄).

For chemical nature is acidic compounds:

                 _{-1            +1}

  SiH₄ + 3HOH = H₂SiO₃ + 4H₂

Differences in the chemical nature of hydrides can be set easily by their behavior during hydrolysis. A characteristic feature is the allocation of hydrides hydrogen hydrolysis. The reaction takes place by redox-reduction mechanism. By extracting hydrogen hydrolysis is complete and irreversible. This basic form alkali hydrides and acid – acid.

Differences chemical nature of acid and basic hydrides clearly manifested in their interaction with each other, for example, under the scheme:

              LiH + BH₃ = Li [ BH₄ ]

           basic.  acidic. tetrahidrydoborat lithium

  

                                                    H H

                             _                      .. ..

                          H: + B: N ® [H: B: H] –

                                                     .. ..

                                                    H H

                                                       

As can be seen amphoteric compound aluminum hydride AlH₃, which depending on the reaction partner may act as a donor and electron pairs (basic compound) and as acceptor (acidic compound):

                  AlH₃ + 3BH₃ = Al (BH₄) ₃

                                          tetrahidrydoborat aluminum

                   KH + AlH₃ = K [AlH₄]

                                        tetrahidrydoalyuminat potassium

Standard electrode potential of about 1/2 H₂/H^– equal – 2.23 V. Accordingly, the H^– ion – one of the strongest reductants. Therefore, ionic, and complex hydrides – are strong reductants and are widely used for various syntheses, for hydrogen. Calcium hydride is used as a dehumidifier (residues – traces of moisture).

 

Compounds of Hydrogen with oxidation state  +1.

Positive polarization of hydrogen atoms is observed in its many compounds with covalent sphere mechanism: under normal conditions this gas (HCl, H₂S, NH₃), liquid (H₂O, HF, HNO₃), solids (H₃PO₄, H₂SiO₃). The properties of these compounds strongly depend on the nature of the item which directly bound hydrogen. In particular, compounds containing sphere mechanism FH, OH, NH, characterized by the formation of hydrogen sphere mechanism. As a result, HF, H2O, NH3 exhibit abnormally high melting point and boiling point compared to the same type of binary compounds of hydrogen, formed by other elements of the group. As a result, the ability to form hydrogen sphere mechanism and enter into donor-acceptor attraction liquid HF, H₂O, NH₃ is a good ionizing solvents.

Hydrogen atoms are also included in the composition of acid salts such as NaHS, NaHCO₃, NaHSO₄.

Hydrogen interacts with many non-metals. Depending on the activity of nonmetals reaction occurs at different speeds. Since hydrogen fluoride interacts with an explosion:

H₂ + F₂ = 2HF

In the dark and without heat reaction

H₂ + Cl₂ = 2HCl

 is slow, when the light is much faster, and in the presence of initiator (heating) – instantly and and with explosion.

With Br₂ and I₂ hydrogen reacts very slowly.

With other non-metals or hydrogen reacts at high t⁰ C and P.

S + H₂ = H₂S,

N₂ + 3H₂ = 2NH₃

Hydrogen is an active reductants. However, atomic hydrogen is more active than molecular, so all the characteristic reactions of  hydrogen with atomic hydrogen is more energetically.

Applications

Large quantities of H₂ are needed in the petroleum and chemical industries. The largest application of H₂ is for the processing (“upgrading”) of fossil fuels, and in the production of ammonia. The key consumers of H₂ in the petrochemical plant include hydrodealkylation, hydrodesulfurization, and hydrocracking. H₂ has several other important uses. H₂ is used as a hydrogenating agent, particularly in increasing the level of saturation of unsaturated fats and oils (found in items such as margarine), and in the production of methanol. It is similarly the source of hydrogen in the manufacture of hydrochloric acid. H₂ is also used as a reducing agent of metallic ores.

Hydrogen is highly soluble in many rare earth and transition metals and is soluble in both nanocrystalline and amorphous metals. Hydrogen solubility in metals is influenced by local distortions or impurities in the crystal lattice. These properties may be useful when hydrogen is purified by passage through hot palladium disks, but the gas’s high solubility is a metallurgical problem, contributing to embrittle of many metals, complicating the design of pipelines and storage tanks.

Apart from its use as a reactant, H2 has wide applications in physics and engineering. It is used as a shielding gas in welding methods such as atomic hydrogen welding. H2 is used as the rotor coolant in electrical generators at power stations, because it has the highest thermal conductivity of any gas. Liquid H2 is used in cryogenic research, including superconductivity studies. Because H2 is lighter than air, having a little more than 1⁄15 of the density of air, it was once widely used as a lifting gas in balloons and airships.

In more recent applications, hydrogen is used pure or mixed with nitrogen (sometimes called forming gas) as a tracer gas for minute leak detection. Applications can be found in the automotive, chemical, power generation, aerospace, and telecommunications industries. Hydrogen is an authorized food additive (E 949) that allows food package leak testing among other anti-oxidizing properties.

Hydrogen’s rarer isotopes also each have specific applications. Deuterium (hydrogen-2) is used iuclear fission applications as a moderator to slow neutrons, and iuclear fusion reactions. Deuterium compounds have applications in chemistry and biology in studies of reaction isotope effects. Tritium (hydrogen-3), produced iuclear reactors, is used in the production of hydrogen bombs, as an isotopic label in the biosciences, and as a radiation source in luminous paints.

Water

Water is a chemical substance with the chemical formula H2O. Its molecule contains one oxygen and two hydrogen atoms connected by covalent bonds. Water is a liquid at ambient conditions, but it often co-exists on Earth with its solid state, ice, and gaseous state, water vapor or steam.

Water covers 70.9% of the Earth’s surface, and is vital for all known forms of life. On Earth, it is found mostly in oceans and other large water bodies, with 1.6% of water below ground in aquifers and 0.001% in the air as vapor, clouds (formed of solid and liquid water particles suspended in air), and precipitation. Oceans hold 97% of surface water, glaciers and polar ice caps 2.4%, and other land surface water such as rivers, lakes and ponds 0.6%. A very small amount of the Earth’s water is contained within biological bodies and manufactured products.

Water on Earth moves continually through a cycle of evaporation or transpiration (evapotranspiration), precipitation, and runoff, usually reaching the sea. Over land, evaporation and transpiration contribute to the precipitation over land.

Clean drinking water is essential to humans and other lifeforms. Access to safe drinking water has improved steadily and substantially over the last decades in almost every part of the world. There is a clear correlation between access to safe water and GDP per capita. However, some observers have estimated that by 2025 more than half of the world population will be facing water-based vulnerability. A recent report (November 2009) suggests that by 2030, in some developing regions of the world, water demand will exceed supply by 50%. Water plays an important role in the world economy, as it functions as a solvent for a wide variety of chemical substances and facilitates industrial cooling and transportation. Approximately 70% of freshwater is consumed by agriculture.

Chemical and physical properties

Impact from a water drop causes an upward “rebound” jet surrounded by circular capillary waves.

