The materials to prepare students for practical lessons of inorganic chemistry

LESSON № 15.

Themes:

1. s-elements I group of the periodic system and property of their  compounds.

2. s-elements ІІ group of the periodic system and property of their compounds. 

 

Plan

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 Metal Reactions

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 oxidation number 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.

All the alkali metals react vigorously or explosively with cold water, producing an aqueous solution of the strongly basic alkali metal hydroxide and releasing hydrogen gas. This reaction becomes more vigorous going down the group: lithium reacts steadily with effervescence, but sodium and potassium can ignite and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers. When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing hydrogen gas; this takes place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water).

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)

note: 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

 

 

PROPERTIES OF ALKALI METAL

The alkali metals have very similar properties: they are all shiny, soft, highly reactive metals at standard temperature and pressure and readily lose their outermost electron to form cations with charge +1. They can all be cut easily with a knife due to their softness, exposing a shiny surface that tarnishes rapidly in air due to oxidation. Because of their high reactivity, they must be stored under oil to prevent reaction with air, and are found naturally only in salts and never as the free element. In the modern IUPAC nomenclature, the alkali metals comprise the group 1 elements, excluding hydrogen (H), which is nominally a group 1 element but not normally considered to be an alkali metal as it rarely exhibits behaviour comparable to that of the alkali metals. All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.

All the discovered alkali metals occur iature: in order of abundance, sodium is the most abundant, followed by potassium, lithium, rubidium, caesium, and finally francium, which is very rare due to its extremely high radioactivity and thus occurs only in traces due to its presence iatural decay chains. Experiments have been conducted to attempt the synthesis of ununennium (Uue), which is likely to be the next member of the group, but they have all met with failure.  However, ununennium may not be an alkali metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements; even if it does turn out to be an alkali metal, it is predicted to have some differences in physical and chemical properties from its lighter homologues.

 Most alkali metals have many different applications. Two of the most well-known applications of the pure elements are rubidium and caesium atomic clocks, of which caesium atomic clocks are the most accurate and precise representation of time. A common application of the compounds of sodium is the sodium-vapour lamp, which emits very efficient light Table salt, or sodium chloride, has been used since antiquity. Sodium and potassium are also essential elements, having major biological roles as electrolytes, and although the other alkali metals are not essential, they also have various effects on the body, both beneficial and harmful.

Characteristics

Chemical

Like other groups, the known members of this family show patterns in electronic configuration, especially the outermost shells, resulting in trends in chemical behavior:

 

 

Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not well established due to its extreme radioactivity; thus, the presentation of its properties here is limited.

All the alkali metals are highly reactive and are never found in elemental forms iature. Because of this, they are usually stored in mineral oil or kerosene (paraffin oil). They react aggressively with the halogens to form the alkali metal halides, which are white ionic crystalline compounds that are all soluble in water except lithium fluoride (LiF). The alkali metals also react with water to form strongly alkaline hydroxides and thus should be handled with great care. The heavier alkali metals react more vigorously than the lighter ones; for example, when dropped into water, caesium produces a larger explosion than potassium. The alkali metals have the lowest first ionisation energies in their respective periods of the periodic table because of their low effective nuclear charge and the ability to attain a noble gas configuration by losing just one electron. The second ionisation energy of all of the alkali metals is very high as it is in a full shell that is also closer to the nucleus; thus, they almost always lose a single electron, forming cations. The alkalides are an exception: they are unstable compounds which contain alkali metals in a −1 oxidation state, which is very unusual as before the discovery of the alkalides, the alkali metals were not expected to be able to form anions and were thought to be able to appear in salts only as cations. The alkalide anions have filled s-subshells, which gives them more stability and allows them to exist. All the stable alkali metals except lithium are known to be able to form alkalides, and the alkalides have much theoretical interest due to their unusual stoichiometry and low ionisation potentials. Alkalides are chemically similar to the electrides, which are salts with trapped electrons acting as anions. A particularly striking example of an alkalide is “inverse sodium hydride”, H+Na−, as opposed to the usual sodium hydride, Na+H−: it is unstable in isolation, due to its high energy resulting from the displacement of two electrons from hydrogen to sodium, although several derivatives are predicted to be metastable or stable.

The chemistry of lithium shows several differences from that of the rest of the group as the small Li+ cation polarises anions and gives its compounds a more covalent character. Lithium and magnesium have a diagonal relationship: because of this, lithium has some similarities to magnesium. For example, lithium forms a stable nitride, a property common among all the alkaline earth metals (magnesium’s group) but unique among the alkali metals. In addition, among their respective groups, only lithium and magnesium form covalent organometallic compounds (e.g. LiMe and MgMe₂). Lithium fluoride is the only alkali metal halide that is not soluble in water, and lithium hydroxide is the only alkali metal hydroxide that is not deliquescent. Francium is also predicted show some differences due to its high atomic weight, causing its electrons to travel at considerable fractions of the speed of light and thus making relativistic effects more prominent. In contrast to the trend of decreasing electronegativities and ionisation energies of the alkali metals, francium’s electronegativity and ionisation energy are predicted to be higher than caesium’s due to the relativistic stabilisation of the 7s electrons; also, its atomic radius is expected to be abnormally low.

Compounds and reactions

Reaction with the group 14 elements

Lithium and sodium react with carbon to form acetylides, Li₂C₂ and Na₂C₂, which can also be obtained by reaction of the metal with acetylene. Potassium, rubidium, and caesium react with graphite; their atoms are intercalated between the hexagonal graphite layers, forming graphite intercalation compounds of formulae MC60 (dark grey, almost black), MC48 (dark grey, almost black), MC36 (blue), MC24 (steel blue), and MC8 (bronze) (M = K, Rb, or Cs). These compounds are over 200 times more electrically conductive than pure graphite, suggesting that the valence electron of the alkali metal is transferred to the graphite layers (e.g. M+C−8). Upon heating of KC8, the elimination of potassium atoms results in the conversion in sequence to KC24, KC36, KC48 and finally KC60. KC8 is a very strong reducing agent and is pyrophoric and explodes on contact with water. While the large alkali metals (K, Rb, and Cs) initially form MC8, the smaller ones initially form MC6.

When the alkali metals react with the heavier elements in the carbon group, ionic substances with cage-like structures are formed, such as the silicide M4Si4 (M = K, Rb, or Cs), which contains M+ and tetrahedral Si₄−4 ions. The chemistry of alkali metal germanides, involving the germanide ion Ge4− and other cluster (Zintl) ions such as Ge₂−4, Ge₄−9, Ge₂−9, and [(Ge₉)₂]6−, is largely analogous to that of the corresponding silicides. Alkali metal stannides are mostly ionic, sometimes with the stannide ion (Sn4−), and sometimes with more complex Zintl ions such as Sn4−9, which appears in tetrapotassium nonastannide (K4Sn9). The monatomic plumbide ion (Pb₄−) is unknown, and indeed its formation is predicted to be energetically unfavourable; alkali metal plumbides have complex Zintl ions, such as Pb₄–9.

Reaction with the pnictogens (alkali metal pnictides)

Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with nitrogen at standard conditions, and its nitride is the only stable alkali metal nitride. Nitrogen is an unreactive gas because breaking the strong triple bond in the dinitrogen molecule (N₂) requires a lot of energy. The formation of an alkali metal nitride would consume the ionisation energy of the alkali metal (forming M+ ions), the energy required to break the triple bond in N2 and the formation of N3− ions, and all the energy released from the formation of an alkali metal nitride is from the lattice energy of the alkali metal nitride. The lattice energy is maximised with small, highly charged ions; the alkali metals do not form highly charged ions, only forming ions with a charge of +1, so only lithium, the smallest alkali metal, can release enough lattice energy to make the reaction with nitrogen exothermic, forming lithium nitride. The reactions of the other alkali metals with nitrogen would not release enough lattice energy and would thus be endothermic, so they do not form nitrides at standard conditions. (Sodium nitride (Na₃N) and potassium nitride (K₃N), while existing, are extremely unstable, being prone to decomposing back into their constituent elements, and cannot be produced by reacting the elements with each other at standard conditions.)

All the alkali metals react readily with phosphorus and arsenic to form phosphides and arsenides with the formula M3Pn (where M represents an alkali metal and Pn represents a pnictogen). This is due to the greater size of the P3− and As3− ions, so that less lattice energy needs to be released for the salts to form. These are not the only phosphides and arsenides of the alkali metals: for example, potassium has nine different known phosphides, with formulae K₃P, K₄P₃, K₅P₄, KP, K₄P₆, K₃P₇, K₃P₁₁, KP_10.3, and KP₁₅. While most metals form arsenides, only the alkali and alkaline earth metals form mostly ionic arsenides. The structure of Na3As is complex with unusually short Na–Na distances of 328–330 pm which are shorter than in sodium metal, and this indicates that even with these electropositive metals the bonding cannot be straightforwardly ionic. Other alkali metal arsenides not conforming to the formula M₃As are known, such as LiAs, which has a metallic lustre and electrical conductivity indicating the presence of some metallic bonding. The antimonides are unstable and reactive as the Sb³− ion is a strong reducing agent; reaction of them with acids form the toxic and unstable gas stibine (SbH₃).  Bismuthides are not even wholly ionic; they are intermetallic compounds containing partially metallic and partially ionic bonds.

Reaction with the chalcogens (alkali metal chalcogenides)

All the alkali metals react vigorously with oxygen at standard conditions. They form various types of oxides, such as simple oxides (containing the O₂− ion), peroxides (containing the O₂²− ion, where there is a single bond between the two oxygen atoms), superoxides (containing the O₂− ion), and many others. Lithium burns in air to form lithium oxide, but sodium reacts with oxygen to form a mixture of sodium oxide and sodium peroxide. Potassium forms a mixture of potassium peroxide and potassium superoxide, while rubidium and caesium form the superoxide exclusively. Their reactivity increases going down the group: while lithium, sodium and potassium merely burn in air, rubidium and caesium are pyrophoric (spontaneously catch fire in air).

The smaller alkali metals tend to polarise the more complex anions (the peroxide and superoxide) due to their small size. This attracts the electrons in the more complex anions towards one of its constituent oxygen atoms, forming an oxide ion and an oxygen atom. This causes lithium to form the oxide exclusively on reaction with oxygen at room temperature. This effect becomes drastically weaker for the larger sodium and potassium, allowing them to form the less stable peroxides. Rubidium and caesium, at the bottom of the group, are so large that even the least stable superoxides can form. Because the superoxide releases the most energy when formed, the superoxide is preferentially formed for the larger alkali metals where the more complex anions are not polarised. (The oxides and peroxides for these alkali metals do exist, but do not form upon direct reaction of the metal with oxygen at standard conditions.) In addition, the small size of the Li+ and O₂− ions contributes to their forming a stable ionic lattice structure. Under controlled conditions, however, all the alkali metals, with the exception of francium, are known to form their oxides, peroxides, and superoxides. The alkali metal peroxides and superoxides are powerful oxidizing agents. Sodium peroxide and potassium superoxide react with carbon dioxide to form the alkali metal carbonate and oxygen gas, which allows them to be used in submarine air purifiers; the presence of water vapour, naturally present in breath, makes the removal of carbon dioxide by potassium superoxide even more efficient.

Rubidium and caesium can form even more complicated oxides than the superoxides. Rubidium can form Rb₆O and Rb₉O₂ upon oxidation in air, while caesium forms an immense variety of oxides, such as the ozonide CsO₃ and several brightly coloured suboxides, such as Cs₇O, Cs₄O, Cs₁₁O₃, Cs₃O (dark-green), CsO, Cs₃O₂, as well as Cs₇O₂. The latter may be heated under vacuum to generate Cs₂O.

The alkali metals can also react analogously with the heavier chalcogens (sulfur, selenium, tellurium, and polonium), and all the alkali metal chalcogenides are known (with the exception of francium’s). Reaction with an excess of the chalcogen can similarly result in lower chalcogenides, with chalcogen ions containing chains of the chalcogen atoms in question. For example, sodium can react with sulfur to form the sulfide (Na₂S) and various polysulfides with the formula Na₂Sx (x from 2 to 6), containing the S^2−x ions. Due to the basicity of the Se²⁻ and Te²⁻ ions, the alkali metal selenides and tellurides are alkaline in solution; when reacted directly with selenium and tellurium, alkali metal polyselenides and polytellurides are formed along with the selenides and tellurides with the Se^2−x and Te^2−x ions. The alkali metal polonides are all ionic compounds containing the Po²⁻ ion; they are very chemically stable and can be produced by direct reaction of the elements at around 300–400 °C.

Reaction with hydrogen and the halogens (alkali metal hydrides and halides)

The alkali metals are among the most electropositive elements on the periodic table and thus tend to bond ionically to the most electronegative elements on the periodic table, the halogens, forming salts known as the alkali metal halides. This includes sodium chloride, otherwise known as common salt. The reactivity becomes higher from lithium to caesium and drops from fluorine to iodine. All of the alkali metal halides have the formula MX where M is an alkali metal and X is a halogen. They are all white ionic crystalline solids. All the alkali metal halides are soluble in water except for lithium fluoride (LiF), which is insoluble in water due to its very high lattice enthalpy. The high lattice enthalpy of lithium fluoride is due to the small sizes of the Li+ and F− ions, causing the electrostatic interactions between them to be strong. The alkali metals also react similarly with hydrogen to form ionic alkali metal hydrides.

Coordination complexes

Alkali metal cations do not usually form coordination complexes with simple Lewis bases due to their low charge of just +1 and their relatively large size; thus the Li+ ion forms most complexes and the heavier alkali metal ions form less and less. In aqueous solution, the alkali metal ions exist as octahedral hexahydrate complexes ([M(H₂O)₆)]+), with the exception of the lithium ion, which due to its small size forms tetrahedral tetrahydrate complexes ([Li(H₂O)₄)]+); the alkali metals form these complexes because their ions are attracted by electrostatic forces of attraction to the polar water molecules. Because of this, anhydrous salts containing alkali metal cations are often used as desiccants. Alkali metals also readily form complexes with crown ethers (e.g. 12-crown-4 for Li+, 15-crown-5 for Na+, and 18-crown-6 for K+) and cryptands due to electrostatic attraction.

Ammonia solutions

Unlike most metals, the alkali metals dissolve slowly in liquid ammonia, forming hydrogen gas and the alkali metal amide (MNH₂, where M represents an alkali metal). The process may be speeded up by a catalyst. The amide salt is quite insoluble and readily precipitates out of solution, leaving intensely coloured ammonia solutions of the alkali metals. The colour is due to the presence of solvated electrons, which contribute to the high electrical conductivity of these solutions. At low concentrations (below 3 M), the solution is dark blue and has ten times the conductivity of aqueous sodium chloride; at higher concentrations (above 3 M), the solution is copper-coloured and has approximately the conductivity of liquid metals like mercury. In addition to the alkali metal amide salt and solvated electrons, such ammonia solutions also contain the alkali metal cation (M+), the neutral alkali metal atom (M), diatomic alkali metal molecules (M₂) and alkali metal anions (M−). These are unstable and eventually become the more thermodynamically stable alkali metal amide and hydrogen gas. Solvated electrons are powerful reducing agents and are often used in chemical synthesis.

Organometallic chemistry

Being the smallest alkali metal, lithium forms the widest variety of and most stable organometallic compounds, which are bonded covalently. Organolithium compounds are electrically non-conducting volatile solids or liquids that melt at low temperatures, and tend to form oligomers with the structure (RLi)x where R is the organic group. As the electropositive nature of lithium puts most of the charge density of the bond on the carbon atom, effectively creating a carbanion, organolithium compounds are extremely powerful bases and nucleophiles. For use as bases, butyllithiums are often used and are commercially available. An example of an organolithium compound is methyllithium ((CH₃Li)x), which exists in tetrameric (x = 4) and hexameric (x = 6) forms.