Snowflakes by Wilson Bentley, 1902

 

Capillary action of water compared to mercury

Water is the chemical substance with chemical formula H₂O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom.

Water appears iature in all three common states of matter and may take many different forms on Earth: water vapor and clouds in the sky; seawater and icebergs in the polar oceans; glaciers and rivers in the mountains; and the liquid in aquifers in the ground.

At high temperatures and pressures, such as in the interior of giant planets, it is argued that water exists as ionic water in which the molecules break down into a soup of hydrogen and oxygen ions, and at even higher pressures as superionic water in which the oxygen crystallises but the hydrogen ions float around freely within the oxygen lattice.

The major chemical and physical properties of water are:

·          Water is a liquid at standard temperature and pressure. It is tasteless and odorless. The intrinsic color of water and ice is a very slight blue hue, although both appear colorless in small quantities. Water vapor is essentially invisible as a gas.

·          Water is transparent in the visible electromagnetic spectrum. Thus aquatic plants can live in water because sunlight can reach them. Ultra-violet and infrared light is strongly absorbed.

·          Since the water molecule is not linear and the oxygen atom has a higher electronegativity than hydrogen atoms, it carries a slight negative charge, whereas the hydrogen atoms are slightly positive. As a result, water is a polar molecule with an electrical dipole moment. Water also can form an unusually large number of intermolecular hydrogen bonds (four) for a molecule of its size. These factors lead to strong attractive forces between molecules of water, giving rise to water’s high surface tension  and capillary forces. The capillary action refers to the tendency of water to move up a narrow tube against the force of gravity. This property is relied upon by all vascular plants, such as trees.

·          Water is a good solvent and is often referred to as the universal solvent. Substances that dissolve in water, e.g., salts, sugars, acids, alkalis, and some gases – especially oxygen, carbon dioxide (carbonation) are known as hydrophilic (water-loving) substances, while those that do not mix well with water (e.g., fats and oils), are known as hydrophobic (water-fearing) substances.

·          All the major components in cells (proteins, DNA and polysaccharides) are also dissolved in water.

·          Pure water has a low electrical conductivity, but this increases significantly with the dissolution of a small amount of ionic material such as sodium chloride.

·          The boiling point of water (and all other liquids) is dependent on the barometric pressure. For example, on the top of Mt. Everest water boils at 68 °C (154 °F), compared to 100 °C (212 °F) at sea level. Conversely, water deep in the oceaear geothermal vents can reach temperatures of hundreds of degrees and remain liquid.

·          Water has the second highest molar specific heat capacity of any known substance, after ammonia, as well as a high heat of vaporization (40.65 kJ•mol−1), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth’s climate by buffering large fluctuations in temperature.

·     
The maximum density of water occurs at 3.98 °C (39.16 °F). It has the anomalous property of becoming less dense, not more, when it is cooled down to its solid form, ice. It expands to occupy 9% greater volume in this solid state, which accounts for the fact of ice floating on liquid water.

Model of hydrogen bonds between molecules of water

ADR label for transporting goods dangerously reactive with water:

·                        Water is miscible with many liquids, such as ethanol, in all proportions, forming a single homogeneous liquid. On the other hand, water and most oils are immiscible usually forming layers according to increasing density from the top. As a gas, water vapor is completely miscible with air.

·                        Water forms an azeotrope with many other solvents.

·                        Water can be split by electrolysis into hydrogen and oxygen.

·                        As an oxide of hydrogen, water is formed when hydrogen or hydrogen-containing compounds burn or react with oxygen or oxygen-containing compounds. Water is not a fuel, it is an end-product of the combustion of hydrogen. The energy required to split water into hydrogen and oxygen by electrolysis or any other means is greater than the energy that can be collected when the hydrogen and oxygen recombine.

·                        Elements which are more electropositive than hydrogen such as lithium, sodium, calcium, potassium and caesium displace hydrogen from water, forming hydroxides. Being a flammable gas, the hydrogen given off is dangerous and the reaction of water with the more electropositive of these elements may be violently explosive.

 

Water is the most common solvent

Water (H₂O) ‑ the most common substance on Earth. Natural water is never clean. Most rainwater is clean, but it contains a number of different minor impurities, which captures the air. Number of contaminants in fresh water is between 0,01 to 0,1%. Sea water contains 3.5% dissolved substances, the main mass of which is NaCl.

Water is the most abundant compound on Earth’s surface, covering about 70 percent of the planet. Iature, water exists in liquid, solid, and gaseous states. It is in dynamic equilibrium between the liquid and gas states at standard temperature and pressure. At room temperature, it is a tasteless and odorless liquid, nearly colorless with a hint of blue. Many substances dissolve in water and it is commonly referred to as the universal solvent. Because of this, water iature and in use is rarely pure and some of its properties may vary slightly from those of the pure substance. However, there are also many compounds that are essentially, if not completely, insoluble in water. Water is the only common substance found naturally in all three common states of matter and it is essential for all life on Earth. Water usually makes up 55% to 78% of the human body.

In keeping with the basic rules of chemical nomenclature, water would have a systematic name of dihydrogen monoxide, but this is not among the names published by the International Union of Pure and Applied Chemistry and, rather than being used in a chemical context, the name is almost exclusively used as a humorous way to refer to water.

Forms of water

Like many substances, water can take numerous forms that are broadly categorized by phase of matter. The liquid phase is the most common among water’s phases (within the Earth’s atmosphere and surface) and is the form that is generally denoted by the word “water.” The solid phase of water is known as ice and commonly takes the structure of hard, amalgamated crystals, such as ice cubes, or loosely accumulated granular crystals, like snow. For a list of the many different crystalline and amorphous forms of solid H₂O, see the article ice. The gaseous phase of water is known as water vapor (or steam), and is characterized by water assuming the configuration of a transparent cloud. (Note that the visible steam and clouds are, in fact, water in the liquid form as minute droplets suspended in the air.) The fourth state of water, that of a supercritical fluid, is much less common than the other three and only rarely occurs iature, in extremely uninhabitable conditions. When water achieves a specific critical temperature and a specific critical pressure (647 K and 22.064 MPa), liquid and gas phase merge to one homogeneous fluid phase, with properties of both gas and liquid. One example of naturally occurring supercritical water is found in the hottest parts of deep water hydrothermal vents, in which water is heated to the critical temperature by scalding volcanic plumes and achieves the critical pressure because of the crushing weight of the ocean at the extreme depths at which the vents are located. Additionally, anywhere there is volcanic activity below a depth of 2.25 km (1.40 mi) can be expected to have water in the supercritical phase.

Vienna Standard Mean Ocean Water is the current international standard for water isotopes. Naturally occurring water is almost completely composed of the neutron-less hydrogen isotope protium. Only 155 ppm include deuterium (₂H or D), a hydrogen isotope with one neutron, and fewer than 20 parts per quintillion include tritium (₃H or T), which has two.