The application of organosodium compounds in chemistry is limited in part due to competition from organolithium compounds, which are commercially available and exhibit more convenient reactivity. The principal organosodium compound of commercial importance is sodium cyclopentadienide. Sodium tetraphenylborate can also be classified as an organosodium compound since in the solid state sodium is bound to the aryl groups. Organometallic compounds of the higher alkali metals are even more reactive than organosodium compounds and of limited utility. A notable reagent is Schlosser’s base, a mixture of n-butyllithium and potassium tert-butoxide. This reagent reacts with propene to form the compound allylpotassium (KCH₂CHCH₂). cis-2-Butene and trans-2-butene equilibrate when in contact with alkali metals. Whereas isomerization is fast with lithium and sodium, it is slow with the higher alkali metals. The higher alkali metals also favor the sterically congested conformation. Several crystal structures of organopotassium compounds have been reported, establishing that they, like the sodium compounds, are polymeric. Organosodium, organopotassium, organorubidium and organocaesium compounds are all mostly ionic and are insoluble (or nearly so) ionpolar solvents.

Physical

The alkali metals are all silver-coloured except for caesium, which has a golden tint. All are soft and have low densities, melting points, and boiling points.

The table below is a summary of the key physical and atomic properties of the alkali metals. Data marked with question marks are either uncertain or are estimations partially based on periodic trends rather than observations.

The alkali metals are more similar to each other than the elements in any other group are to each other.  For instance, when moving down the table, all known alkali metals show increasing atomic radius, decreasing electronegativity,  increasing reactivity, and decreasing melting and boiling points. In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium.

Atomic and ionic radii

The atomic radii of the alkali metals increase going down the group. Because of the shielding effect, when an atom has more than one electron shell, each electron feels electric repulsion from the other electrons as well as electric attraction from the nucleus. In the alkali metals, the outermost electron only feels a net charge of +1, as some of the nuclear charge (which is equal to the atomic number) is cancelled by the inner electrons; the number of inner electrons of an alkali metal is always one less than the nuclear charge. Therefore, the only factor which affects the atomic radius of the alkali metals is the number of electron shells. Since this number increases down the group, the atomic radius must also increase down the group.

The ionic radii of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different electron shell than the inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the effective nuclear charge has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases.

First ionisation energy

Periodic trend for ionisation energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases.

The first ionisation energy of an element or molecule is the energy required to move the most loosely held electron from one mole of gaseous atoms of the element or molecules to form one mole of gaseous ions with electric charge +1. The factors affecting the first ionisation energy are the nuclear charge, the amount of shielding by the inner electrons and the distance from the most loosely held electron from the nucleus, which is always an outer electron in main group elements. The first two factors change the effective nuclear charge the most loosely held electron feels. Since the outermost electron of alkali metals always feel the same effective nuclear charge (+1), the only factor which affects the first ionisation energy is the distance from the outermost electron to the nucleus. Since this distance increases down the group, the outermost electron feels less attraction from the nucleus and thus the first ionisation energy decreases. (This trend is broken in francium due to the relativistic stabilization and contraction of the 7s orbital, bringing francium’s valence electron closer to the nucleus than would be expected from non-relativistic calculations. This makes francium’s outermost electron feel more attraction from the nucleus, increasing its first ionisation energy slightly beyond that of caesium.)

The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled electron shell and is thus difficult to remove.

Reactivity

The reactivities of the alkali metals increase going down the group. This is the result of a combination of two factors: the first ionisation energies and atomisation energies of the alkali metals. Because the first ionisation energy of the alkali metals decreases down the group, it is easier for the outermost electron to be removed from the atom and participate in chemical reactions, thus increasing reactivity down the group. The atomisation energy measures the strength of the metallic bond of an element, which falls down the group as the atoms increase in radius and thus the metallic bond must increase in length, making the delocalised electrons further away from the attraction of the nuclei of the heavier alkali metals. Adding the atomisation and first ionisation energies gives a quantity closely related to (but not equal to) the activation energy of the reaction of an alkali metal with another substance. This quantity decreases going down the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity increases down the group.

Electronegativity

The variation of Pauling electronegativity (y-axis) as one descends the main groups of the periodic table from the second to the sixth period

Electronegativity is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself. If the bond between sodium and chlorine in sodium chloride were covalent, the pair of shared electrons would be attracted to the chlorine because the effective nuclear charge on the outer electrons is +7 in chlorine but is only +1 in sodium. The electron pair is attracted so close to the chlorine atom that they are practically transferred to the chlorine atom (an ionic bond). However, if the sodium atom was replaced by a lithium atom, the electrons will not be attracted as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more strongly attracted to the closer effective nuclear charge from lithium. Hence, the larger alkali metal atoms (further down the group) will be less electronegative as the bonding pair is less strongly attracted towards them.

Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example, lithium iodide (LiI) will dissolve in organic solvents, a property of most covalent compounds. Lithium fluoride (LiF) is the only alkali halide that is not soluble in water,  and lithium hydroxide (LiOH) is the only alkali metal hydroxide that is not deliquescent.

Melting and boiling points

The melting point of a substance is the point where it changes state from solid to liquid while the boiling point of a substance (in liquid state) is the point where the vapor pressure of the liquid equals the environmental pressure surrounding the liquid and all the liquid changes state to gas. As a metal is heated to its melting point, the metallic bonds keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal’s boiling point. Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali metals decreases down the group. This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons. As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points. (The increased nuclear charge is not a relevant factor due to the shielding effect.)

Density

The alkali metals all have the same crystal structure (body-centred cubic) and thus the only relevant factors are the number of atoms that can fit into a certain volume and the mass of one of the atoms, since density is defined as mass per unit volume. The first factor depends on the volume of the atom and thus the atomic radius, which increases going down the group; thus, the volume of an alkali metal atom increases going down the group. The mass of an alkali metal atom also increases going down the group. Thus, the trend for the densities of the alkali metals depends on their atomic weights and atomic radii; if figures for these two factors are known, the ratios between the densities of the alkali metals can then be calculated. The resultant trend is that the densities of the alkali metals increase down the table, with an exception at potassium. Due to having the lowest atomic weight of all the elements in their period and having the largest atomic radius for their periods, the alkali metals are the least dense metals in the periodic table. Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water.

All the alkali metals have odd atomic numbers; hence, their isotopes must be either odd-odd (both proton and neutroumber are odd) or odd-even (protoumber is odd, but neutroumber is even). Odd-odd nuclei have even mass numbers, while odd-eveuclei have odd mass numbers. Odd-odd primordial nuclides are rare because most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.

Due to the great rarity of odd-odd nuclei, almost all the primordial isotopes of the alkali metals are odd-even (the exceptions being the light stable isotope lithium-6 and the long-lived radioisotope potassium-40). For a given odd mass number, there can be only a single beta-stable nuclide, since there is not a difference in binding energy between even-odd and odd-even comparable to that between even-even and odd-odd, leaving other nuclides of the same mass number (isobars) free to beta decay toward the lowest-mass nuclide. An effect of the instability of an odd number of either type of nucleons is that odd-numbered elements, such as the alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 monoisotopic elements that have only a single stable isotope, all but one have an odd atomic number and all but one also have an eveumber of neutrons. Beryllium is the single exception to both rules, due to its low atomic number.

All of the alkali metals except lithium and caesium have at least one naturally occurring radioisotope: sodium-22 and sodium-24 are trace radioisotopes produced cosmogenically, potassium-40 and rubidium-87 have very long half-lives and thus occur naturally, and all isotopes of francium are radioactive. Caesium was also thought to be radioactive in the early 20th century, although it has no naturally occurring radioisotopes. (Francium had not been discovered yet at that time.) The natural radioisotope of potassium, potassium-40, makes up about 0.012% of natural potassium, and thus natural potassium is weakly radioactive. This natural radioactivity became a basis for a mistaken claim of the discovery for element 87 (the next alkali metal after caesium) in 1925.

Caesium-137, with a half-life of 30.17 years, is one of the two principal medium-lived fission products, along with strontium-90, which are responsible for most of the radioactivity of spent nuclear fuel after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left from the Chernobyl accident. 137Cs undergoes high-energy beta decay and eventually becomes stable barium-137. It is a strong emitter of gamma radiation. 137Cs has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay. 137Cs has been used as a tracer in hydrologic studies, analogous to the use of tritium. Small amounts of caesium-134 and caesium-137 were released into the environment during nearly all nuclear weapon tests and some nuclear accidents, most notably the Goiânia accident and the Chernobyl disaster. As of 2005, caesium-137 is the principal source of radiation in the zone of alienation around the Chernobyl nuclear power plant.

Extensions

Empirical (Na–Cs, Mg–Ra) and predicted (Fr–Uhp, Ubn–Uhh) atomic radius of the alkali and alkaline earth metals from the third to the ninth period, measured in angstroms.

Although francium is the heaviest alkali metal that has been discovered, there has been some theoretical work predicting the physical and chemical characteristics of the hypothetical heavier alkali metals. Being the first period 8 element, the undiscovered element ununennium (element 119) is predicted to be the next alkali metal after francium and behave much like their lighter congeners; however, it is also predicted to differ from the lighter alkali metals in some properties. Its chemistry is predicted to be closer to that of potassiumor rubidium instead of caesium or francium. This is unusual as periodic trends, ignoring relativistic effects would predict ununennium to be even more reactive than caesium and francium. This lowered reactivity is due to the relativistic stabilisation of ununennium’s valence electron, increasing ununennium’s first ionisation energy and decreasing the metallic and ionic radii; this effect is already seen for francium. This assumes that ununennium will behave chemically as an alkali metal, which, although likely, may not be true due to relativistic effects. The relativistic stabilisation of the 8s orbital also increases ununennium’s electron affinity far beyond that of caesium and francium; indeed, ununennium is expected to have an electron affinity higher than all the alkali metals lighter than it. Relativistic effects also cause a very large drop in the polarisability of ununennium. On the other hand, ununennium is predicted to continue the trend of melting points decreasing going down the group, being expected to have a melting point between 0 °C and 30 °C.

Empirical (Na–Fr) and predicted (Uue) electron affinity of the alkali metals from the third to the eighth period, measured in electron volts

The stabilisation of ununennium’s valence electron and thus the contraction of the 8s orbital cause its atomic radius to be lowered to 240 pm, very close to that of rubidium (247 pm), so that the chemistry of ununennium in the +1 oxidation state should be more similar to the chemistry of rubidium than to that of francium. On the other hand, the ionic radius of the Uue+ ion is predicted to be larger than that of Rb+, because the 7p orbitals are destabilised and are thus larger than the p-orbitals of the lower shells. Ununennium may also show the +3 oxidation state, which is not seen in any other alkali metal, in addition to the +1 oxidation state that is characteristic of the other alkali metals and is also the main oxidation state of all the known alkali metals: this is because of the destabilisation and expansion of the 7p3/2 spinor, causing its outermost electrons to have a lower ionisation energy than what would otherwise be expected. Indeed, many ununennium compounds are expected to have a large covalent character, due to the involvement of the 7p3/2 electrons in the bonding.

Not as much work has been done predicting the properties of the alkali metals beyond ununennium. Although a simple extrapolation of the periodic table would put element 169, unhexennium, under ununennium, Dirac-Fock calculations predict that the next alkali metal after ununennium may actually be element 165, unhexpentium, which is predicted to have the electron configuration [Uuo] 5g18 6f14 7d10 8s2 8p1/22 9s1. Further calculations show that unhexpentium would follow the trend of increasing ionisation energy beyond caesium, having an ionisation energy comparable to that of sodium, and that it should also continue the trend of decreasing atomic radii beyond caesium, having an atomic radius comparable to that of potassium. However, the 7d electrons of unhexpentium may also be able to participate in chemical reactions along with the 9s electron, possibly allowing oxidation states beyond +1 and perhaps even making unhexpentium behave more like a boron group element than an alkali metal.

The probable properties of the alkali metals beyond unhexpentium have not been explored yet as of 2012. In periods 8 and above of the periodic table, relativistic and shell-structure effects become so strong that extrapolations from lighter congeners become completely inaccurate. In addition, the relativistic and shell-structure effects (which stabilise the s-orbitals and destabilise and expand the d-, f-, and g-orbitals of higher shells) have opposite effects, causing even larger difference between relativistic and non-relativistic calculations of the properties of elements with such high atomic numbers. Due to the alkali and alkaline earth metals both being s-block elements, these predictions for the trends and properties of ununennium and unhexpentium also mostly apply to the corresponding alkaline earth metals unbinilium (Ubn) and unhexhexium (Uhh).

Other similar substances

Hydrogen

The element hydrogen, with one electron per neutral atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not normally considered to be an alkali metal; when it is considered to be an alkali metal, it is because of its atomic properties and not its chemical properties. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule (H₂); however, the alkali metals only form diatomic molecules (such as dilithium, Li2) at high temperatures, when they are in the gaseous state.

Hydrogen, like the alkali metals, has one valence electron and reacts easily with the halogens, but the similarities end there. Its placement above lithium is primarily due to its electron configuration and not its chemical properties. It is sometimes placed above carbon due to their similar electronegativities or fluorine due to their similar chemical properties.

The first ionisation energy of hydrogen (1312.0 kJ/mol) is much higher than that of the alkali metals. As only one additional electron is required to fill in the outermost shell of the hydrogen atom, hydrogen often behaves like a halogen, forming the negative hydride ion, and is sometimes considered to be a halogen. (The alkali metals can also form negative ions, known as alkalides, but these are little more than laboratory curiosities, being unstable.) Under extremely high pressures, such as those found at the cores of Jupiter and Saturn, hydrogen does become metallic and behaves like an alkali metal; in this phase, it is known as metallic hydrogen.

Sodium

Sodium compounds have been known since ancient times; salt (sodium chloride) has been an important commodity in human activities, as testified by the English word salary, referring to salarium, the wafers of salt sometimes given to Roman soldiers along with their other wages. In medieval Europe a compound of sodium with the Latiame of sodanum was used as a headache remedy. Pure sodium was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda (now called sodium hydroxide), a very similar method to the one used to isolate potassium earlier that year.

Potassium

While potash has been used since ancient times, it was not understood for most of its history to be a fundamentally different substance from sodium mineral salts. Georg Ernst Stahl obtained experimental evidence which led him to suggest the fundamental difference of sodium and potassium salts in 1702, and Henri Louis Duhamel du Monceau was able to prove this difference in 1736. The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus Antoine Lavoisier did include the alkali in his list of chemical elements in 1789. Pure potassium was first isolated in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH, potassium hydroxide) by the use of electrolysis of the molten salt with the newly invented voltaic pile. Potassium was the first metal that was isolated by electrolysis. Later that same year, Davy reported extraction of sodium from the similar substance caustic soda (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different. Rubidium

Rubidium was discovered in 1861 in Heidelberg, Germany by Robert Bunsen and Gustav Kirchhoff, the first people to suggest finding new elements by spectrum analysis, in the mineral lepidolite through the use of a spectroscope. Because of the bright red lines in its emission spectrum, they chose a name derived from the Latin word rubidus, meaning dark red or bright red. Rubidium’s discovery succeeded that of caesium, also discovered by Bunsen and Kirchhoff through spectroscopy.

Caesium

In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in the mineral water from Dürkheim, Germany. Due to the bright-blue lines in its emission spectrum, they chose a name derived from the Latin word caesius, meaning sky-blue. Caesium was the first element to be discovered spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.

Francium

There were at least four erroneous and incomplete discoveries before Marguerite Perey of the Curie Institute in Paris, France discovered francium in 1939 by purifying a sample of actinium-227, which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one that was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227. Perey then attempted to determine the proportion of beta decay to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure that she later revised to 1%. It was the last element discovered in nature, rather than by synthesis.