Heavy water is water with a higher-than-average deuterium content, up to 100%. Chemically, it is similar but not identical to normal water. This is because the nucleus of deuterium is twice as heavy as protium, and this causes noticeable differences in bonding energies. Because water molecules exchange hydrogen atoms with one another, hydrogen deuterium oxide (DOH) is much more common in low-purity heavy water than pure dideuterium monoxide (D₂O). Humans are generally unaware of taste differences, but sometimes report a burning sensation or sweet flavor. Rats, however, are able to avoid heavy water by smell. Toxic to many animals, heavy water is used in the nuclear reactor industry to moderate (slow down) neutrons. Light water reactors are also common, where “light” simply designates normal water.

Light water more specifically refers to deuterium-depleted water (DDW), water in which the deuterium content has been reduced below the standard 155 ppm level.

Physics and chemistry of water

Water is the chemical substance with chemical formula H₂O: one molecule of water has two hydrogen atoms covalently bonded to a single oxygen atom. Water is a tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water vapor is essentially invisible as a gas.

Water is primarily a liquid under standard conditions, which is not predicted from its relationship to other analogous hydrides of the oxygen family in the periodic table, which are gases such as hydrogen sulfide. The elements surrounding oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, all combine with hydrogen to produce gases under standard conditions. The reason that water forms a liquid is that oxygen is more electronegative than all of these elements with the exception of fluorine. Oxygen attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the hydrogen atoms, and a net negative charge on the oxygen atom. The presence of a charge on each of these atoms gives each water molecule a net dipole moment. Electrical attraction between water molecules due to this dipole pulls individual molecules closer together, making it more difficult to separate the molecules and therefore raising the boiling point. This attraction is known as hydrogen bonding. The molecules of water are constantly moving in relation to each other, and the hydrogen bonds are continually breaking and reforming at timescales faster than 200 femtoseconds. However, this bond is sufficiently strong to create many of the peculiar properties of water, such as those that make it integral to life. Water can be described as a polar liquid that slightly dissociates disproportionately into the hydronium ion (H₃O+ (aq)) and an associated hydroxide ion (OH− (aq)).

2 H₂O (l) is in equilibrium with H₃O+ (aq) + OH− (aq)

The dissociation constant for this dissociation is commonly symbolized as Kw and has a value of about 10−14 at 25 °C; see “Water (data page)” and “Self-ionization of water” for more information.

Percentage of elements in water by mass: 11.1% hydrogen, 88.9% oxygen.

The self-diffusion coefficient of water is 2.299·10−9 m²·s−1

Water, ice and vapor

Heat capacity and heats of vaporization and fusion

Heat of vaporization of water from melting to critical temperature

Water has a very high specific heat capacity – the second highest among all the heteroatomic species (afterammonia), as well as a high heat of vaporization (40.65 kJ/mol or 2257 kJ/kg at the normal boiling point), both of which are a result of the extensive hydrogen bonding between its molecules. These two unusual properties allow water to moderate Earth’s climate by buffering large fluctuations in temperature. According to Josh Willis, ofNASA’s Jet Propulsion Laboratory, the oceans absorb one thousand times more heat than the atmosphere (air) and are holding 80 to 90% of the heat of global warming.

The specific enthalpy of fusion of water is 333.55 kJ/kg at 0 °C. Of common substances, only that of ammonia is higher. This property confers resistance to melting on the ice of glaciers and drift ice. Before and since the advent of mechanical refrigeration, ice was and still is in common use for retarding food spoilage.

Note that the specific heat capacity of ice at −10 °C is about 2.05 J/(g·K) and that the heat capacity of steam at 100 °C is about 2.080 J/(g·K).

An example of intermolecular hydrogen bonding in aself-assembled dimer complex reported by Meijer and coworkers. The hydrogen bonds are the dotted lines.

A hydrogen atom attached to a relatively electronegative atom is a hydrogen bond donor. This electronegative atom is usually fluorine,oxygen, or nitrogen. An electronegative atom such as fluorine, oxygen, or nitrogen is a hydrogen bond acceptor, whether it is bonded to a hydrogen atom or not. An example of a hydrogen bond donor is ethanol, which has a hydrogen bonded to oxygen; an example of a hydrogen bond acceptor which does not have a hydrogen atom bonded to it is the oxygen atom on diethyl ether.

A hydrogen attached to carbon can also participate in hydrogen bonding when the carbon atom is bound to electronegative atoms, as is the case in chloroform, CHCl₃. The electronegative atom attracts the electron cloud from around the hydrogeucleus and, by decentralizing the cloud, leaves the atom with a positive partial charge. Because of the small size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, represents a large charge density. A hydrogen bond results when this strong positive charge density attracts a lone pair of electrons on another heteroatom, which becomes the hydrogen-bond acceptor.

The hydrogen bond is often described as an electrostatic dipole-dipole interaction. However, it also has some features of covalent bonding: it is directional and strong, produces interatomic distances shorter than sum of van der Waals radii, and usually involves a limited number of interaction partners, which can be interpreted as a type of valence. These covalent features are more substantial when acceptors bind hydrogens from more electronegative donors.

The partially covalent nature of a hydrogen bond raises the following questions: “To which molecule or atom does the hydrogeucleusbelong?” and “Which should be labeled ‘donor’ and which ‘acceptor’?” Usually, this is simple to determine on the basis of interatomic distances in the X−H…Y system: X−H distance is typically ≈110 pm, whereas H…Y distance is ≈160 to 200 pm. Liquids that display hydrogen bonding are called associated liquids.

Hydrogen bonds can vary in strength from very weak (1–2 kJ mol⁻¹) to extremely strong (161.5 kJ mol⁻¹ in the ion HF−2). Typicalenthalpies in vapor include:

·             F−H^…:F (161.5 kJ/mol or 38.6 kcal/mol)

·             O−H^…:N (29 kJ/mol or 6.9 kcal/mol)

·             O−H^…:O (21 kJ/mol or 5.0 kcal/mol)

·             N−H^…:N (13 kJ/mol or 3.1 kcal/mol)

·             N−H^…:O (8 kJ/mol or 1.9 kcal/mol)

HO−H^…:OH+3 (18 kJ/mol or 4.3 kcal/mol; data obtained using molecular dynamics as detailed in the reference and should be compared to 7.9 kJ/mol for bulk water, obtained using the same molecular dynamics.)

Quantum chemical calculations of the relevant interresidue potential constants (compliance constants) revealed large differences between individual H bonds of the same type. For example, the central interresidue N−H•••N hydrogen bond between guanine and cytosine is much stronger in comparison to the N−H•••N bond between the adenine-thymine pair.

The length of hydrogen bonds depends on bond strength, temperature, and pressure. The bond strength itself is dependent on temperature, pressure, bond angle, and environment (usually characterized by local dielectric constant). The typical length of a hydrogen bond in water is 197 pm. The ideal bond angle depends on the nature of the hydrogen bond donor. The following hydrogen bond angles between a hydrofluoric acid donor and various acceptors have been determined experimentally:

The most ubiquitous and perhaps simplest example of a hydrogen bond is found between water molecules. In a discrete water molecule, there are two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them; the simplest case, when only two molecules are present, is called the water dimer and is often used as a model system. When more molecules are present, as is the case with liquid water, more bonds are possible because the oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with a hydrogen on another water molecule. This can repeat such that every water molecule is H-bonded with up to four other molecules, as shown in the figure (two through its two lone pairs, and two through its two hydrogen atoms). Hydrogen bonding strongly affects the crystal structure of ice, helping to create an open hexagonal lattice. The density of ice is less than the density of water at the same temperature; thus, the solid phase of water floats on the liquid, unlike most other substances.

Liquid water’s high boiling point is due to the high number of hydrogen bonds each molecule can form, relative to its low molecular mass. Owing to the difficulty of breaking these bonds, water has a very high boiling point, melting point, and viscosity compared to otherwise similar liquids not conjoined by hydrogen bonds. Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds of a water molecule is up to four. For example, hydrogen fluoride—which has three lone pairs on the F atom but only one H atom—can form only two bonds; (ammonia has the opposite problem: three hydrogen atoms but only one lone pair).

H−F^…H−F^…H−F

The exact number of hydrogen bonds formed by a molecule of liquid water fluctuates with time and depends on the temperature. The number of hydrogen bonds may also be affected by the presence of oxygen diffusion-enhancing compounds such as trans sodium crocetinate (TSC), which have been shown to encourage the formation of hydrogen bonds. From TIP4P liquid water simulations at 25 °C, it was estimated that each water molecule participates in an average of 3.59 hydrogen bonds. At 100 °C, this number decreases to 3.24 due to the increased molecular motion and decreased density, while at 0 °C, the average number of hydrogen bonds increases to 3.69. A more recent study found a much smaller number of hydrogen bonds: 2.357 at 25 °C. The differences may be due to the use of a different method for defining and counting the hydrogen bonds.

Where the bond strengths are more equivalent, one might instead find the atoms of two interacting water molecules partitioned into two polyatomic ions of opposite charge, specifically hydroxide(OH−) and hydronium (H₃O+). (Hydronium ions are also known as ‘hydroxonium’ ions.)

It can be that a single hydrogen atom participates in two hydrogen bonds, rather than one. This type of bonding is called “bifurcated” (split in two or ‘two-forked’). It can exist for instance in complex natural or synthetic organic molecules. It was suggested that a bifurcated hydrogen atom is an essential step in water reorientation.

Acceptor-type hydrogen bonds (terminating on an oxygen’s lone pairs) are more likely to form bifurcation (it is called overcoordinated oxygen, OCO) than are donor-type hydrogen bonds, beginning on the same oxygen’s hydrogens.

Many polymers are strengthened by hydrogen bonds in their main chains. Among the synthetic polymers, the best known example is nylon, where hydrogen bonds occur in the repeat unit and play a major role in crystallization of the material. The bonds occur between carbonyl and amine groups in the amide repeat unit. They effectively link adjacent chains to create crystals, which help reinforce the material. The effect is greatest in aramid fibre, where hydrogen bonds stabilize the linear chains laterally. The chain axes are aligned along the fibre axis, making the fibres extremely stiff and strong. Hydrogen bonds are also important in the structure of cellulose and derived polymers in its many different forms iature, such as wood and natural fibres such as cotton and flax.

The hydrogen bond networks make both natural and synthetic polymers sensitive to humidity levels in the atmosphere because water molecules can diffuse into the surface and disrupt the network. Some polymers are more sensitive than others. Thus nylons are more sensitive than aramids, and nylon 6 more sensitive thaylon-11.

·  Dramatically higher boiling points of NH₃, H₂O, and HF compared to the heavier analogues PH_3, H₂S, and HCl.

·  Increase in the melting point, boiling point, solubility, and viscosity of many compounds can be explained by the concept of hydrogen bonding.

·  Viscosity of anhydrous phosphoric acid and of glycerol

·  Dimer formation in carboxylic acids and hexamer formation in hydrogen fluoride, which occur even in the gas phase, resulting in gross deviations from the ideal gas law.

·  Pentamer formation of water and alcohols in apolar solvents.

·  High water solubility of many compounds such as ammonia is explained by hydrogen bonding with water molecules.

·  Negative azeotropy of mixtures of HF and water

·  Deliquescence of NaOH is caused in part by reaction of OH− with moisture to form hydrogen-bonded H₃O−2 species. An analogous process happens between NaNH₂ and NH₃, and between NaF and HF.

·  The fact that ice is less dense than liquid water is due to a crystal structure stabilized by hydrogen bonds.

The presence of hydrogen bonds can cause an anomaly in the normal succession of states of matter for certain mixtures of chemical compounds as temperature increases or decreases. These compounds can be liquid until a certain temperature, then solid even as the temperature increases, and finally liquid again as the temperature rises over the “anomaly interval”

Smart rubber utilizes hydrogen bonding as its sole means of bonding, so that it can “heal” when torn, because hydrogen bonding can occur on the fly between two surfaces of the same polymer.

Strength of nylon and cellulose fibres.

Wool, being a protein fibre is held together by hydrogen bonds, causing wool to recoil when stretched. However, washing at high temperatures can permanently break the hydrogen bonds and a garment may permanently lose its shape.

Electrical properties

Electrical conductivity

Pure water containing no exogenous ions is an excellent insulator, but not even “deionized” water is completely free of ions. Water undergoes auto-ionization in the liquid state, when two water molecules form one hydroxide anion (OH−) and one hydronium cation (H₃O+).

Because water is such a good solvent, it almost always has some solute dissolved in it, often a salt. If water has even a tiny amount of such an impurity, then it can conduct electricity far more readily.

It is known that the theoretical maximum electrical resistivity for water is approximately 182 kΩ•m at 25 °C. This figure agrees well with what is typically seen on reverse osmosis, ultra-filtered and deionized ultra-pure water systems used, for instance, in semiconductor manufacturing plants. A salt or acid contaminant level exceeding even 100 parts per trillion (ppt) in otherwise ultra-pure water begins to noticeably lower its resistivity by up to several kΩ•m.

In pure water, sensitive equipment can detect a very slight electrical conductivity of 0.055 µS/cm at 25 °C. Water can also be electrolyzed into oxygen and hydrogen gases but in the absence of dissolved ions this is a very slow process, as very little current is conducted. In ice, the primary charge carriers are protons (see proton conductor)

Electrolysis

Water can be split into its constituent elements, hydrogen and oxygen, by passing an electric current through it. This process is called electrolysis. Water molecules naturally dissociate into H+and OH− ions, which are attracted toward the cathode and anode, respectively. At the cathode, two H+ ions pick up electrons and form H₂ gas. At the anode, four OH− ions combine and releaseO2 gas, molecular water, and four electrons. The gases produced bubble to the surface, where they can be collected. The standard potential of the water electrolysis cell (when heat is added to the reaction) is a minimum of 1.23 V at 25 °C. The operating potential is actually 1.48 V (or above) in practical electrolysis when heat input is negligible.