Eka-francium

The next element below francium (eka-francium) is very likely to be ununennium (Uue), element 119, although this is not completely certain due to relativistic effects. The synthesis of ununennium was first attempted in 1985 by bombarding a target of einsteinium-254 with calcium-48 ions at the superHILAC accelerator at Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.

It is highly unlikely that this reaction will be able to create any atoms of ununennium in the near future, given the extremely difficult task of making sufficient amounts of 254Es, which is favoured for production of ultraheavy elements because of its large mass, relatively long half-life of 270 days, and availability in significant amounts of several micrograms, to make a large enough target to increase the sensitivity of the experiment to the required level; einsteinium has not been found iature and has only been produced in laboratories. However, given that ununennium is only the first period 8 element on the extended periodic table, it may well be discovered in the near future through other reactions; indeed, another attempt to synthesise ununennium by bombarding a berkelium target with titanium ions is under way at the GSI Helmholtz Centre for Heavy Ion Research in Darmstadt, Germany. Currently, none of the period 8 elements have been discovered yet, and it is also possible, due to drip instabilities, that only the lower period 8 elements, up to around element 128, are physically possible. No attempts at synthesis have been made for any heavier alkali metals, such as unhexpentium, due to their extremely high atomic number.

Occurrence

Estimated abundances of the chemical elements in the Solar system. Hydrogen and helium are most common, from the Big Bang. The next three elements (lithium, beryllium, and boron) 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. Iron is especially common because it represents the minimum energy nuclide that can be made by fusion of helium in supernovae.

The Oddo-Harkins rule holds that elements with even atomic numbers are more common that those with odd atomic numbers, with the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability. All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the noble gases and the alkaline earth metals) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones as the alkali metals from rubidium onward can only be synthesized in supernovae and not in stellar nucleosynthesis. Lithium is also much less abundant than sodium and potassium as it is poorly synthesized in both Big Bang nucleosynthesis and in stars: the Big Bang could only produce trace quantities of lithium, beryllium and boron due to the absence of a stable nucleus with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this bottleneck by the triple-alpha process, fusing three helium nuclei to form carbon, and skipping over those three elements.

On Earth

The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the formation and evolution of the solar system. In turn, the natural history of the Earth caused parts of this planet to have differing concentrations of the elements. The mass of the Earth is approximately 5.98×1024 kg. It is composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to mass segregation, the core region is believed to be primarily composed of iron (88.8%), with smaller amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements.

The alkali metals, due to their high reactivity, do not occur naturally in pure form iature. They are lithophiles and therefore remain close to the Earth’s surface because they combine readily with oxygen and so associate strongly with silica, forming relatively low-density minerals that do not sink down into the Earth’s core. Potassium, rubidium and caesium are also incompatible elements due to their low ionic radii.

Sodium and potassium are very abundant in earth, both being among the ten most common elements in Earth’s crust; sodium makes up approximately 2.6% of the Earth’s crust measured by weight, making it the sixth most abundant element overall and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth’s crust and is the seventh most abundant element. Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater. Other solid deposits include halite, amphibole, cryolite, nitratine, and zeolite.

Lithium, due to its relatively low reactivity, can be found in seawater in large amounts; it is estimated that seawater is approximately 0.14 to 0.25 parts per million (ppm) or 25 micromolar.

Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite. Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.

Francium-223, the only naturally occurring isotope of francium, is the product of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals. In a given sample of uranium, there is estimated to be only one francium atom for every 1018 uranium atoms. It has been calculated that there is at most 30 g of francium in the earth’s crust at any time, due to its extremely short half-life of 22 minutes.

Production and isolation

Salt flats are rich in lithium, such as these in Salar del Hombre Muerto, Argentina (left) and Uyuni, Bolivia (right). The lithium-rich brine is concentrated by pumping it into solar evaporation ponds (visible in Argentina image).

The production of pure alkali metals is difficult due to their extreme reactivity with commonly used substances, such as water. The alkali metals are so reactive that they cannot be displaced by other elements and must be isolated through high-energy methods such as electrolysis.

Lithium salts have to be extracted from the water of mineral springs, brine pools, and brine deposits. The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride.

Potassium occurs in many minerals, such as sylvite (potassium chloride). It is occasionally produced through separating the potassium from the chlorine in potassium chloride, but is more often produced through electrolysis of potassium hydroxide, found extensively in places such as Canada, Russia, Belarus, Germany, Israel, United States, and Jordan, in a method similar to how sodium was produced in the late 1800s and early 1900s. It can also be produced from seawater. Sodium occurs mostly in seawater and dried seabed, but is now produced through electrolysis of sodium chloride by lowering the melting point of the substance to below 700 °C through the use of a Downs cell. Extremely pure sodium can be produced through the thermal decomposition of sodium azide.

For several years in the 1950s and 1960s, a by-product of the potassium production called Alkarb was a main source for rubidium. Alkarb contained 21% rubidium while the rest was potassium and a small fraction of caesium. Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite. Today, a common method for separating rubidium from potassium and caesium is the fractional crystallization of a rubidium and caesium alum (Cs,Rb)Al(SO₄)₂·12H₂O, which yields pure rubidium alum after approximately 30 different reactions. The limited applications and the lack of a mineral rich in rubidium limits the production of rubidium compounds to 2 to 4 tonnes per year. Caesium, however, is not produced from the above reaction. Instead, the mining of pollucite ore is the main method of obtaining pure caesium, extracted from the ore mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.

Francium-223, the only naturally occurring isotope of francium, is produced naturally as the product of the alpha decay of actinium-227. Francium can be found in trace amounts in uranium and thorium minerals; it has been calculated that at most there are 30 g of francium in the earth’s crust at any given time. As a result of its extreme rarity iature, most francium is synthesized in the nuclear reaction 197Au + 18O → 210Fr + 5 n, yielding francium-209, francium-210, and francium-211. The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms, which were synthesized using the nuclear reaction given above.

From their silicate ores, all the alkali metals may be obtained the same way: sulfuric acid is first used to dissolve the desired alkali metal ion and aluminium(III) ions from the ore (leaching), whereupon basic precipitation removes aluminium ions from the mixture by precipitating it as the hydroxide. The remaining insoluble alkali metal carbonate is then precipitated selectively; the salt is then dissolved in hydrochloric acid. The result is then left to evaporate and the alkali metal can then be isolated through electrolysis.

Lithium and sodium are typically isolated through electrolysis from their liquid chlorides, with calcium chloride typically added to lower the melting point of the mixture. The heavier alkali metals, however, is more typically isolated in a different way, where a reducing agent (typically sodium for potassium and magnesium or calcium for the heaviest alkali metals) is used to reduce the alkali metal chloride. The liquid or gaseous product (the alkali metal) then undergoes fractional distillation for purification.

Applications

All of the discovered alkali metals excluding francium have many applications. Lithium is often used in batteries, and lithium oxide can help process silica. Lithium can also be used to make lubricating greases, air treatment, and aluminium production.

Pure sodium has many applications, including use in sodium-vapour lamps, which produce very efficient light compared to other types of lighting, and can help smooth the surface of other metals. Sodium compounds have many applications as well, the most well-known compound being table salt. Sodium is also used in soap as salts of fatty acids.

Potassium compounds are often used as fertilisers as potassium is an important element for plant nutrition. Other potassium ions are often used to hold anions. Potassium hydroxide is a very strong base, and is used to control the pH of various substances.

Rubidium and caesium are often used in atomic clocks. Caesium atomic clocks are extraordinarily accurate; if a clock had been made at the time of the dinosaurs, it would be off by less than four seconds (after 80 million years). For that reason, caesium atoms are used as the definition of the second. Rubidium ions are often used in purple fireworks, and caesium is often used in drilling fluids in the petroleum industry.

Francium has no commercial applications, but because of francium’s relatively simple atomic structure, among other things, it has been used in spectroscopy experiments, leading to more information regarding energy levels and the coupling constants between subatomic particles. Studies on the light emitted by laser-trapped francium-210 ions have provided accurate data on transitions between atomic energy levels, similar to those predicted by quantum theory.

Biological role and precautions

Lithium naturally only occurs in traces in biological systems and has no known biological role, but does have effects on the body when ingested. Lithium carbonate is used as a mood stabiliser in psychiatry to treat bipolar disorder (manic-depression) in daily doses of about 0.5 to 2 grams, although there are side-effects. Excessive ingestion of lithium causes drowsiness, slurred speech and vomiting, among other symptoms, and poisons the central nervous system,  which is dangerous as the required dosage of lithium to treat bipolar disorder is only slightly lower than the toxic dosage. Its biochemistry, the way it is handled by the human body and studies using rats and goats suggest that it is an essential trace element, although the natural biological function of lithium in humans has yet to be identified.

Sodium and potassium occur in all known biological systems, generally functioning as electrolytes inside and outside cells. 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 (also known as common salt) 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 1.5 grams per day, but most people in the United States consume more than 2.3 grams per day, the minimum amount that promotes hypertension; this in turn causes 7.6 million premature deaths worldwide.

Potassium is the major cation (positive ion) inside animal cells, while sodium is the major cation outside animal cells. The concentration differences of these charged particles causes a difference in electric potential between the inside and outside of cells, known as the membrane potential. The balance between potassium and sodium is maintained by ion pumps in the cell membrane. The cell membrane potential created by potassium and sodium ions allows the cell to generate an action potential ‑ a “spike” of electrical discharge. The ability of cells to produce electrical discharge is critical for body functions such as neurotransmission, muscle contraction, and heart function.

Rubidium has no known biological role, but may help stimulate metabolism, and, similarly to caesium replace potassium in the body causing potassium deficiency. Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium, allowing the caesium to replace the potassium in the body, causing potassium deficiency. Exposure to large amounts of caesium compounds can cause hyperirritability and spasms, but as such amounts would not ordinarily be encountered iatural sources, caesium is not a major chemical environmental pollutant. The median lethal dose (LD50) value for caesium chloride in mice is 2.3 g per kilogram, which is comparable to the LD50 values of potassium chloride and sodium chloride. Caesium chloride has been promoted as an alternative cancer therapy, but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment. Radioisotopes of caesium require special precautions: the improper handling of caesium-137 gamma ray sources can lead to release of this radioisotope and radiation injuries. Perhaps the best-known case is the Goiânia accident of 1987, in which an improperly-disposed-of radiation therapy system from an abandoned clinic in the city of Goiânia, Brazil, was scavenged from a junkyard, and the glowing caesium salt sold to curious, uneducated buyers. This led to four deaths and serious injuries from radiation exposure. Together with caesium-134, iodine-131, and strontium-90, caesium-137 was among the isotopes distributed by the Chernobyl disaster which constitute the greatest risk to health.

Francium has no biological role and is most likely to be toxic due to its extreme radioactivity, causing radiation poisoning, but since the greatest quantity of francium ever assembled to date is about 300,000 neutral atoms, it is unlikely that most people will ever encounter francium.

 

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

Structural Biochemistry/Alkali Metals

The alkali metals are a series of chemical elements forming Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals). The alkali metals provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behavior down the group.

Alkali Metals generally exist as cations with a 1+ charge, and exists as lustrous metals that conducts both heat and electricity well. They are water soluble, basic and very reactive. Alkali metals are the least electronegative group, and tend to react best with nonmetals and water.

Potassium and Sodium are vital in biochemical systems. Valinomycin complexes selectively with potassium for transporting molecules across the cell membrane. Crown ethers and cryptands are structures that complex with alkali metal ions; supramolecular chemists are studying these interactions.

The sodium and potassium pump are important in maintaining equilibrium amounts both inside and outside the cell. Sodium is concentrated outside, whereas potassium is concentrated inside. Both of these metals also help to communicate electrical signals ierves and in the heart. Humans obtain sodium by eating foods with table salt and baking soda. To obtain potassium, we can eat bananas, oranges, and avocados.

Lithium is the lightest of the alkali metals. Traces of lithium ion occur in animal tissues, but it has no known physiological role. Lithium salts were introduced into psychiatry in 1949 for the treatment of mania. It is believed that the lithium ion at concentrations of 1 to 10 mEq per liter inhibits the depolarization-provoked and calcium-dependent release of norepinephrine and dopamine, but not 5-HT, from nerve terminals. Lithium ion is small ( its radius is around 0.6 Angstrom) and mobile ( fast exchange ligands ). However, it weakly binds to ligands and is easily hydrated by water. In medicine, Lithum is used to control bipolar affective disorder such as manic depression.

Sodium reacts spontaneously with water and rapidly with oxygen. It is a silvery, soft solid in its pure form. There is an abundant amount of the sodium element, most commonly recognized as sodium chloride, or table salt. Sodium chloride is an important nutrition for animals.

Potassium reacts very spontaneously with water, igniting the hydrogen, causing it to burst into lavender-colored flames. It is the sixth most abundant element on earth and is used in fertilizer and drain cleaners.

Rubidium has similar characteristics as potassium but is much more reactive.

Cesium is a silver metal that is one of the five elements that are liquid at room temperature. It also has similar characteristics to potassium but is extremely reactive.

Francium is the second rarest natural and most unstable element. All the francium isotopes are radioactive and have short half-lives.

The alkaline earth metals are a group of chemical elements in the periodic table with very similar properties. They are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure and readily lose their two outermost electrons to form cations with charge 2+ and an oxidation state, or oxidatioumber of +2. In the modern IUPAC nomenclature, the alkaline earth metals comprise the group 2 elements.

The chemical elements found in Group II of the periodic table are called alkaline earth metals. The six chemical elements that constitute the periodic table are beryllium, magnesium, calcium, strontium, barium and radium. These elements have +2 oxidatioumber, making them highly reactive and thus are not found free iature but as carbonates or sulfates. They share several similar characteristics as the alkali metals of Group I in the periodic table, and thus are placed just next to alkali metals. The most striking resemblance is their high reactivity.

Alkaline earth metals have two electrons in their valence shell of their atoms, and their general configuration is written as [Noble gas]ns2.

Element Smbol Atomic Number Electronic Configuration

Beryllium Be 4 [He]2s2

Magnesium Mg 12 [Ne]3s2

Calcium Ca 20 [Ar]4s2

Strontium Sr 38 [Kr]5s2

Barium Ba 56 [Xe]6s2

Radium Ra 88 [Rn]7s2

Atomic and Ionic Radius

Compared to the atomic and ionic radii of alkali metals, the radii of alkaline earth metals is smaller due to the higher nuclear charge which causes the electrons to be attracted to the nucleus, thereby resulting in reduction of atomic and ionic size. As we move from Be to Ra in the group, it is observed that the atomic radius increases due to the increase in the atomic number, number of shells and screening effect.

Reactivity

Alkali earth metals are not as reactive as alkali metals, however, they are more reactive than other elements in the periodic table. They react with water to form metal hydroxide and hydrogen gas. Due to the presence of two electrons in their valence shells, they are seen to form cations with a charge of 2. Since these elements have high reactivity, they are found in the form of compounds such as sulfates and carbonates instead of the free form. Magnesium in its pure form is combustible iature and burns with an intense white light when brought in contact with air (as it combines with air to form magnesium oxide). The other elements of this group also form oxides when reacted with oxygen.

Boiling and Melting Points

The temperature at which a solid element turns into liquid state and liquid element into vapor state is called melting and boiling points respectively. In case of alkaline metals, as the atomic number increases, the boiling point and melting point are seen to decrease in the group. However, in case of alkaline earth metals no regular pattern or trend is seen. Beryllium has the highest melting and boiling point in the group and Magnesium has the lowest. However, since the atomic radii of alkaline earth metals are smaller than those of alkali metals, the boiling and melting points of Group II is higher than Group I elements. Thus, we can use the boiling and melting points test to differentiate between Group I and II elements.