One of the important properties of water is that it has a high dielectric constant. This constant shows its ability to make electrostatic bonds with other molecules, meaning it can eliminate the attraction of the opposite charges of the surrounding ions.

Polarity and hydrogen bonding

An important feature of water is its polar nature. The water molecule forms an angle, with hydrogen atoms at the tips and oxygen at the vertex. This angle formed is 104.3 degrees as opposed to the typical tetrahedral angle of 109 degrees. Because oxygen has a higher electronegativity than hydrogen, the side of the molecule with the oxygen atom has a partial negative charge. Also the presence of the lone pairs tend to push the oxygen away. An object with such a charge difference is called a dipole meaning two poles. The oxygen end is partially negative and the hydrogen end is partially positive, because of this the direction of the dipole moment points from the oxygen towards the center of the hydrogens. The charge differences cause water molecules to be attracted to each other (the relatively positive areas being attracted to the relatively negative areas) and to other polar molecules. This attraction contributes to hydrogen bonding, and explains many of the properties of water, such as solvent action.

A water molecule can form a maximum of four hydrogen bonds because it can accept two and donate two hydrogen atoms. Other molecules like hydrogen fluoride, ammonia, methanol form hydrogen bonds but they do not show anomalous behavior of thermodynamic, kinetic or structural properties like those observed in water. The answer to the apparent difference between water and other hydrogen bonding liquids lies in the fact that apart from water none of the hydrogen bonding molecules can form four hydrogen bonds, either due to an inability to donate/accept hydrogens or due to steric effects in bulky residues. In water, local tetrahedral order due to the four hydrogen bonds gives rise to an open structure and a 3-dimensional bonding network, resulting in the anomalous decrease of density when cooled below 4 °C.

Although hydrogen bonding is a relatively weak attraction compared to the covalent bonds within the water molecule itself, it is responsible for a number of water’s physical properties. One such property is its relatively high melting and boiling point temperatures; more energy is required to break the hydrogen bonds between molecules. The similar compound hydrogen sulfide (H₂S), which has much weaker hydrogen bonding, is a gas at room temperature even though it has twice the molecular mass of water. The extra bonding between water molecules also gives liquid water a large specific heat capacity. This high heat capacity makes water a good heat storage medium (coolant) and heat shield.

Cohesion and adhesion

Water molecules stay close to each other (cohesion), due to the collective action of hydrogen bonds between water molecules. These hydrogen bonds are constantly breaking, with new bonds being formed with different water molecules; but at any given time in a sample of liquid water, a large portion of the molecules are held together by such bonds.

Water also has high adhesion properties because of its polar nature. On extremely clean/smooth glass the water may form a thin film because the molecular forces between glass and water molecules (adhesive forces) are stronger than the cohesive forces. In biological cells and organelles, water is in contact with membrane and protein surfaces that are hydrophilic; that is, surfaces that have a strong attraction to water. Irving Langmuir observed a strong repulsive force between hydrophilic surfaces. To dehydrate hydrophilic surfaces—to remove the strongly held layers of water of hydration—requires doing substantial work against these forces, called hydration forces. These forces are very large but decrease rapidly over a nanometer or less. They are important in biology, particularly when cells are dehydrated by exposure to dry atmospheres or to extracellular freezing.

Surface tension

Water has a high surface tension of 72.8 mN/m at room temperature, caused by the strong cohesion between water molecules, the highest of the commoon-ionic, non-metallic liquids. This can be seen when small quantities of water are placed onto a sorption-free (non-adsorbent and non-absorbent) surface, such as polyethylene or Teflon, and the water stays together as drops. Just as significantly, air trapped in surface disturbances forms bubbles, which sometimes last long enough to transfer gas molecules to the water. Another surface tension effect is capillary waves, which are the surface ripples that form around the impacts of drops on water surfaces, and sometimes occur with strong subsurface currents flowing to the water surface. The apparent elasticity caused by surface tension drives the waves.

Capillary action

Due to an interplay of the forces of adhesion and surface tension, water exhibits capillary action whereby water rises into a narrow tube against the force of gravity. Water adheres to the inside wall of the tube and surface tension tends to straighten the surface causing a surface rise and more water is pulled up through cohesion. The process continues as the water flows up the tube until there is enough water such that gravity balances the adhesive force.

Surface tension and capillary action are important in biology. For example, when water is carried through xylem up stems in plants, the strong intermolecular attractions (cohesion) hold the water column together and adhesive properties maintain the water attachment to the xylem and prevent tension rupture caused by transpiration pull.

Water as a solvent

Water is also a good solvent, due to its polarity. Substances that will mix well and dissolve in water (e.g. salts) are known as hydrophilic (“water-loving”) substances, while those that do not mix well with water (e.g. fats and oils), are known as hydrophobic (“water-fearing”) substances. The ability of a substance to dissolve in water is determined by whether or not the substance can match or better the strong attractive forces that water molecules generate between other water molecules. If a substance has properties that do not allow it to overcome these strong intermolecular forces, the molecules are “pushed out” from the water, and do not dissolve. Contrary to the common misconception, water and hydrophobic substances do not “repel”, and the hydration of a hydrophobic surface is energetically, but not entropically, favorable.

When an ionic or polar compound enters water, it is surrounded by water molecules (Hydration). The relatively small size of water molecules typically allows many water molecules to surround one molecule of solute. The partially negative dipole ends of the water are attracted to positively charged components of the solute, and vice versa for the positive dipole ends.

In general, ionic and polar substances such as acids, alcohols, and salts are relatively soluble in water, and non-polar substances such as fats and oils are not. Non-polar molecules stay together in water because it is energetically more favorable for the water molecules to hydrogen bond to each other than to engage invan der Waals interactions with non-polar molecules.

An example of an ionic solute is table salt; the sodium chloride, NaCl, separates into Na+ cations and Cl− anions, each being surrounded by water molecules. The ions are then easily transported away from their crystalline lattice into solution. An example of a nonionic solute is table sugar. The water dipoles make hydrogen bonds with the polar regions of the sugar molecule (OH groups) and allow it to be carried away into solution.

Water in acid-base reactions

Chemically, water is amphoteric: it can act as either an acid or a base in chemical reactions. According to the Brønsted-Lowry definition, an acid is defined as a species which donates a proton (a H+ ion) in a reaction, and a base as one which receives a proton. When reacting with a stronger acid, water acts as a base; when reacting with a stronger base, it acts as an acid. For instance, water receives an H+ ion from HCl when hydrochloric acid is formed:

HCl (acid) + H2O (base) = H3O+ + Cl−

In the reaction with ammonia, NH3, water donates a H+ ion, and is thus acting as an acid:

NH3 (base) + H2O (acid) = NH+4 + OH−

Because the oxygen atom in water has two lone pairs, water often acts as a Lewis base, or electron pair donor, in reactions with Lewis acids, although it can also react with Lewis bases, forming hydrogen bonds between the electron pair donors and the hydrogen atoms of water. HSAB theory describes water as both a weak hard acid and a weak hard base, meaning that it reacts preferentially with other hard species:

H+ (Lewis acid) + H₂O (Lewis base) → H₃O⁺

Fe³⁺ (Lewis acid) + H₂O (Lewis base) → Fe(H₂O)³⁺

6Cl− (Lewis base) + H₂O (Lewis acid) → Cl(H₂O)⁻⁶

When a salt of a weak acid or of a weak base is dissolved in water, water can partially hydrolyze the salt, producing the corresponding base or acid, which gives aqueous solutions of soap andbaking soda their basic pH:

Na₂CO₃ + H₂O = NaOH + NaHCO₃

Ligand chemistry

Water’s Lewis base character makes it a common ligand in transition metal complexes, examples of which range from solvated ions, such as Fe(H₂O)³⁺, to perrhenic acid, which contains two water molecules coordinated to a rhenium atom, to various solid hydrates, such as CoCl₂·6H₂O. Water is typically a monodentate ligand, it forms only one bond with the central atom.

Organic chemistry

As a hard base, water reacts readily with organic carbocations, for example in hydration reaction, in which a hydroxyl group (OH−) and an acidic proton are added to the two carbon atoms bonded together in the carbon-carbon double bond, resulting in an alcohol. When addition of water to an organic molecule cleaves the molecule in two, hydrolysis is said to occur. Notable examples of hydrolysis are saponification of fats and digestion of proteins and polysaccharides. Water can also be a leaving group in S_N2 substitution and E2 elimination reactions, the latter is then known asdehydration reaction.

Acidity iature

Pure water has the concentration of hydroxide ions (OH−) equal to that of the hydronium (H₃O+) or hydrogen (H+) ions, which gives pH of 7 at 298 K. In practice, pure water is very difficult to produce. Water left exposed to air for any length of time will dissolve carbon dioxide, forming a dilute solution of carbonic acid, with a limiting pH of about 5.7. As cloud droplets form in the atmosphere and as raindrops fall through the air minor amounts of CO2 are absorbed, and thus most rain is slightly acidic. If high amounts of nitrogen and sulfur oxides are present in the air, they too will dissolve into the cloud and rain drops, producing acid rain.

Water in redox reactions

Water contains hydrogen in oxidation state +1 and oxygen in oxidation state −2. Because of that, water oxidizes chemicals with reduction potential below the potential of H+/H₂, such as hydrides,alkali and alkaline earth metals (except for beryllium), etc. Some other reactive metals, such as aluminum, are oxidized by water as well, but their oxides are not soluble, and the reaction stops because of passivation. Note, however, that rusting of iron is a reaction between iron and oxygen, dissolved in water, not between iron and water.

2 Na + 2 H₂O → 2 NaOH + H₂

Water itself can be oxidized, emitting oxygen gas, but very few oxidants react with water even if their reduction potential is greater than the potential of O₂/O₂−. Almost all such reactions require acatalyst.

4 AgF₂ + 2 H₂O → 4 AgF + 4 HF + O₂

Geochemistry

Action of water on rock over long periods of time typically leads to weathering and water erosion, physical processes that convert solid rocks and minerals into soil and sediment, but under some conditions chemical reactions with water occur as well, resulting in metasomatism or mineral hydration, a type of chemical alteration of a rock which produces clay minerals iature and also occurs when Portland cement hardens.

Water ice can form clathrate compounds, known as clathrate hydrates, with a variety of small molecules that can be embedded in its spacious crystal lattice. The most notable of these ismethane clathrate, 4CH₄·23H₂O, naturally found in large quantities on the ocean floor.

Transparency

Water is relatively transparent to visible light, near ultraviolet light, and far-red light, but it absorbs most ultraviolet light, infrared light, and microwaves. Most photoreceptors and photosynthetic pigments utilize the portion of the light spectrum that is transmitted well through water. Microwave ovens take advantage of water’s opacity to microwave radiation to heat the water inside of foods. The very weak onset of absorption in the red end of the visible spectrum lends water its intrinsic blue hue.

Heavy water and isotopologues

Several isotopes of both hydrogen and oxygen exist, giving rise to several known isotopologues of water.Hydrogen occurs naturally in three isotopes. The most common (¹H) accounting for more than 99.98% of hydrogen in water, consists of only a single proton in its nucleus. A second, stable isotope, deuterium (chemical symbol D or ²H), has an additional neutron. Deuterium oxide, D2O, is also known as heavy water because of its higher density. It is used iuclear reactors as aneutron moderator. The third isotope, tritium, has 1 proton and 2 neutrons, and is radioactive, decaying with a half-life of 4500 days. T2O exists iature only in minute quantities, being produced primarily via cosmic ray-induced nuclear reactions in the atmosphere. Water with one deuterium atom HDO occurs naturally in ordinary water in low concentrations (~0.03%) and D₂O in far lower amounts (0.000003%).

The most notable physical differences between H₂O and D₂O, other than the simple difference in specific mass, involve properties that are affected by hydrogen bonding, such as freezing and boiling, and other kinetic effects. The difference in boiling points allows the isotopologues to be separated. The self-diffusion coefficient of H₂O at 25°C is 23% higher than the value of D₂O.

Consumption of pure isolated D₂O may affect biochemical processes – ingestion of large amounts impairs kidney and central nervous system function. Small quantities can be consumed without any ill-effects, and even very large amounts of heavy water must be consumed for any toxicity to become apparent.

Oxygen also has three stable isotopes, with 16O present in 99.76%, 17O in 0.04%, and 18O in 0.2% of water molecules.

The first decomposition of water into hydrogen and oxygen, by electrolysis, was done in 1800 by an English chemist William Nicholson. In 1805, Joseph Louis Gay-Lussac and Alexander von Humboldt showed that water is composed of two parts hydrogen and one part oxygen.

Gilbert Newton Lewis isolated the first sample of pure heavy water in 1933.

The properties of water have historically been used to define various temperature scales. Notably, the Kelvin, Celsius, Rankine, and Fahrenheit scales were, or currently are, defined by the freezing and boiling points of water. The less common scales of Delisle, Newton, Réaumur and Rømer were defined similarly. The triple point of water is a more commonly used standard point today.

Systematic naming

The accepted IUPAC name of water is oxidane or simply water, or its equivalent in different languages, although there are other systematic names which can be used to describe the molecule.

The simplest systematic name of water is hydrogen oxide. This is analogous to related compounds such as hydrogen peroxide, hydrogen sulfide, and deuterium oxide (heavy water). Another systematic name, oxidane, is accepted by IUPAC as a parent name for the systematic naming of oxygen-based substituent groups, although even these commonly have other recommended names. For example, the name hydroxyl is recommended over oxidanyl for the –OH group. The name oxane is explicitly mentioned by the IUPAC as being unsuitable for this purpose, since it is already the name of a cyclic ether also known as tetrahydropyran.