Ionization Energy

Since Group II elements have a larger radii, they have higher ionization energies as compared to Group I elements. As we move down Group II, the atomic number and size increases, as the number of shells increase and the magnitude of screening effect is also seen to be higher. However, this causes the ionization energy to decrease down the group.

Metallic Character and Density

Alkaline earth metals are electro-positive iature, which increases down the group. However, Group II elements are not as electro-positive as Group I elements due to their higher ionization energies. These elements are harder and denser as compared to alkali metals. This is because the smaller atomic size causes the electrons to be packed more closely, thereby forming strong metallic bonds. This is why these metals are harder and denser than alkali metals.

Physical and Chemical Behavior

Group II elements are mostly good conductors of electricity since they are highly metallic iature. They are shiny and generally white or silvery in color. They are soft metals but harder and denser than alkali metals. When freshly cut, these metals have a gray-white appearance, which is seen to tarnish as soon as it is exposed to air. Chemically, these elements are strong reducing agents as they have two valence electrons in their valence shell, which they readily give away during chemical bond formation. The free elements are soluble in liquid ammonia, imparting a metallic, copper-like appearance. These solutions are very helpful in various chemical processes.

Of the six elements in Group II, calcium is the most widely available element and not only ranks fifth among the elements found in the Earth’s crust, but also ranks fifth among the elements in the human body. Magnesium ranks as the eighth most abundant element in the Earth’s crust and seventh most abundant element in the human body. Barium and beryllium are poisonous and radium, being radioactive it’s exposure is harmful to humans.

Each of the six alkaline earth metals in Group II is unique and has different usages. Some are found in the human body, whereas others are harmful to the human body but useful for other purposes like radioactive processes and so on. These six metals also burn with distinctive flames and have unique flame coloration, which is why they are used in fireworks.

The alkaline earth metals are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). This group lies in the s-block of the periodic table as all alkaline earth metals have their outermost electron in an s-orbital.

All the discovered alkaline earth metals occur in nature. Experiments have been conducted to attempt the synthesis of element 120, which is likely to be the next member of the group, but they have all met with failure. However, element 120 may not be an alkaline earth metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements.

Characteristics

Chemical

Like other groups, the members of this family show patterns in its electronic configuration, especially the outermost shells, resulting in trends in chemical behavior:

Most of the chemistry has been observed only for the first five members of the group. The chemistry of radium is not well established due to its radioactivity; thus, the presentation of its properties here is limited.

The alkaline earth metals are all silver-colored, soft, and have relatively low densities, melting points, and boiling points. In chemical terms, all of the alkaline metals react with the halogens to form the alkaline earth metal halides, all of which are ionic crystalline compounds (except for beryllium chloride, which is covalent). All the alkaline earth metals except beryllium also react with water to form strongly alkaline hydroxides and thus should be handled with great care. The heavier alkaline earth metals react more vigorously than the lighter ones. The alkaline metals have the second-lowest first ionization energies in their respective periods of the periodic table because of their somewhat low effective nuclear charges and the ability to attain a full outer shell configuration by losing just two electrons. The second ionization energy of all of the alkaline metals is also somewhat low.

Beryllium is an exception: It does not react with water or steam, and its halides are covalent. If beryllium did form compounds with an ionization state of +2, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since beryllium has a high charge density. All compounds that include beryllium have a covalent bond. Even the compound beryllium fluoride, which is the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted.

All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions.

Compounds and reactions

The alkaline earth metals all react with the halogens to form ionic halides, such as calcium chloride (CaCl₂), as well as reacting with oxygen to form oxides such as strontium oxide (SrO). Calcium, strontium, and barium react with water to produce hydrogen gas and their respective hydroxides, and also undergo transmetalation reactions to exchange ligands.

 

Physical and atomic

The table below is a summary of the key physical and atomic properties of the alkaline earth metals.

 

Chemical Reactions and Compounds

The alkaline earth metals react directly with most nonmetallic elements. forming Except for beryllium, the alkaline earths react directly with hydrogen gas to form hydrides, MH₂; M = Mg, Ca, Sr, Ba, or Ra. Beryllium hydride, BeH₂ can also be prepared, but not directly from the elements. Alkaline-earth metals combine readily with oxygen from the air to form oxides, MO. This follows the general reaction:

2M(s) + O₂(g) → 2MO₂(s)      M = Be, Mg, Ca, Sr, Ba, or Ra      (1)

The following video shows the reaction of magnesium with oxygen:

Magnesium is burned in air, and emits a bright white flame. A white powder of MgO remains after the reaction described by the equation:

2Mg(s) + O₂(g) → 2MgO₂(s)

It should also be noted that while MgO is the main product, nitrogen is also present in the air, and so some magnesium nitride is also produced according to the chemical equation:

3Mg(s) + N₂(g) → Mg₃N₂(s)

These oxides will coat the surface of the metal and prevent other substances from contacting and reacting with it. A good example of the effect of such an oxide coating is the reaction of alkaline-earth metals with water. Beryllium and magnesium react much more slowly than the others because their oxides are insoluble and prevent water from contacting the metal.

Alkaline-earth metals react directly with halogens to form salts:

M(s) + Cl₂(g) → MCl₂(s)      M = Be, Mg, Ca, Sr, Ba, or Ra      (2)

Salt obtained by evaporating seawater (sea salt) contains a good deal of magnesium chloride and calcium chloride as well as sodium chloride. It also has small traces of iodide salts, accounting for the absence of simple goiter in communities which obtain their salt from the oceans. Simple goiter is an enlargement of the thyroid gland caused by iodine deficiency.

Alkaline earths also form sulfides: MS. In all these compounds the alkaline-earth elements occur as dipositive ions, Mg²⁺, Ca²⁺, Sr²⁺, or Ba²⁺.

Similar compounds of Be can be formed by roundabout means, but not by direct combination of the elements. Moreover, the Be compounds are more covalent than ionic. The Be²⁺ ion has a very small radius (31 pm) and is therefore capable of distorting (polarizing) the electron cloud of an anion in its vicinity. Therefore all bonds involving Be have considerable covalent character, and the chemistry of Be is significantly different from that of the other members of group IIA.

As in the case of the alkali metals, the most important and abundant alkaline earths, Mg and Ca, are in the third and fourth periods. Be is rare, although its strength and low density make it useful in certain special alloys. Sr and Ba occur naturally as the relatively insoluble sulfates SrSO₄ (strontianite) and BaSO₄ (barite), but these two elements are of minor commercial importance.

The most common ores of Mg and Ca are dolomite, MgCO₃•CaCO₃, after which an entire mountain range in Italy is named, and limestone, CaCO₃, an important building material. Mg is also recovered from seawater on a wide scale. The oxides of the alkaline earths are commonly obtained by heating the carbonates. For example, lime, CaO, is obtained from limestone as follows:

CaCO₃(s) CaO(s) + CO₂(g)

 

Except for BeO, which is covalently bonded, alkaline-earth oxides contain O^2– ions and are strongly basic. When treated with water (a process known as slaking), they are converted to hydroxides:

 

CaO(s) + H₂O(l) → Ca(OH)₂(s)

 

Ca(OH)₂ (slaked lime) is an important strong base for industrial applications, because it is cheaper than NaOH.

MgO has an extremely high melting point (2800°C) because of the close approach and large charges of its constituent Mg²⁺ and O^2– ions in the crystal lattice. As a solid it is a good electrical insulator, and so it is used to surround metal-resistance heating wires in electric ranges. MgO is also used to line high- temperature furnaces. When converted to the hydroxide, Mg finds a different use. Mg(OH)₂ is quite insoluble in water, and so it does not produce a high enough concentration of hydroxide ions to be caustic. It is basic, however, and gram for gram caeutralize nearly twice the quantity of acid that NaOH can. Consequently a suspension of Mg(OH)₂ in water (milk of magnesia) makes an excellent antacid, for those who can stand its taste.

Because the carbonate ion behaves as a Brönstedt-Lowry base, carbonate salts dissolve in acidic solutions. Iature, water often becomes acidic because the acidic oxide CO₂ is present in the atmosphere. When CO₂ from the air dissolves in water, it can help dissolve limestone:

 

CO₂(g) + H₂O(l) + CaCO₃(s) Ca²⁺(aq) + HCO₃^–(aq)

 

This reaction often occurs underground as rainwater saturated with CO₂ seeps through a layer of limestone. Caves from which the limestone has been dissolved are often prevalent in areas where there are large deposits of CaCO₃. In addition, the groundwater and well water in such areas becomes hard. Hard water contains appreciable concentrations of Ca²⁺, Mg²⁺ , and certain other metal ions. These form insoluble compounds with soap, causing curdy, scummy precipitates. Hard water can be softened by adding Na₂CO₃, washing soda, which precipitates CaCO₃, or by ion exchange, a process in which the undesirable Ca²⁺ and Mg²⁺ ions are replaced in solution by Na⁺ ions, which do not precipitate soap. Most home water softeners work on the latter principle.

Nuclear stability

All of the alkaline earth metals except magnesium and strontium have at least one naturally occurring radioisotope: beryllium-7, beryllium-10, and calcium-41 are trace radioisotopes, calcium-48 and barium-130 have very long half-lives and thus occur naturally, and all isotopes of radium are radioactive. Calcium-48 is the lightest nuclide to undergo double beta decay.

The natural radioisotope of calcium, calcium-48, makes up about 0.1874% of natural calcium, and thus natural calcium is weakly radioactive. Barium-130 makes up approximately 0.1062% of natural barium, and thus barium is weakly radioactive as well.

History

Etymology

The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These oxides are basic (alkaline) when combined with water. “Earth” is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating—properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier’s idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths, thus supporting Lavoisier’s hypothesis and causing the group to be named the alkaline earth metals.

Discovery

The calcium compounds calcite and lime have been known and used since prehistoric times. The same is true for the beryllium compounds beryl and emerald. The other compounds of the alkaline earth metals were discovered starting in the early 15th century. The magnesium compound magnesium sulfate was first discovered in 1618 by farmer at Epsom in England. Strontium carbonate was discovered in minerals in the Scottish village of Strontian in 1790. The last element was the least abundant radioactive radium which was extracted from uraninite in 1898.

All elements except beryllium were isolated by electrolysis of molten compounds. Magnesium, calcium and strontium were first produced by Humphry Davy in 1808, while beryllium was independently isolated Friedrich Wöhler and Antoine Bussy in 1828 by reacting berylium compounds with potassium. In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne also by electrolysis.

Beryllium

Emerald, a variety of beryl

Beryl, a mineral which contains beryllium, has been known since the time of the Ptolemaic dynasty in Egypt. Although it was originally thought that beryl was an aluminium silicate, beryl was later found to contain a then-unknown element when, in 1797, Louis-Nicolas Vauquelin dissolved aluminium hydroxide from beryl in an alkali. In 1828, Friedrich Wöhler and Antoine Bussy independently isolated this new element, beryllium, by the same method, which involved a reaction of beryllium chloride with metallic potassium; this reaction was not able to produce large ingots of beryllium. It was not until 1898, when Paul Lebeau performed an electrolysis of a mixture of beryllium fluoride and sodium fluoride that large pure samples of beryllium were produced.

Magnesium

Magnesium was first produced by Sir Humphry Davy in England in 1808 using electrolysis of a mixture of magnesia and mercuric oxide. Antoine Bussy prepared it in coherent form in 1831. Davy’s first suggestion for a name was magnium, but the name magnesium is now used.

Properties

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.

Compounds

Magnesium forms a variety of industrially and biologically important compounds, including magnesium oxide, various salts and others.

Isotopes

Magnesium has three stable isotopes: 24Mg, 25Mg and 26Mg. All are present in significant amounts (see table of isotopes above). About 79% of Mg is 24Mg. The isotope 28Mg is radioactive and in the 1950s to 1970s was made commercially by several nuclear power plants for use in scientific experiments. This isotope has a relatively short half-life (21 hours) and so its use was limited by shipping times.

26Mg has found application in isotopic geology, similar to that of aluminium. 26Mg is a radiogenic daughter product of 26Al, which has a half-life of 717,000 years. Large enrichments of stable 26Mg have been observed in the Ca-Al-rich inclusions of some carbonaceous chondrite meteorites. The anomalous abundance of 26Mg is attributed to the decay of its parent 26Al in the inclusions. Therefore, the meteorite must have formed in the solar nebula before the 26Al had decayed. Hence, these fragments are among the oldest objects in the solar system and have preserved information about its early history.

It is conventional to plot 26Mg/24Mg against an Al/Mg ratio. In an isochron dating plot, the Al/Mg ratio plotted is27Al/24Mg. The slope of the isochron has no age significance, but indicates the initial 26Al/27Al ratio in the sample at the time when the systems were separated from a common reservoir.

China is the dominant supplier of magnesium, with approximately 80% of the world market share. China is almost completely reliant on thesilicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon) to obtain the metal.

In the United States, magnesium is principally obtained by electrolysis of fused magnesium chloride from brines, wells, and sea water. At thecathode, the Mg²⁺  ion is reduced by two electrons to magnesium metal:

Mg²⁺ + 2 e− → Mg

At the anode, each pair of Cl− ions is oxidized to chlorine gas, releasing two electrons to complete the circuit:

2 Cl− → Cl₂ (g) + 2 e−

China is the dominant supplier of magnesium, with approximately 80% of the world market share. China is almost completely reliant on the silicothermic Pidgeon process (the reduction of the oxide at high temperatures with silicon) to obtain the metal.

In the United States, magnesium is principally obtained by electrolysis of fused magnesium chloride from brines, wells, and sea water. At the cathode, the Mg²⁺ ion is reduced by two electrons to magnesium metal:

Mg²⁺ + 2 e− → Mg

At the anode, each pair of Cl− ions is oxidized to chlorine gas, releasing two electrons to complete the circuit:

2 Cl− → Cl₂ (g) + 2 e−

A new process, solid oxide membrane technology, involves the electrolytic reduction of MgO. At the cathode Mg²⁺ ion is reduced by two electrons to magnesium metal. The electrolyte is Yttria-stabilized zirconia(YSZ). The anode is a liquid metal. At the YSZ/liquid metal anode O²⁻ is reduced. A layer of graphite borders the liquid metal anode, and at this interface carbon and oxygen react to form carbon monoxide. When silver is used as the liquid metal anode, there is no reductant carbon or hydrogeeeded, and only oxygen gas is evolved at the anode. It has been reported that this method provides a 40% reduction in cost per pound over the electrolytic reduction method. This method is more environmentally sound than others because there is much less carbon dioxide emitted.

The United States has traditionally been the major world supplier of this metal, supplying 45% of world production even as recently as 1995. Today, the US market share is at 7%, with a single domestic producer left, US Magnesium, a Renco Group company in Utah born from now-defunct Magcorp.

Biological roles

Because of the important interaction between phosphate and magnesium ions, magnesium ions are essential to the basic nucleic acid chemistry of life, and thus are essential to all cells of all known living organisms. Over 300 enzymes require the presence of magnesium ions for their catalytic action, including all enzymes utilizing or synthesizing ATP, or those that use other nucleotides to synthesize DNA and RNA. ATP exists in cells normally as a chelate of ATP and a magnesium ion.

Plants have an additional use for magnesium in that chlorophylls are magnesium-centered porphyrins. Magnesium deficiency in plants causes late-season yellowing between leaf veins, especially in older leaves, and can be corrected by applying Epsom salts (which is rapidly leached), or else crushed dolomitic limestone to the soil.

Magnesium is a vital component of a healthy human diet. Human magnesium deficiency (including conditions that show few overt symptoms) is relatively rare although only 32% of people in the United States meet the RDA-DRI; low levels of magnesium in the body has been associated with the development of a number of human illnesses such as asthma, diabetes, and osteoporosis. Taken in the proper amount magnesium plays a role in preventing both stroke and heart attack. The symptoms of people with fibromyalgia, migraines, and girls going through their premenstrual syndrome are less severe and magnesium can shorten the length of the migraine symptoms.