The polarized form of the water molecule, H⁺OH⁻, is also called hydron hydroxide by IUPAC nomenclature.

Dihydrogen monoxide (DHMO) is a rarely used name of water. This term has been used in various hoaxes that call for this “lethal chemical” to be banned, such as in the dihydrogen monoxide hoax. Other systematic names for water include hydroxic acid, hydroxylic acid, and hydrogen hydroxide. Both acid and alkali names exist for water because it is amphoteric (able to react both as an acid or an alkali). None of these exotic names are used widely.

 

Natural water contains salts of Ca and Mg (Cl^–, SO₄^2-, HCO₃^–) is called strong.

Hardness of water is determined by the concentration of calcium and magnesium ions in mmol/liter.

Water hardness of less than 2mmol / l is called soft, 2 to 10 – average more than 10 – hard.

Hardness of water is divided into:

• temporary (carbonate);

• ongoing (non carbonate).

Temporary water hardness in water caused by the presence of hydrogen carbonate and permanent – chlorides, sulfates of calcium and magnesium.

When boiling water hydrocarbons decompose and temporary hardness is removed:

Са(НСО₃)₂ ® СаСО₃¯ + Н₂О + СО₂­

Carbonate hardness can be removed and added to it, calcium hydroxide or lime milk:

Са(НСО₃)₂ + Са(ОН)₂ = 2СаСО₃¯ + 2Н₂О

Standing water hardness in both cases remain. It can eliminate the addition of soda to the water:

Na₂CO₃ + CaCl₂ = 2NaCl + CaCO₃¯   or

CO₃^2- + Ca²⁺ = CaCO₃¯

CO₃^2- + Mg²⁺ = MgСО₃¯

   Permanent and temporary hardness also eliminated when water flow through cation changing.

All dissolved substances can be removed from water by distillation (distillation) or ion exchange.

Taste and odor

Water can dissolve many different substances, giving it varying tastes and odors. Humans and other animals have developed senses which enable them to evaluate the potability of water by avoiding water that is too salty or putrid. The taste of spring water and mineral water, often advertised in marketing of consumer products, derives from the minerals dissolved in it. However, pure H2O is tasteless and odorless. The advertised purity of spring and mineral water refers to absence of toxins, pollutants and microbes.

Hydrogen peroxide

Hydrogen peroxide (H₂O₂) is an oxidizer commonly used as a bleach. It is the simplest peroxide (a compound with an oxygen-oxygen single bond). Hydrogen peroxide is a clear liquid, slightly more viscous than water, that appears colorless in dilute solution. It is used as a disinfectant, antiseptic, oxidizer, and in rocketry as a propellant. The oxidizing capacity of hydrogen peroxide is so strong that it is considered a highly reactive oxygen species.

Hydrogen peroxide is naturally produced in organisms as a by-product of oxidative metabolism. Nearly all living things (specifically, all obligate and facultative aerobes) possess enzymes known as peroxidases, which harmlessly and catalytically decompose low concentrations of hydrogen peroxide to water and oxygen.

Structure and properties

H₂O₂ adopts a nonplanar structure of C2 symmetry. Although chiral, the molecule undergoes rapid racemization. The flat shape of the anti conformer would minimize steric repulsions, the 90° torsion angle of the syn conformer would optimize mixing between the filled p-type orbital of the oxygen (one of the lone pairs) and the LUMO of the vicinal O-H bond. The observed anticlinal “skewed” shape is a compromise between the two conformers.

Despite the fact that the O-O bond is a single bond, the molecule has a high barrier to complete rotation of 29.45 kJ/mol (compared with 12.5 kJ/mol for the rotational barrier of ethane). The increased barrier is attributed to repulsion between one lone pair and other lone pairs. The bond angles are affected by hydrogen bonding, which is relevant to the structural difference between gaseous and crystalline forms; indeed a wide range of values is seen in crystals containing molecular H₂O₂.

H₂ + O₂ → H₂O₂

Physical properties of hydrogen peroxide solutions

The properties of aqueous solutions of hydrogen peroxide differ from those of the neat material, reflecting the effects of hydrogen bonding between water and hydrogen peroxide molecules. Hydrogen peroxide and water form a eutectic mixture, exhibiting freezing-point depression. Whereas pure water melts and freezes at approximately 273K, and pure hydrogen peroxide just 0.4K below that, a 50% (by volume) solution melts and freezes at 221 K.

Decomposition

Hydrogen peroxide decomposes (disproportionates) exothermically into water and oxygen gas spontaneously:

2 H₂O₂ → 2 H₂O + O₂

Redox reactions

In acidic solutions, H₂O₂ is one of the most powerful oxidizers known—stronger than chlorine, chlorine dioxide, and potassium permanganate. Also, through catalysis, H2O2 can be converted into hydroxyl radicals (.OH), which are highly reactive.

In aqueous solutions, hydrogen peroxide can oxidize or reduce a variety of inorganic ions. When it acts as a reducing agent, oxygen gas is also produced.

In acidic solutions Fe²⁺ is oxidized to Fe³⁺ (hydrogen peroxide acting as an oxidizing agent),

2 Fe²⁺(aq) + H₂O₂ + 2 H⁺(aq) → 2 Fe³⁺(aq) + 2H₂O(l)

and sulfite (SO₃²⁻) is oxidized to sulfate (SO₄²⁻). However, potassium permanganate is reduced to Mn²⁺ by acidic H₂O₂. Under alkaline conditions, however, some of these reactions reverse; for example, Mn²⁺ is oxidized to Mn⁴⁺ (as MnO₂).

Other examples of hydrogen peroxide’s action as a reducing agent are reaction with sodium hypochlorite or potassium permanganate, which is a convenient method for preparing oxygen in the laboratory.

NaOCl + H₂O₂ → O₂ + NaCl + H₂O

2 KMnO₄ + 3 H₂O₂ → 2 MnO₂ + 2 KOH + 2 H₂O + 3 O₂

Exercise (Classifying coumpounds):

1. Which of the following substances is made of molecules?

NaBr

KBr

BaBr2

ClBr

CaBr2

2. Which of the following substances must be covalent?

KBr

Cs2O

F2

NaF

CaCl2

3. Which of the following is NOT a characteristic of ionic compounds?

	solid forms conduct electricity poorly		liquid forms conduct electricity well
	always contain both cations and anions		usually made of metals and nonmetals
	electrons are shared between atoms

4. Which of the following compounds contains a polyatomic ion?

CO2

Na2SO4

LiCl

C6H12O6

KBr

5. Which of the following is an inorganic compound?

H2CO2

NaC2O4

HC2H3O2

NaCN

C6H12O6

6. Which of the following is a binary covalent compound?

Na2O

NO2

C2H5O

NaNO3

CaO

Exercise (Polyatomic atoms)