Adult human bodies contain about 24 grams of magnesium, with 60% in the skeleton, 39% intracellular (20% in skeletal muscle), and 1% extracellular. Serum levels are typically 0.7–1.0 mmol/L or 1.8–2.4 mEq/L. Serum magnesium levels may appear normal even in cases of underlying intracellular deficiency, although no known mechanism maintains a homeostatic level in the blood other than renal excretion of high blood levels.

Intracellular magnesium is correlated with intracellular potassium. Magnesium is absorbed in the gastrointestinal tract, with more absorbed when status is lower. Magnesium competes with calcium in the human body, in this way it actually keeps calcium in check. However, this can cause a calcium deficiency if calcium levels are already low. Low and high protein intake inhibit magnesium absorption, and other factors such as phosphate, phytate, and fat affect absorption. Excess dietary magnesium is excreted in feces, urine, and sweat. Magnesium status may be assessed roughly through serum and erythrocyte Mg concentrations and urinary and fecal excretion, but intravenous magnesium loading tests are likely the most accurate and practical in most people. In these tests, magnesium is injected intravenously; a retention of 20% or more indicates deficiency. Other nutrient deficiencies are identified through biomarkers, but none are established for magnesium.

The UK recommended daily values for magnesium is 300 mg for men and 270 mg for women. Spices, nuts, cereals, coffee, cocoa, tea, and vegetables are rich sources of magnesium. Green leafy vegetables such as spinach are also rich in magnesium as they contain chlorophyll. Observations of reduced dietary magnesium intake in modern Western countries compared to earlier generations may be related to food refining and modern fertilizers that contaio magnesium.

Numerous pharmaceutical preparations of magnesium, as well as magnesium dietary supplements are available. Magnesium oxide, one of the most common forms in magnesium dietary supplements because it has high magnesium content per weight, has been reported the least bioavailable. Magnesium citrate has been reported as more bioavailable than oxide or amino-acid chelate (glycinate) forms.

Excess magnesium in the blood is freely filtered at the kidneys, and for this reason it is difficult to overdose on magnesium from dietary sources alone. With supplements, overdose is possible, however, particularly in people with poor renal function; occasionally, with use of high cathartic doses of magnesium salts, severe hypermagnesemia has been reported to occur even without renal dysfunction. Alcoholism can produce a magnesium deficiency, which is easily reversed by oral or parenteral administration, depending on the degree of deficiency.

 

Calcium

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).

Calcium provides an important link between tectonics, climate and the carbon cycle. In the simplest terms, uplift of mountains exposes Ca-bearing rocks to chemical weathering and releases Ca²⁺ into surface water. This Ca²⁺ eventually is transported to the ocean where it reacts with dissolved CO₂ to form limestone. Some of this limestone settles to the sea floor where it is incorporated into new rocks. Dissolved CO₂, along with carbonate and bicarbonate ions, are referred to as dissolved inorganic carbon (DIC).The actual reaction is more complicated and involves the bicarbonate ion (HCO₃⁻) that forms when CO₂ reacts with water at seawater pH:

Ca²⁺ + 2HCO₃⁻ → CaCO₃ (limestone) + CO₂ + H₂O

Note that at ocean pH most of the CO₂ produced in this reaction is immediately converted back into HCO− The reaction results in a net transport of one molecule of CO₂ from the ocean/atmosphere into the lithosphere.

The result is that each Ca²⁺ ion released by chemical weathering ultimately removes one CO₂ molecule from the surficial system (atmosphere, ocean, soils and living organisms), storing it in carbonate rocks where it is likely to stay for hundreds of millions of years. The weathering of calcium from rocks thus scrubs CO₂ from the ocean and atmosphere, exerting a strong long-term effect on climate.Analogous cycles involving magnesium, and to a much smaller extent strontium and barium, have the same effect.

As the weathering of limestone (CaCO₃) liberates equimolar amounts of Ca²⁺ and CO₂, it has no net effect on the CO₂ content of the atmosphere and ocean. The weathering of silicate rocks like granite, on the other hand, is a net CO₂ sink because it produces abundant Ca²⁺ but very little CO₂.

Recommended adequate intake by the IOM for calcium:
Age	Calcium (mg/day)
0–6 months	200
7–12 months	260
1–3 years	700
4–8 years	1000
9–18 years	1300
19–50 years	1000
51–70 years (male)	1000
51–70 years (female)	1200
71+ years	1200

Nutrition

Calcium is an important component of a healthy diet and a mineral necessary for life. The National Osteoporosis Foundation says, “Calcium plays an important role in building stronger, denser bones early in life and keeping bones strong and healthy later in life.”

Approximately 99 percent of the body’s calcium is stored in the bones and teeth. The rest of the calcium in the body has other important uses, such as some exocytosis, especially neurotransmitter release, and muscle contraction. In the electrical conduction system of the heart, calcium replaces sodium as the mineral that depolarizes the cell, proliferating the action potential. In cardiac muscle, sodium influx commences an action potential, but during potassium efflux, the cardiac myocyte experiences calcium influx, prolonging the action potential and creating a plateau phase of dynamic equilibrium. Long-term calcium deficiency can lead to rickets and poor blood clotting and in case of a menopausal woman, it can lead to osteoporosis, in which the bone deteriorates and there is an increased risk of fractures. While a lifelong deficit can affect bone and tooth formation, over-retention can cause hypercalcemia (elevated levels of calcium in the blood), impaired kidney function and decreased absorption of other minerals. Several sources suggest a correlation between high calcium intake (2000 mg per day, or twice the U.S. recommended daily allowance, equivalent to six or more glasses of milk per day) and prostate cancer. High calcium intakes or high calcium absorption were previously thought to contribute to the development of kidney stones. However, a high calcium intake has been associated with a lower risk for kidney stones in more recent research. Vitamin D is needed to absorb calcium.

Dairy products, such as milk and cheese, are a well-known source of calcium. Some individuals are allergic to dairy products and even more people, in particular those of non Indo-European descent, are lactose-intolerant, leaving them unable to consume non-fermented dairy products in quantities larger than about half a liter per serving. Others, such as vegans, avoid dairy products for ethical and health reasons.

Many good vegetable sources of calcium exist, including seaweeds such as kelp, wakame and hijiki; nuts and seeds like almonds, hazelnuts, sesame, and pistachio; blackstrap molasses; beans (especially soy beans); figs; quinoa; okra; rutabaga; broccoli; dandelion leaves; and kale. In addition, for some drinks (like soy milk or orange juice) it is typical to be fortified with calcium.

Numerous vegetables, notably spinach, chard and rhubarb have a high calcium content, but they may also contain varying amounts of oxalic acid that binds calcium and reduces its absorption. The same problem may to a degree affect the absorption of calcium from amaranth, collard greens, and chicory greens. This process may also be related to the generation of calcium oxalate.

An overlooked source of calcium is eggshell, which can be ground into a powder and mixed into food or a glass of water.

The calcium content of most foods can be found in the USDA National Nutrient Database.

Osteoporosis

Such studies often do not test calcium alone, but rather combinations of calcium and vitamin D. Randomized controlled trials found both positive and negative effects. The different results may be explained by doses of calcium and underlying rates of calcium supplementation in the control groups.

Cancer

A meta-analysis by the international Cochrane Collaboration of two randomized controlled trials found that calcium “might contribute to a moderate degree to the prevention of adenomatous colonic polyps”.

More recent studies were conflicting, and one that was positive for effect (Lappe, et al.) did control for a possible anti-carcinogenic effect of vitamin D, which was found to be an independent positive influence from calcium-alone on cancer risk.

A randomized controlled trial found that 1000 mg of elemental calcium and 400 IU of vitamin D3 had no effect on colorectal cancer.

A randomized controlled trial found that 1400–1500 mg supplemental calcium and 1100 IU vitamin D3 reduced aggregated cancers with a relative risk of 0.402.

An observational cohort study found that high calcium and vitamin D intake was associated with “lower risk of developing premenopausal breast cancer.”

 

Hazards and toxicity

Compared with other metals, the calcium ion and most calcium compounds have low toxicity. This is not surprising given the very high natural abundance of calcium compounds in the environment and in organisms. Calcium poses few serious environmental problems, with kidney stones the most common side-effect in clinical studies. Acute calcium poisoning is rare, and difficult to achieve unless calcium compounds are administered intravenously. For example, the oral median lethal dose (LD50) for rats for calcium carbonate and calcium chloride are 6.45and 1.4 g/kg, respectively.

Calcium metal is hazardous because of its sometimes-violent reactions with water and acids. Calcium metal is found in some drain cleaners, where it functions to generate heat and calcium hydroxide that saponifies the fats and liquefies the proteins (e.g., hair) that block drains. When swallowed calcium metal has the same effect on the mouth, esophagus and stomach, and can be fatal.

Excessive consumption of calcium carbonate antacids/dietary supplements (such as Tums) over a period of weeks or months can cause milk-alkali syndrome, with symptoms ranging from hypercalcemia to potentially fatal renal failure. What constitutes “excessive” consumption is not well known and, it is presumed, varies a great deal from person to person. Persons consuming more than 10 grams/day of CaCO₃ (=4 g Ca) are at risk of developing milk-alkali syndrome, but the condition has been reported in at least one person consuming only 2.5 grams/day of CaCO₃ (=1 g Ca), an amount usually considered moderate and safe.

Oral calcium supplements diminish the absorption of thyroxine when taken within four to six hours of each other. Thus, people taking both calcium and thyroxine run the risk of inadequate thyroid hormone replacement and thence hypothyroidism if they take them simultaneously or near-simultaneously.

Excessive calcium supplementation can be detrimental to cardiovascular health, especially in men.

Strontium

In 1790, physician Adair Crawford, who had been working with barium, realized that Strontian ores showed different properties than other supposed ores of barium. Therefore, he concluded that these ores contained new minerals, which were named strontites in 1793 by Thomas Charles Hope, a chemistry professor at the University of Glasgow, who confirmed Crawford’s discovery. Strontium was eventually isolated in 1808 by Sir Humphry Davy by electrolysis of a mixture of strontium chloride and mercuric oxide. The discovery was announced by Davy on 30 June 1808 at a lecture to the Royal Society.

Applications

Consuming 75% of production, the primary use for strontium is in glass for colour television cathode ray tubes. It prevents X-ray emission. All parts of the CRT must absorb X-rays. In the neck and the funnel of the tube, lead glass is used for this purpose, but this type of glass shows a browning effect due to the interaction of the X-rays with the glass. Therefore, the front panel has to use a different glass mixture, in which strontium and barium are the X-ray-absorbing materials. The average values for the glass mixture determined for a recycling study in 2005 is 8.5% strontium oxide and 10% barium oxide. The amount of strontium used for the production of cathode ray tube is declining because the CRTs are replaced by other display methods. This decline has a significant influence on the mining and refining of strontium.

Because strontium is so similar to calcium, it is incorporated in the bone. All four stable isotopes are incorporated, in roughly similar proportions, as they are found in nature. However, the actual distribution of the isotopes tends to vary greatly from one geographical location to another. Thus, analyzing the bone of an individual can help determine the region it came from. This approach helps to identify the ancient migration patterns as well as the origin of commingled human remains in battlefield burial sites. Strontium, thus, helps forensic scientists too.

87Sr/86Sr ratios are commonly used to determine the likely provenance areas of sediment iatural systems, especially in marine and fluvial environments. Dasch (1969) showed that surface sediments of Atlantic displayed 87Sr/86Sr ratios that could be regarded as bulk averages of the 87Sr/86Sr ratios of geological terranes from adjacent landmasses. A good example of a fluvial-marine system to which Sr isotope provenance studies have been successfully employed is the River Nile-Mediterranean system, Due to the differing ages of the rocks that constitute the majority of the Blue and White Nile, catchment areas of the changing provenance of sediment reaching the River Nile delta and East Mediterranean Sea can be discerned through Sr isotopic studies. Such changes are climatically controlled in the Late Quaternary.

More recently, 87Sr/86Sr ratios have also been used to determine the source of ancient archaeological materials such as timbers and corn in Chaco Canyon, New Mexico. 87Sr/86Sr ratios in teeth may also be used to track animal migrations or in criminal forensics.

Pyrotechnics

Strontium carbonate or other strontium salts are used in the manufacture of fireworks, as they impart a deep red color to the firework. This application consumes about 5% of the world’s production.

Uses for radioactive strontium

89Sr is the active ingredient in Metastron (the generic version of Metastron, Generic Strontium Chloride Sr-89 Injection, its manufactured by Bio-Nucleonics Inc.), a radiopharmaceutical used for bone pain secondary to metastatic bone cancer. The strontium acts like calcium and is preferentially incorporated into bone at sites of increased osteogenesis. This localization focuses the radiation exposure on the cancerous lesion.

90Sr has been used as a power source for radioisotope thermoelectric generators (RTGs). 90Sr produces approximately 0.93 watts of heat per gram (it is lower for the form of 90Sr used in RTGs, which is strontium fluoride). However, 90Sr has a lifetime approximately 3 times shorter and has a lower density than 238Pu, another RTG fuel. The main advantage of 90Sr is that it is cheaper than 238Pu and is found iuclear waste. The Soviet Union deployed nearly 1000 of these RTGs on its northern coast as a power source for lighthouses and meteorology stations.

90Sr is also used in cancer therapy. Its beta emission and long half-life is ideal for superficial radiotherapy.

Niche applications

Strontium chloride is sometimes used in toothpastes for sensitive teeth. One popular brand includes 10% total strontium chloride hexahydrate by weight.

Small amounts are used in the refining of zinc, to remove small amounts of lead impurities.

Research trends

Strontium titanate has an extremely high refractive index and an optical dispersion greater than that of diamond, making it useful in a variety of optics applications. This quality has also led to its being cut into gemstones, in particular as a diamond simulant. However, it is very soft and easily scratches so it is rarely used.

Ferrite magnets

·          Strontium aluminate is used as a bright phosphor with long persistence of phosphorescence.

·          Strontium oxide is sometimes used to improve the quality of some pottery glazes.

·          Strontium ranelate is used in the treatment of osteoporosis. It is a prescription drug in the EU, but not in the USA.

·          Strontium barium niobate can be used in outdoors holographic 3D displays as a “screen”.

·          Strontium metal is used in strontium 90%-aluminium 10% alloys of an eutectic composition for the modification of aluminium-silicon casting alloys. AJ62, a durable, creep-resistant magnesium alloy used in car and motorcycle engines by BMW, contains 2% strontium by weight.

·          Strontium is used in scientific studies of neurotransmitter release ieurons. Like calcium, strontium facilitates synaptic vesicle fusion with the synaptic membrane. But, unlike calcium, strontium causes asynchronous vesicle fusion. Therefore, replacing calcium in a culture medium with strontium allows scientists to measure the effects of a single-vesicle fusion event, e.g., the size of the postsynaptic response elicited by the neurotransmitter content of a single vesicle.

The important concept for isotopic tracing is that Sr derived from any mineral through weathering reactions will have the same 87Sr/86Sr as the mineral. Therefore, differences in 87Sr/86Sr among ground waters require either (a) differences in mineralogy along contrasting flowpaths or (b) differences in the relative amounts of Sr weathered from the same suite of minerals. This latter situation can arise in several ways. First, differences in initial water chemistry within a homogeneous rock unit will affect the relative weathering rates of the minerals. For example, sections of the soil zone affected by evaporative concentration of recharge waters or by differences in pCO₂ can be expected to have different 87Sr/86Sr. Secondly, differences in the relative mobilities of water at scales ranging from inter-grain pores to the catchment scale may also profoundly affect 87Sr/86Sr. For example, the chemical composition and the resultant 87Sr/86Sr in immobile waters at a plagioclase-hornblende grain boundary versus a quartz-mica boundary will be different. Third, a difference in the relative “effective” surface areas of minerals in one portion of the rock unit will also cause differences in chemistry and isotopic composition; “poisoning” of reactive surfaces by organic coatings is an example of this kind of process. In a fundamental sense, because the waters in shallow systems are not in chemical equilibrium with the rocks, it is unrealistic to expect that waters along flowpaths within even a constant-mineralogy unit should have a constant 87Sr/86Sr. Instead, the waters moving along specific flowpaths slowly react with the rocks and gradually approach chemical equilibrium over long time-periods.

Compounds

The mineral celestine (SrSO₄) illustrates the fact that most Sr compounds are colourless or white.

Strontium forms a variety of salts, the properties of which are always intermediate between those of barium and calcium. The salts tend to be colourless. The sulfate and carbonate are poorly soluble, hence their occurrence as minerals. Most compounds are derived from the carbonate or the sulfide, which is obtained from the minerals. Typical for an alkaline earth derivative, the sulfide hydrolyzes readily:

SrS + 2 H₂O → Sr(OH)₂ + H₂S

Similar reactions are used in the production of commercially useful compounds, including the most useful strontium compound, strontium carbonate.

SrS + H₂O + CO₂ → SrCO₃ + H₂S

Strontium nitrate can also be prepared in this way.

The human body absorbs strontium as if it were calcium. Due to the chemical similarity of the elements, the stable forms of strontium might not pose a significant health threat — in fact, the levels found naturally may actually be beneficial – but the radioactive 90Sr can lead to various bone disorders and diseases, including bone cancer. The strontium unit is used in measuring radioactivity from absorbed 90Sr.

Acantharea a relative large group of marine radiolarian protozoa produce intricate mineral skeletons composed of strontium sulfate. In biological systems calcium is substituted in a small extent by strontium. In the human body most of the absorbed strontium is deposited in the bones. The ratio of strontium to calcium in human bones is between 1:1000 and 1:2000 roughly in the same range as in the blood serum.

A recent in-vitro study conducted the NY College of Dental Sciences using strontium on osteoblasts showed marked improvement on bone-building osteoblasts.

The drug strontium ranelate, made by combining strontium with ranelic acid, was found to aid bone growth, increase bone density, and lessen vertebral, peripheral, and hip fractures. Women receiving the drug showed a 12.7% increase in bone density. Women receiving a placebo had a 1.6% decrease. Half the increase in bone density (measured by X-ray densitometry) is attributed to the higher atomic weight of Sr compared with calcium, whereas the other half a true increase in bone mass. Strontium ranelate is registered as a prescription drug in Europe and many countries worldwide. It must be prescribed by a doctor, must be delivered by a pharmacist, and requires strict medical supervision.

There is a long history of medical research regarding strontium’s benefits, beginning in the 1950s. Studies indicate a lack of undesirable side-effects. Several other salts of strontium such as strontium citrate and strontium carbonate are available in the United States under the Dietary Supplements Health and Education Act of 1994, providing close to the recommended strontium content, about 680 milligrams per day, of strontium ranelate. Their long-term safety and efficacy have not been evaluated on humans in large-scale medical trials. However, some companies do manufacture strontium pills for increasing bone health.

Barium

Barite, the material in which barium was first known to be in

Barite, a mineral containing barium, was first recognized as containing a new element in 1774 by Carl Scheele, although he was only able to isolate barium oxide. Barium oxide was isolated again two years later by Johan Gottlieb Gahn. Later in the 18th century, William Withering noticed a heavy mineral in the Cumberland lead mines, which are now known to contain barium. Barium itself was finally isolated in 1808 when Sir Humphry Davy used electrolysis with molten salts, and Davy named the element barium, after baryta. Later, Robert Bunsen and Augustus Matthiessen isolated pure barium by electrolysis of a mixture of barium chloride and ammonium chloride.

Radium

While studying uraninite, on 21 December 1898, Marie and Pierre Curie discovered that even after uranium had decayed, the material created was still radioactive. The material behaved somewhat similarly to barium compounds, although some properties, such as the color of the flame test and spectral lines, were much different. They announced the discovery of a new element on 26 December 1898 to the French Academy of Sciences. Radium was named in 1899 from the word radius, meaning ray, as radium emitted power in the form of rays.

Occurrence

Series of alkaline earth metals

Beryllium is the chemical element with the symbol Be and atomic number 4. Because any beryllium synthesized in stars is short-lived, it is a relatively rare element in both the universe and in the crust of the Earth. It is a divalent element which occurs naturally only in combination with other elements in minerals. Notable gemstones which contain beryllium include beryl (aquamarine, emerald) and chrysoberyl. As a free element it is a steel-gray, strong, lightweight and brittle alkaline earth metal.

Beryllium increases hardness and resistance to corrosion when alloyed with aluminium, cobalt, copper (notably beryllium copper), iron and nickel. In structural applications, high flexural rigidity, thermal stability, thermal conductivity and low density (1.85 times that of water) make beryllium a quality aerospace material for high-speed aircraft, missiles, space vehicles and communication satellites. Because of its low density and atomic mass, beryllium is relatively transparent to X-rays and other forms of ionizing radiation; therefore, it is the most common window material for X-ray equipment and in particle physics experiments. The high thermal conductivities of beryllium and beryllium oxide have led to their use in heat transport and heat sinking applications.

The commercial use of beryllium presents technical challenges due to the toxicity (especially by inhalation) of beryllium-containing dusts. Beryllium is corrosive to tissue, and can cause a chronic life-threatening allergic disease called berylliosis in some people. The element is not known to be necessary or useful for either plant or animal life.

Beryllium occurs in the earth’s crust at a concentration of two to six parts per million (ppm), much of which is in soils, where it has a concentration of six ppm. Beryllium is one of the rarest elements in seawater, even rarer than elements such as scandium, with a concentration of 0.2 parts per trillion. However, in freshwater, beryllium is somewhat more common, with a concentration of 0.1 parts per billion.

Precautions

Approximately 35 micrograms of beryllium is found in the human body, but this amount is not considered harmful.  Beryllium is chemically similar to magnesium and therefore can displace it from enzymes, which causes them to malfunction. Chronic berylliosis is a pulmonary and systemic granulomatous disease caused by inhalation of dust or fumes contaminated with beryllium; either large amounts over a short time or small amounts over a long time can lead to this ailment. Symptoms of the disease can take up to five years to develop; about a third of patients with it die and the survivors are left disabled. The International Agency for Research on Cancer (IARC) lists beryllium and beryllium compounds as Category 1 carcinogens.

Acute beryllium disease in the form of chemical pneumonitis was first reported in Europe in 1933 and in the United States in 1943. A survey found that about 5% of workers in plants manufacturing fluorescent lamps in 1949 in the United States had beryllium-related lung diseases. Chronic berylliosis resembles sarcoidosis in many respects, and the differential diagnosis is often difficult. It killed some early workers iuclear weapons design, such as Herbert L. Anderson.

Early researchers tasted beryllium and its various compounds for sweetness in order to verify its presence. Modern diagnostic equipment no longer necessitates this highly risky procedure and no attempt should be made to ingest this highly toxic substance. Beryllium and its compounds should be handled with great care and special precautions must be taken when carrying out any activity which could result in the release of beryllium dust (lung cancer is a possible result of prolonged exposure to beryllium laden dust). Although the use of beryllium compounds in fluorescent lighting tubes was discontinued in 1949, potential for exposure to beryllium exists in the nuclear and aerospace industries and in the refining of beryllium metal and melting of beryllium-containing alloys, the manufacturing of electronic devices, and the handling of other beryllium-containing material.

A successful test for beryllium in air and on surfaces has been recently developed and published as an international voluntary consensus standard ASTM D7202. The procedure uses dilute ammonium bifluoride for dissolution and fluorescence detection with beryllium bound to sulfonated hydroxybenzoquinoline, allowing up to 100 times more sensitive detection than the recommended limit for beryllium concentration in the workplace. Fluorescence increases with increasing beryllium concentration. The new procedure has been successfully tested on a variety of surfaces and is effective for the dissolution and ultratrace detection of refractory beryllium oxide and siliceous beryllium (ASTM D7458).

Magnesium and calcium are both incredibly abundant in the earth’s crust, with calcium being the fifth most abundant element, and magnesium the eighth. While none of the alkaline earth metals are ever found in their elemental state, magnesium and calcium are found in many rocks and minerals. Magnesium is often found in carnellite, magnesite, and dolomite, while calcium is often found in chalk, limestone, gypsum, and anhydrite.

Strontium is also incredibly common on earth, being the fifteenth most abundant element in the crust. Most strontium in the crust is in the minerals celestite and strontianite. Barium is slightly less common, and much of it is in the mineral barite.

Radium, being a decay product of uranium, is found in all uranium-bearing ores. Due to its relatively short half-life, no radium that was present when the earth was formed is still around today, and instead has all come from the gradual decay of the uranium.

Production

Emerald, a variety of beryl, is a naturally occurring compound of beryllium

Most beryllium is extracted from beryllium hydroxide. One way to create this substance is a sintering method, which is done by mixing beryl, sodium fluorosilicate, and soda at high temperatures, which forms sodium fluoroberyllate, aluminium oxide and silicon dioxide. An solution of sodium fluoroberyllate and sodium hydroxide in water is then used to form beryllium hydroxide by precipitation. Another method used is known as the melt method. In it, beryl is heated to high temperatures while in a powdered form, and is then cooled with water. It is then heated again slightly while in sulfuric acid, eventually yielding beryllium hydroxide. The created beryllium hydroxide from either method is then used to create beryllium fluoride and beryllium chloride through a somewhat long process. Electrolysis or heating of these compounds can then be used to obtain beryllium.

Strontium carbonate is generally extracted from the mineral celestite. This can be done through two methods: either leaching the celestite with sodium carbonate, or by a more complicated method involving coal.

Barium can be produced from barite ore. Once the ore has been mined, it has to be separated from quartz, sometimes by froth flotation methods, resulting in relatively pure barite. Carbon is then used to reduce the baryte into barium sulfide. The barium sulfide can then be dissolved with other elements to form other compounds, such as barium nitrate, which in turn can be thermally decompressed into barium oxide, which eventually can yield pure barium after a reaction with aluminium. The most important supplier of barium is China, which produces more than 50% of the world’s barium.

Applications

Beryllium is mostly used for military applications, but there are other uses of beryllium as well. In electronics, beryllium is used as a p-type dopant in some semiconductors, and beryllium oxide is used as a high-strength electrical insulator and heat conductor. Due to its light weight and other properties, beryllium is also used in mechanics when stiffness, light weight, and dimensional stability are required at wide temperature ranges.

Magnesium has many different uses. One of its most common uses was in industry, where it has many structural advantages over other materials such as aluminium, although this usage has fallen out of favor recently due to magnesium’s flammability. Magnesium is also often alloyed with aluminium or zinc to form materials with more desirable properties than any pure metal. Magnesium has many other uses in industrial applications, such as having a role in the production of iron and steel, and the production of titanium.

Calcium also has many uses. One of its uses is as a reducing agent in the separation of other metals form ore, such as uranium. It is also used in the production of the alloys of many metals, such as aluminium and copper alloys, and is also used to deoxidize alloys as well. Calcium also has a role in the making of cheese, mortars, and cement.

Strontium and barium do not have as many applications as the lighter alkaline earth metals, but still have uses. Strontium carbonate is often used in the manufacturing of red fireworks, and pure strontium is used in the study of neurotransmitter release ieurons. Barium has some use in vacuum tubes to remove gases, and barium sulfate has many uses in the petroleum industry, as well as other industries.

Due to its radioactivity, radium no longer has many applications, but it used to have many. Radium used to be used often in luminous paints, although this use was stopped after workers got sick. As people used to think that radioactivity was a good thing, radium used to be added to drinking water, toothpaste, and many other products, although they are also not used anymore due to their health effects. Radium is no longer even used for its radioactive properties, as there are more powerful and safer emitters than radium.

Biological role and precautions

Magnesium and calcium are ubiquitous and essential to all known living organisms. They are involved in more than one role, with, for example, magnesium or calcium ion pumps playing a role in some cellular processes, magnesium functioning as the active center in some enzymes, and calcium salts taking a structural role, most notably in bones.

Strontium plays an important role in marine aquatic life, especially hard corals, which use strontium to build their exoskeletons. It and barium have some uses in medicine, for example “barium meals” in radiographic imaging, whilst strontium compounds are employed in some toothpastes. Excessive amounts of strontium-90 are toxic due to its radioactivity.

Beryllium and radium, however, are toxic. Beryllium’s low aqueous solubility means it is rarely available to biological systems; it has no known role in living organisms, and when encountered by them, is usually highly toxic. Radium has a low availability and is highly radioactive, making it toxic to life.

Extensions

The next alkaline earth metal after radium is thought to be element 120, although this may not be true due to relativistic effects. The synthesis of element 120 was first attempted in March 2007, when a team at the Flerov Laboratory of Nuclear Reactions in Dubna bombarded plutonium-244 with iron-58 ions; however, no atoms were produced, leading to a limit of 400 fb for the cross-section at the energy studied. In April 2007, a team at the GSI attempted to create element 120 by bombarding uranium-238 with nickel-64, although no atoms were detected, leading to a limit of 1.6 pb for the reaction. Synthesis was again attempted at higher sensitivities, although no atoms were detected. Other reactions have been tried, although all have been met with failure.

The chemistry of element 120 is predicted to be closer to that of calcium or strontium instead of barium or radium. This is unusual as periodic trends would predict element 120 to be more reactive than barium and radium. This lowered reactivity is due to the expected energies of element 120’s valence electrons, increasing element 120’s ionization energy and decreasing the metallic and ionic radii.

Historical applications

Some of the few practical uses of radium are derived from its radioactive properties. More recently discovered radioisotopes, such as 60Co and 137Cs, are replacing radium in even these limited uses because several of these isotopes are more powerful emitters, safer to handle, and available in more concentrated form.

Luminescent paint

Radium was formerly used in self-luminous paints for watches, nuclear panels, aircraft switches, clocks, and instrument dials. A typical self-luminous watch that uses radium paint contains around 1 microgram of radium. In the mid-1920s, a lawsuit was filed against the United States Radium Corporation by five dying “Radium Girl” dial painters who had painted radium-based luminous paint on the dials of watches and clocks. The dial painters routinely licked their brushes to give them a fine point, thereby ingesting radium. Their exposure to radium caused serious health effects which included sores, anemia, and bone cancer. This is because radium is treated as calcium by the body, and deposited in the bones, where radioactivity degrades marrow and can mutate bone cells.

During the litigation, it was determined that the company’s scientists and management had taken considerable precautions to protect themselves from the effects of radiation, yet had not seen fit to protect their employees. Worse, for several years the companies had attempted to cover up the effects and avoid liability by insisting that the Radium Girls were instead suffering from syphilis. This complete disregard for employee welfare had a significant impact on the formulation of occupational disease labor law.

As a result of the lawsuit, the adverse effects of radioactivity became widely known, and radium-dial painters were instructed in proper safety precautions and provided with protective gear. In particular, dial painters no longer licked paint brushes to shape them (which caused some ingestion of radium salts). Radium was still used in dials as late as the 1960s, but there were no further injuries to dial painters. This highlighted that the harm to the Radium Girls could easily have been avoided.

From the 1960s the use of radium paint was discontinued. In many cases luminous dials were implemented with non-radioactive fluorescent materials excited by light; such devices glow in the dark after exposure to light, but the glow fades. Where indefinite self-luminosity in darkness was required, safer radioactive promethium paint was initially used, later replaced by tritium which continues to be used today. Tritium emits beta radiation which cannot penetrate the skin, rather than the penetrating gamma radiation of radium and is regarded as safer. It has a half-life of 12 years.