1. The SCN^– ion forms a bright red complex with iron(II) ions. The SCN^– ion is called the:

thiocyanate ion

thiocyanite ion

sulfurocyanide ion

cyanate ion

cyanide ion

thiocyanide ion

2. Sodium bicarbonate is “baking soda”, the active ingredient in many antacids. The formula of the bicarbonate ion is:

HCO32-

HCO3–

H2CO3–

(CO3)24-

CO32-

3. If TeO₄^2- is the tellurate ion, the name of the TeO₃^2- ion is expected to be:

hypotellurite

pertellurate

hemitellurate

telluride

tellurite

4. Periodate ion is used in staining procedures that improve the quality of microscopic images. The formula of the periodate ion is:

IO4–

HIO3–

IO4-2

IO32-

IO3–

5. Veterinarians administer an 0.2 percent potassium permanganate solution to cows and horses as an antidote for plant alkaloid poisoning. A bottle that contains potassium permanganate solution might be labelled:

KMgO4

K2MnO4

K2MnO3

KMnO4

KMnO3

6. Ammonium nitrate is used as a fertilizer by farmers and as an explosive by terrorists. The formula of ammonium nitrate is

NH4NO3

NH3NO3

(NH4)2NO2

NH3NO2

AmNO2

7. If the selenate ion has formula SeO₄^2-, the selenite ion might be:

SeO4+2

SeO32-

SeO4–

SeO52-

SeO43-

8. Stable compounds containing the iodite ion have never been prepared. The formula of the iodite ion could be:

IO2-

HIO3–

IO2-2

IO2–

IO3–

9. Iodate ion is used as a disinfectant. The formula of the iodate ion is

IO32-

IO3–

IO4–

HIO3–

IO4-2

10. Phosphate ion is the most common:

-2 anion

-3 anion

-1 anion

+1 cation

+1 anion

Empirical and Molecular Formulas

The empirical formula of a compound gives the simplest whole number ratio of different types of atoms in the compound.  All salt formulas are empirical formulas.  On the other hand, the molecular formula of a compound may or may not be the same as its empirical formula.  For example, the molecular formula of butane is C4H10 while its empirical formula is C2H5.  The molecular formula gives the true number of each kind of atom in a molecule.

Empirical formulas may be easily determined from experimental data. 

Usually you must first determine how many grams of each type of atom are in the compound.  If percent composition data is given, assume that you have 100.0 g of the compound; then the number of grams of each element is equal to the percentage for that element.

The next task is convert the grams of each element to moles of the element.  Be sure to keep at least three significant figures in your answers.

The final step is to write the molar amounts of each element as subscripts in the formula.  Then divide all molar subscripts by the smallest value in the set.  At this point, the subscripts may all be very close to whole numbers; if so, you are finished. If one (or more) of the subscripts is not close to a whole number, multiply all molar subscripts by the simple factor which makes all subscripts whole numbers.

Once the empirical formula is determined, the molecular formula is easily found if the molar mass (molecular weight) of the molecule is also known.  You first calculate the molar mass of the empirical formula.  Then you divide the molar mass of the molecule by the molar mass of the empirical formula.  The division should give a simple whole number.  That number is the factor by which all subscripts in the empirical formula must be multiplied to obtain the molecular formula.

Exercises

1.  The molecular formula of the antifreeze ethylene glycol is C2H6O2.  What is the empirical formula?

2.  A well-known reagent in analytical chemistry, dimethylglyoxime, has the empirical formula C2H4NO.  If its molar mass is 116.1 g/mol, what is the molecular formula of the compound?

3.  Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy.  One of them is a blue solid that comes apart, reversibly, in the gas phase.  It contains 36.84% N.  What is the empirical formula of this oxide?

4.  A sample of indium chloride weighing 0.5000 g is found to contain 0.2404 g of chlorine  What is the empirical formula of the indium compound?

Answers:

1. CH3O

2.  Molar mass of empirical formula is 58.06 g/mol.  Thus molecular formula is C4H8N2O2.

3.  The ratios are .  Since 1.50 is not close to a whole number, we multiply both subscripts by 2.  The empirical formula is thus N2O3.  (The name is dinitrogen trioxide.)

4.  InCl3.

Naming Inorganic Compounds: Answer Sheet

1. Complete the following chart of corresponding ioames and formulas.

2. Complete the following chart of corresponding compound names and formulas. Circle the names of all non-ionic (i.e., molecular) compounds. (NOTE: In these answers I have used shading (instead of circles) to indicate the molecular compounds (diphosphorus pentoxide, etc.).)

Naming Inorganic Compounds

To name a compound you must first decide whether the substance is an ionic or molecular compound. Ionic compounds are easily recognized since they usually contain both metallic and non-metallic elements. The most common exception to this rule are ionic compounds containing the ammonium ion, NH4+, such as (NH₄)₂CO₃ or NH₄Br which contaio metal ions. Molecular compounds typically contain only non-metallic atoms (and metalloids).

Conventions for naming ionic compounds are given in Chang, pgs 53-56. To successfully follow the rules, however, you must be first learn the names of common ions (Chang, Tables 2.2 and 2.3, pg 55). Names of ionic compounds do not give the number of each type of ion in the formula: the chemist is supposed to be able to figure that out from his/her knowledge of ion charges and the requirement that salts be neutral (and thus have a sum of zero for the ion charges in the formula).

Binary compounds of the non-metals are named following the guidelines given in Chang on pages 56-58. Note that wheaming these molecular compounds, the number of atoms of a given type is commonly indicated with a prefix (di-, tri-, tetra, etc.).

Exercises

To name a compound you must first decide whether the substance is an ionic or molecular compound. Ionic compounds are easily recognized since they usually contain both metallic and non-metallic elements. The most common exception to this rule are ionic compounds containing the ammonium ion, NH₄⁺, such as (NH₄)2CO3 or NH₄Br which contaio metal ions. Molecular compounds typically contain only non-metallic atoms (and metalloids).

Conventions for naming ionic compounds are given in Chang, pp. 59-62. To successfully follow the rules, however, you must be first learn the names of common ions (Chang, Tables 2.2 and 2.3, p. 60). Names of ionic compounds do not give the number of each type of ion in the formula: the chemist is supposed to be able to figure that out from his/her knowledge of ion charges and the requirement that salts be neutral (and thus have a sum of zero for the ion charges in the formula).

Binary compounds of the non-metals are named following the guidelines given in Chang on pp. 62-64. Note that wheaming these molecular compounds, the number of atoms of a given type is commonly indicated with a prefix (di-, tri-, tetra, etc.).

1. Complete the following chart of corresponding ioames and formulas.

 

2. Complete the following chart of corresponding compound names and formulas. Circle the names of all non-ionic (i.e., molecular) compounds.

 

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

The following website shows the explosive reaction of alkali metals and water. It’s cool stuff! Check it out!

9. http://video.google.com/videosearch?hl=en&q=alkali+metals&um=1&ie=UTF-8&sa=X&oi=video_result_group&resnum=4&ct=title

10. http://video.google.com/videoplay?do…66654801392897

 

Prepared by PhD Falfushynska H.