Clocks, watches, and instruments dating from the first half of the 20th century, often in military applications, may have been painted with radioactive luminous paint. They are usually no longer luminous; however, this is not due to radioactive decay of the radium (which has a half-life of 1600 years) but to the fluorescence of the zinc sulfide fluorescent medium being worn out by the radiation from the radium. The appearance of an often thick layer of green or yellowish brown paint in devices from this period suggests a radioactive hazard. The radiation dose from an intact device is relatively low and usually not an acute risk; but the paint is dangerous if released and inhaled or ingested.

Recreational use

Radium was once an additive in products such as toothpaste, hair creams, and even food items due to its supposed curative powers. Such products soon fell out of vogue and were prohibited by authorities in many countries after it was discovered they could have serious adverse health effects. (See, for instance, Radithor or Revigator types of “Radium water” or “Standard Radium Solution for Drinking”.) Spas featuring radium-rich water are still occasionally touted as beneficial, such as those in Misasa, Tottori, Japan. In the U.S., nasal radium irradiation was also administered to children to prevent middle-ear problems or enlarged tonsils from the late 1940s through the early 1970s.

Medical use

Radium (usually in the form of radium chloride) was used in medicine to produce radon gas which in turn was used as a cancer treatment; for example, several of these radon sources were used in Canada in the 1920s and 1930s. The isotope 223Ra (under the trade name Alpharadin) is currently under investigation for use in medicine as a cancer treatment of bone metastasis.

Howard Atwood Kelly, one of the founding physicians of Johns Hopkins Hospital, was a major pioneer in the medical use of radium to treat cancer. His first patient was his own aunt in 1904, who died shortly after surgery. Kelly was known to use excessive amounts of radium to treat various cancers and tumors. As a result, some of his patients died from high amounts of radium exposure. His method of radium application was inserting a radium capsule near the affected area then sewing the radium “points” directly to the tumor. This was the same method used to treat Henrietta Lacks, the host of the original HeLa cells, for cervical cancer.

Research

In 1909, the famous Rutherford experiment used radium as an alpha source to probe the atomic structure of gold. This experiment led to the Rutherford model of the atom and revolutionized the field of nuclear physics. When mixed with beryllium, it is a neutron source. This type of neutron source were for a long time the main source for neutrons in research.

Precautions

Radium is highly radioactive and its decay product, radon gas, is also radioactive. Since radium is chemically similar to calcium, it has the potential to cause great harm by replacing calcium in bones. Exposure to radium can cause cancer and other disorders, because radium and its decay product radon emit alpha particles upon their decay, which kill and mutate cells. At the time of the Manhattan Project in 1944, the “tolerance dose” for workers was set at 0.1 microgram of ingested radium.

Some of the biological effects of radium were apparent from the start. The first case of so-called “radium-dermatitis” was reported in 1900, only 2 years after the element’s discovery. The French physicist Antoine Becquerel carried a small ampoule of radium in his waistcoat pocket for 6 hours and reported that his skin became ulcerated. Marie Curie experimented with a tiny sample that she kept in contact with her skin for 10 hours, and noted that an ulcer appeared several days later. Handling of radium has been blamed for Curie’s death due to aplastic anemia. Stored radium should be ventilated to prevent accumulation of radon. Emitted energy from the decay of radium also ionizes gases, fogs photographic emulsions, and produces many other detrimental effects.

WATER HARDNESS

ORIGIN OF WATER “HARDNESS”

 

Carbon dioxide reacts with water to form carbonic acid (1) which at ordinary environmental pH exists mostly as bicarbonate ion (2). Microscopic marine organisms take this up as carbonate (4) to form calcite skeletons which, over millions of years, have built up extensive limestone deposits. Groundwaters, made slightly acidic by CO₂ (both that absorbed from the air and from the respiration of soil bacteria) dissolve the limestone (3), thereby acquiring calcium and bicarbonate ions and becoming “hard”. If the HCO₃– concentration is sufficiently great, the combination of processes (2) and (4) causes calcium carbonate (“lime scale”) to precipitate out on surfaces such as the insides of pipes. (Calcium bicarbonate itself does not form a solid, but always precipitates as CaCO₃.)

Mechanisms and theory

The key to understanding the mechanism behind hardness is understanding the metallic microstructure, or the structure and arrangement of the atoms at the atomic level. In fact, most important metallic properties critical to the manufacturing of today’s goods are determined by the microstructure of a material. At the atomic level, the atoms in a metal are arranged in an orderly three-dimensional array called a crystal lattice. In reality, however, a given specimen of a metal likely never contains a consistent single crystal lattice. A given sample of metal will contain many grains, with each grain having a fairly consistent array pattern. At an even smaller scale, each grain contains irregularities.

There are two types of irregularities at the grain level of the microstructure that are responsible for the hardness of the material. These irregularities are point defects and line defects. A point defect is an irregularity located at a single lattice site inside of the overall three-dimensional lattice of the grain. There are three main point defects. If there is an atom missing from the array, a vacancy defect is formed. If there is a different type of atom at the lattice site that should normally be occupied by a metal atom, a substitutional defect is formed. If there exists an atom in a site where there should normally not be, an interstitial defect is formed. This is possible because space exists between atoms in a crystal lattice. While point defects are irregularities at a single site in the crystal lattice, line defects are irregularities on a plane of atoms. Dislocations are a type of line defect involving the misalignment of these planes. In the case of an edge dislocation, a half plane of atoms is wedged between two planes of atoms. In the case of a screw dislocation two planes of atoms are offset with a helical array running between them.

Dislocations provide a mechanism for planes of atoms to slip and thus a method for plastic or permanent deformation. Planes of atoms can flip from one side of the dislocation to the other effectively allowing the dislocation to traverse through the material and the material to deform permanently. The movement allowed by these dislocations causes a decrease in the material’s hardness.

The way to inhibit the movement of planes of atoms, and thus make them harder, involves the interaction of dislocations with each other and interstitial atoms. When a dislocation intersects with a second dislocation, it cao longer traverse through the crystal lattice. The intersection of dislocations creates an anchor point and does not allow the planes of atoms to continue to slip over one another A dislocation can also be anchored by the interaction with interstitial atoms. If a dislocation comes in contact with two or more interstitial atoms, the slip of the planes will again be disrupted. The interstitial atoms create anchor points, or pinning points, in the same manner as intersecting dislocations.

By varying the presence of interstitial atoms and the density of dislocations, a particular metal’s hardness can be controlled. Although seemingly counter-intuitive, as the density of dislocations increases, there are more intersections created and consequently more anchor points. Similarly, as more interstitial atoms are added, more pinning points that impede the movements of dislocations are formed. As a result, the more anchor points added, the harder the material will become.

Conventional water softening

Most conventional water-softening devices depend on a process known as ion-exchange in which “hardness” ions trade places with sodium and chloride ions that are loosely bound to an ion-exchange resin or a zeolite (many zeolite minerals occur iature, but specialized ones are often made artificially.)

 

Ion exchange in water softening

The illustration depicts a negatively-charged zeolite to which [positive] sodium ions are attached. Calcium or magnesium ions in the water displace sodium ions, which are released into the water. In a similar way, positively-charged zeolites bind negatively-charged chloride ions (Cl–), which get displaced by bicarbonate ions in the water. As the zeolites become converted to their Ca²⁺ and HCO³– forms they gradually lose their effectiveness and must be regenerated. This is accomplished by passing a concentrated brine solution though them, causing the above reaction to be reversed. Herein lies one of the drawbacks of this process: most of the salt employed in the regeneration process gets flushed out of the system and and is usually released into the soil or drainage system— something that can have damaging consequences to the environment, especially in arid regions. For this reason, many jurisdications prohibit such release, and require users to dispose of the spent brine at an approved site or to use a commercial service company.

Riverside County CA water softener restrictions “Alternative” water softening methods

The great economic importance of water softening has created a large and thriving industry that utilizes a number of proven methods based on well established scientific principles. It has also unfortunately attracted a variety of operators offering technologies that are purported to be better, less expensive, easier to install, or “chemical-free”, but which have never been validated scientifically and whose principles of operation are largely unexplained by the known laws of chemistry. This does not mean that such schemes cannot work (after all, we can use theory to show that under idealized conditions, water caever boil and it caever rain!), but it should inspire a good degree of skepticism. Most of the statements supporting alternative water treatement methods come from those who have a commercial interest in these devices, they are not supported by credible and independently verifiable performance data, and the explanations they offer for how they work reveal such a weak understanding of basic chemistry on the part of their authors that it is difficult to have much confidence in them.

Some dubious water-treatment processes and products

Magnetic water treatment and related pseudoscience

“Catalytic” water treatment schemes

Against this, there is some anecdotal evidence that certain magnetic and electromagnetic devices can be effective in preventing scale formation in hard water systems. It is very difficult to judge such claims, which are almost never based on tests that are well enough described to allow others to evaluate them and to verify the results. While the lack of “scientific” evidence does not in itself invalidate a claim for the efficacy of a device, it should make one hesitate to accept it without some guarantee of performance.

 

Hard water is water that has high mineral content (in contrast with “soft water”).

Hard drinking water is generally not harmful to one’s health, but can pose serious problems in industrial settings, where water hardness is monitored to avoid costly breakdowns in boilers, cooling towers, and other equipment that handles water. In domestic settings, hard water is often indicated by a lack of suds formation when soap is agitated in water. Wherever water hardness is a concern, water softening is commonly used to reduce hard water’s adverse effects.Sources of hardness

Water’s hardness is determined by the concentration of multivalent cations in the water. Multivalent cations are cations (positively charged metal complexes) with a charge greater than 1+. Usually, the cations have the charge of 2+. Common cations found in hard water include Ca²⁺ and Mg²⁺. These ions enter a water supply by leaching from minerals within an aquifer. Common calcium-containing minerals are calcite and gypsum. A common magnesium mineral is dolomite (which also contains calcium). Rainwater and distilled water are soft, because they also contain few ions.

The following equilibrium reaction describes the dissolving/formation of calcium carbonate scales:

CaCO₃ + CO₂ + H₂O ⇋ Ca²⁺ + 2HCO₃⁻

Calcium carbonate scales formed in water-heating systems are called limescale.

Calcium and magnesium ions can sometimes be removed by water softeners.

Temporary hardness

Temporary hardness is a type of water hardness caused by the presence of dissolved bicarbonate minerals (calcium bicarbonate and magnesium bicarbonate). When dissolved these minerals yield calcium and magnesium cations (Ca²⁺, Mg²⁺) and carbonate and bicarbonate anions (CO₃^2-, HCO₃^–). The presence of the metal cations makes the water hard. However, unlike the permanent hardness caused by sulfate and chloride compounds, this “temporary” hardness can be reduced either by boiling the water, or by the addition of lime (calcium hydroxide) through the softening process of lime softening. Boiling promotes the formation of carbonate from the bicarbonate and precipitates calcium carbonate out of solution, leaving water that is softer upon cooling.

Permanent hardness

Permanent hardness is hardness (mineral content) that cannot be removed by boiling. When this is the case, it is usually caused by the presence of calcium sulfate and/or magnesium sulfates in the water, which precipitates out as the temperature increases. Ions causing permanent hardness of water can be removed using a water softener, or ion exchange column.

Total Permanent Hardness = Calcium Hardness + Magnesium Hardness

The calcium and magnesium hardness is the concentration of calcium and magnesium ions expressed as equivalent of calcium carbonate.

Total permanent water hardness expressed as equivalent of CaCO₃ can be calculated with the following formula: Total Permanent Hardness (CaCO₃) = 2.5(Ca⁺⁺) + 4.1(Mg⁺⁺⁾

Effects of hard water

With hard water, soap solutions form a white precipitate (soap scum) instead of producing lather. This effect arises because the 2+ ions destroy the surfactant properties of the soap by forming a solid precipitate (the soap scum). A major component of such scum is calcium stearate, which arises from sodium stearate, the main component of soap:

2 C₁₇H₃₅COO^– + Ca²⁺ → (C₁₇H₃₅COO)₂Ca

Hardness can thus be defined as the soap-consuming capacity of a water sample, or the capacity of precipitation of soap as a characteristic property of water that prevents the lathering of soap. Synthetic detergents do not form such scums.

A portion of the ancient Roman Eifel aqueduct in Germany

Hard water also forms deposits that clog plumbing. These deposits, called “scale“, are composed mainly of calcium carbonate (CaCO₃), magnesium hydroxide (Mg(OH)₂), and calcium sulfate (CaSO₄). Calcium and magnesium carbonates tend to be deposited as off-white solids on the surfaces of pipes and the surfaces of heat exchangers. This precipitation (formation of an insoluble solid) is principally caused by thermal decomposition of bi-carbonate ions but also happens to some extent even in the absence of such ions. The resulting build-up of scale restricts the flow of water in pipes. In boilers, the deposits impair the flow of heat into water, reducing the heating efficiency and allowing the metal boiler components to overheat. In a pressurized system, this overheating can lead to failure of the boiler. The damage caused by calcium carbonate deposits varies depending on the crystalline form, for example, calcite or aragonite.

The presence of ions in an electrolyte, in this case, hard water, can also lead to galvanic corrosion, in which one metal will preferentially corrode when in contact with another type of metal, when both are in contact with an electrolyte. The softening of hard water by ion exchange does not increase its corrosivity per se. Similarly, where lead plumbing is in use, softened water does not substantially increase plumbo-solvency.

In swimming pools, hard water is manifested by a turbid, or cloudy (milky), appearance to the water. Calcium and magnesium hydroxides are both soluble in water. The solubility of the hydroxides of the alkaline-earth metals to which calcium and magnesium belong (group 2 of the periodic table) increases moving down the column. Aqueous solutions of these metal hydroxides absorb carbon dioxide from the air, forming the insoluble carbonates, giving rise to the turbidity. This often results from the alkalinity (the hydroxide concentration) being excesively high (pH > 7.6). Hence, a common solution to the problem is to, while maintaining the chlorine concentration at the proper level, raise the acidity (lower the pH) by the addition of hydrochloric acid, the optimum value being in the range of 7.2 to 7.6.

Softening

For the reasons discussed above, it is often desirable to soften hard water. Most detergents contain ingredients that counteract the effects of hard water on the surfactants. For this reason, water softening is often unnecessary. Where softening is practiced, it is often recommended to soften only the water sent to domestic hot water systems so as to prevent or delay inefficiencies and damage due to scale formation in water heaters. A common method for water softening involves the use of ion exchange resins, which replace ions like Ca²⁺ by twice the number of monocations such as sodium or potassium ions.

Health considerations

The World Health Organization says that “there does not appear to be any convincing evidence that water hardness causes adverse health effects in humans”. In fact, the National Research Council has found that hard water can actually serve as a dietary supplement for calcium and magnesium.

Some studies have shown a weak inverse relationship between water hardness and cardiovascular disease in men, up to a level of 170 mg calcium carbonate per litre of water. The World Health Organization has reviewed the evidence and concluded the data were inadequate to allow for a recommendation for a level of hardness.

Recommendations have been made for the maximum and minimum levels of calcium (40–80 ppm) and magnesium (20–30 ppm) in drinking water, and a total hardness expressed as the sum of the calcium and magnesium concentrations of 2–4 mmol/L.

Other studies have shown weak correlations between cardiovascular health and water hardness.

Some studies correlate domestic hard water usage with increased eczema in children.

The Softened-Water Eczema Trial (SWET), a multicenter randomized controlled trial of ion-exchange softeners for treating childhood eczema, was undertaken in 2008. However, no meaningful difference in symptom relief was found between children with access to a home water softener and those without.

There are a large number of hardness testing methods available (e.g. Vickers, Brinell, Rockwell, Meyer and Leeb). Although it is impossible in many cases to give an exact conversion, it is possible to give an approximate material-specific comparison table e.g. for steels.

THE VICKERS HARDNESS test was developed in 1921 by Robert L. Smith and George E. Sandland at Vickers Ltd as an alternative to the Brinell method to measure the hardness of materials. The Vickers test is often easier to use than other hardness tests since the required calculations are independent of the size of the indenter, and the indenter can be used for all materials irrespective of hardness. The basic principle, as with all common measures of hardness, is to observe the questioned material’s ability to resist plastic deformation from a standard source. The Vickers test can be used for all metals and has one of the widest scales among hardness tests. The unit of hardness given by the test is known as the Vickers Pyramid Number (HV) or Diamond Pyramid Hardness (DPH). The hardness number can be converted into units of pascals, but should not be confused with a pressure, which also has units of pascals. The hardness number is determined by the load over the surface area of the indentation and not the area normal to the force, and is therefore not a pressure.

THE BRINELL SCALE characterizes the indentation hardness of materials through the scale of penetration of an indenter, loaded on a material test-piece. It is one of several definitions of hardness in materials science.

Proposed by Swedish engineer Johan August Brinell in 1900, it was the first widely used and standardised hardness test in engineering and metallurgy. The large size of indentation and possible damage to test-piece limits its usefulness.

The typical test uses a 10 millimetres (0.39 in) diameter steel ball as an indenter with a 3,000 kgf (29 kN; 6,600 lbf) force. For softer materials, a smaller force is used; for harder materials, a tungsten carbide ball is substituted for the steel ball. The indentation is measured and hardness calculated as:

where:

P = applied force (kgf)

D = diameter of indenter (mm)

d = diameter of indentation (mm)

The BHN can be converted into the ultimate tensile strength (UTS), although the relationship is dependent on the material, and therefore determined empirically. The relationship is based on Meyer’s index (n) from Meyer’s law. If Meyer’s index is less than 2.2 then the ratio of UTS to BHN is 0.36. If Meyer’s index is greater than 2.2, then the ratio increases.

BHN is designated by the most commonly used test standards (ASTM E10-12 and ISO 6506–1:2005) as HBW (H from hardness, B from brinell and W from the material of the indenter, tungsten (wolfram) carbide). In former standards HB or HBS were used to refer to measurements made with steel indenters.

HBW is calculated in both standards using the SI units as

where:

F = applied force (N)

D = diameter of indenter (mm)

d = diameter of indentation (mm)

When quoting a Brinell hardness number (BHN or more commonly HB), the conditions of the test used to obtain the number must be specified. The standard format for specifying tests can be seen in the example “HBW 10/3000”. “HBW” means that a tungsten carbide (from the chemical symbol for tungsten) ball indenter was used, as opposed to “HBS”, which means a hardened steel ball. The “10” is the ball diameter in millimeters. The “3000” is the force in kilograms force.

The hardness may also be shown as XXX HB YYD². The XXX is the force to apply (in kgf) on a material of type YY (5 for aluminum alloys, 10 for copper alloys, 30 for steels). Thus a typical steel hardness could be written: 250 HB 30D². It could be a maximum or a minimum.

THE ROCKWELL SCALE is a hardness scale based on indentation hardness of a material. The Rockwell test determines the hardness by measuring the depth of penetration of an indenter under a large load compared to the penetration made by a preload. There are different scales, denoted by a single letter, that use different loads or indenters. The result is a dimensionless number noted as HRA, where A is the scale letter.

When testing metals, indentation hardness correlates linearly with tensile strength. This important relation permits economically important nondestructive testing of bulk metal deliveries with lightweight, even portable equipment, such as hand-held Rockwell hardness testers

History

The differential depth hardness measurement was conceived in 1908 by a Viennese professor Paul Ludwik in his book Die Kegelprobe (crudely, “the cone trial”). The differential-depth method subtracted out the errors associated with the mechanical imperfections of the system, such as backlash and surface imperfections. The Brinell hardness test, invented in Sweden, was developed earlier—in 1900—but it was slow, not useful on fully hardened steel, and left too large an impression to be considered nondestructive.

Hugh M. Rockwell (1890–1957) and Stanley P. Rockwell (1886–1940) from Connecticut in the United States co-invented the “Rockwell hardness tester,” a differential-depth machine. They applied for a patent on July 15, 1914. The requirement for this tester was to quickly determine the effects of heat treatment on steel bearing races. The application was subsequently approved on February 11, 1919, and holds U.S. Patent 1,294,171. At the time of invention, both Hugh and Stanley Rockwell worked for the New Departure Manufacturing Co. of Bristol, CT. New Departure was a major ball bearing manufacturer which in 1916 became part of United Motors and, shortly thereafter, General Motors Corp.

After leaving the Connecticut company, Stanley Rockwell, then in Syracuse, NY, applied for an improvement to the original invention on September 11, 1919, which was approved on November 18, 1924. The new tester holds U.S. Patent 1,516,207. Rockwell moved to West Hartford, CT, and made an additional improvement in 1921. Stanley collaborated with instrument manufacturer Charles H. Wilson of the Wilson-Mauelen Company in 1920 to commercialize his invention and develop standardized testing machines. Stanley started a heat-treating firm circa 1923, the Stanley P. Rockwell Company, which still exists in Hartford, CT. The later-named Wilson Mechanical Instrument Company has changed ownership over the years, and was most recently acquired by Instron Corp. in 1993.

The determination of the Rockwell hardness of a material involves the application of a minor load followed by a major load, and theoting the depth of penetration from a dial, on which a harder material gives a higher number. The chief advantage of Rockwell hardness is its ability to display hardness values directly, thus obviating tedious calculations involved in other hardness measurement techniques.

It is typically used in engineering and metallurgy. Its commercial popularity arises from its speed, reliability, robustness, resolution and small area of indentation.

In order to get a reliable reading the thickness of the test-piece should be at least 10 times the depth of the indentation. Also, readings should be taken from a flat perpendicular surface, because convex surfaces give lower readings. A correction factor can be used if the hardness of a convex surface is to be measured.

·         Except for testing thin materials in accordance with A623 the steel indenter balls have been replaced by tungsten carbide balls of the varying diameters. When a ball indenter is used, the letter “W is used to indicate a tungsten/carbide ball was used, and the letter “S” indicates the use of a steel ball. E.g.: 70 HRCW indicates the reading was 70 in the Rockwell C scale using a Tungsten/Carbide indenter.

·         The superficial Rockwell scales use lower loads and shallower impressions on brittle and very thin materials. The 45N scale employs a 45-kgf load on a diamond cone-shaped Brale indenter, and can be used on dense ceramics. The 15T scale employs a 15-kgf load on a ¹⁄₁₆-inch-diameter (1.588 mm) hardened steel ball, and can be used on sheet metal.

Readings below HRC 20 are generally considered unreliable, as are readings much above HRB 100.

 

THE MEYER HARDNESS TEST is a rarely used hardness test based upon projected area of an impression. This is a more fundamental measurement of hardness than other hardness tests which are based on the surface area of an indentation. The principle behind the test is that the mean pressure required to test the material is the measurement of the hardness of the material. The mean pressure is calculated by dividing the load by the projected area of the indentation. The result is called the Meyer hardness, which has units of megapascals (MPa).

An advantage of the Meyer test is that it is less sensitive to the applied load, especially compared to the Brinell hardness test. For cold worked materials the Meyer hardness is relatively constant and independent of load, whereas for the Brinell hardness test it decreases with higher loads. For annealed materials the Meyer hardness increases continuously with load due to strain hardening.

Based on Meyer’s law hardness values from this test can be converted into Brinell hardness values, and vice-versa.

The Meyer hardness test was devised by Prof. Eugene Meyer of the Materials Testing Laboratory at the Imperial School of Technology, Charlottenburg, Germany, circa 1908.

Measurement

Hardness can be quantified by instrumental analysis. The total water hardness is the sum of the molar concentrations of Ca²⁺ and Mg²⁺, in mol/L or mmol/L units. Although water hardness usually measures only the total concentrations of calcium and magnesium (the two most prevalent divalent metal ions), iron, aluminium, and manganese can also be present at elevated levels in some locations. The presence of iron characteristically confers a brownish (rust-like) colour to the calcification, instead of white (the color of most of the other compounds).

Water hardness is ofteot expressed as a molar concentration, but rather in various units, such as degrees of general hardness (dGH), German degrees (°dH), parts per million (ppm, mg/L, or American degrees), grains per gallon (gpg), English degrees (°e, e, or °Clark), or French degrees (°F). The table below shows conversion factors between the various units.

The various alternative units represent an equivalent mass of calcium oxide (CaO) or calcium carbonate (CaCO₃) that, when dissolved in a unit volume of pure water, would result in the same total molar concentration of Mg²⁺ and Ca²⁺. The different conversion factors arise from the fact that equivalent masses of calcium oxide and calcium carbonates differ, and that different mass and volume units are used. The units are as follows:

·            Parts per million (ppm) is usually defined as 1 mg/L CaCO₃ (the definition used below). It is equivalent to mg/L without chemical compound specified, and to American degree.

·            Grains per Gallon (gpg) is defined as 1 grain (64.8 mg) of calcium carbonate per U.S. gallon (3.79 litres), or 17.118 ppm.

·            a mmol/L is equivalent to 100.09 mg/L CaCO₃ or 40.08 mg/L Ca²⁺.

·            A degree of General Hardness (dGH or ‘German degree (°dH, deutsche Härte)’ is defined as 10 mg/L CaO or 17.848 ppm.

·            A Clark degree (°Clark) or English degrees (°e or e) is defined as one grain (64.8 mg) of CaCO₃ per Imperial gallon (4.55 litres) of water, equivalent to 14.254 ppm.

·            A French degree (°F or f) is defined as 10 mg/L CaCO₃, equivalent to 10 ppm. The lowercase f is often used to prevent confusion with degrees Fahrenheit.

Hard/soft classification

Because it is the precise mixture of minerals dissolved in the water, together with the water’s pH and temperature, that determines the behavior of the hardness, a single-number scale does not adequately describe hardness. However, the United States Geological Survey uses the following classification into hard and soft water,

Indices

Several indices are used to describe the behaviour of calcium carbonate in water, oil, or gas mixtures.

Langelier Saturation Index (LSI)

The Langelier Saturation Index (sometimes Langelier Stability Index) is a calculated number used to predict the calcium carbonate stability of water. It indicates whether the water will precipitate, dissolve, or be in equilibrium with calcium carbonate. In 1936, Wilfred Langelier developed a method for predicting the pH at which water is saturated in calcium carbonate (called pHs). The LSI is expressed as the difference between the actual system pH and the saturation pH:

LSI = pH (measured) — pHs

·            For LSI > 0, water is super saturated and tends to precipitate a scale layer of CaCO₃.

·            For LSI = 0, water is saturated (in equilibrium) with CaCO₃. A scale layer of CaCO₃ is neither precipitated nor dissolved.

·            For LSI < 0, water is under saturated and tends to dissolve solid CaCO₃.

If the actual pH of the water is below the calculated saturation pH, the LSI is negative and the water has a very limited scaling potential. If the actual pH exceeds pHs, the LSI is positive, and being supersaturated with CaCO₃, the water has a tendency to form scale. At increasing positive index values, the scaling potential increases.

In practice, water with an LSI between -0.5 and +0.5 will not display enhanced mineral dissolving or scale forming properties. Water with an LSI below -0.5 tends to exhibit noticeably increased dissolving abilities while water with an LSI above +0.5 tends to exhibit noticeably increased scale forming properties.

It is also worth noting that the LSI is temperature sensitive. The LSI becomes more positive as the water temperature increases. This has particular implications in situations where well water is used. The temperature of the water when it first exits the well is often significantly lower than the temperature inside the building served by the well or at the laboratory where the LSI measurement is made. This increase in temperature can cause scaling, especially in cases such as hot water heaters. Conversely, systems that reduce water temperature will have less scaling.

Water Analysis:

pH = 7.5

TDS = 320 mg/L

Calcium = 150 mg/L (or ppm) as CaCO3

Alkalinity = 34 mg/L (or ppm) as CaCO3

LSI Formula: LSI = pH - pHs

pHs = (9.3 + A + B) - (C + D) where:

A = (Log10[TDS] - 1)/10 = 0.15

B = -13.12 x Log10(oC + 273) + 34.55 = 2.09 at 25°C and 1.09 at 82°C

C = Log10[Ca2+ as CaCO3] - 0.4 = 1.78         (Ca2+ as CaCO3 is also called Calcium Hardness and is calculated as = 2.5(Ca2+)

D = Log10[alkalinity as CaCO3] = 1.53

Ryznar Stability Index (RSI)

The Ryznar stability index (RSI) uses a database of scale thickness measurements in municipal water systems to predict the effect of water chemistry.

Ryznar saturation index (RSI) was developed from empirical observations of corrosion rates and film formation in steel mains. It is defined as:

RSI = 2 pHs – pH (measured)

·                     For 6,5 < RSI < 7 water is considered to be approximately at saturation equilibrium with calcium carbonate

·                     For RSI > 8 water is under saturated and, therefore, would tend to dissolve any existing solid CaCO₃

·                     For RSI < 6,5 water tends to be scale forming

MEASURING HARDNESS

There are three main types of hardness measurements: scratch, indentation, and rebound. Within each of these classes of measurement there are individual measurement scales. For practical reasons conversion tables are used to convert between one scale and another.

·        Scratch hardness

Scratch hardness is the measure of how resistant a sample is to fracture or permanent plastic deformation due to friction from a sharp object. The principle is that an object made of a harder material will scratch an object made of a softer material. When testing coatings, scratch hardness refers to the force necessary to cut through the film to the substrate. The most common test is Mohs scale, which is used in mineralogy. One tool to make this measurement is the sclerometer.

Another tool used to make these tests is the pocket hardness tester. This tool consists of a scale arm with graduated markings attached to a four wheeled carriage. A scratch tool with a sharp rim is mounted at a predetermined angle to the testing surface. In order to use it a weight of known mass is added to the scale arm at one of the graduated markings, the tool is then drawn across the test surface. The use of the weight and markings allows a known pressure to be applied without the need for complicated machinery.

·        Indentation hardness

Indentation hardness measures the resistance of a sample to material deformation due to a constant compression load from a sharp object; they are primarily used in engineering and metallurgy fields. The tests work on the basic premise of measuring the critical dimensions of an indentation left by a specifically dimensioned and loaded indenter.

Common indentation hardness scales are Rockwell, Vickers, Shore, and Brinell.

·        Rebound hardness

Rebound hardness, also known as dynamic hardness, measures the height of the “bounce” of a diamond-tipped hammer dropped from a fixed height onto a material. This type of hardness is related to elasticity. The device used to take this measurement is known as a scleroscope.

Two scales that measures rebound hardness are the Leeb rebound hardness test and Bennett hardness scale.

·        Hardening

There are five hardening processes: Hall-Petch strengthening, work hardening, solid solution strengthening, precipitation hardening, and martensitic transformation.

·        Physics

Diagram of a stress-strain curve, showing the relationship between stress (force applied per unit area) and strain or deformation of a ductile metal.

In solid mechanics, solids generally have three responses to force, depending on the amount of force and the type of material:

They exhibit elasticity—the ability to temporarily change shape, but return to the original shape when the pressure is removed. “Hardness” in the elastic range—a small temporary change in shape for a given force—is known as stiffness in the case of a given object, or a high elastic modulus in the case of a material.

They exhibit plasticity—the ability to permanently change shape in response to the force, but remain in one piece. The yield strength is the point at which elastic deformation gives way to plastic deformation. Deformation in the plastic range is non-linear, and is described by the stress-strain curve. This response produces the observed properties of scratch and indentation hardness, as described and measured in materials science. Some materials exhibit both elasticity and viscosity when undergoing plastic deformation; this is called viscoelasticity.

They fracture—split into two or more pieces.

Strength is a measure of the extent of a material’s elastic range, or elastic and plastic ranges together. This is quantified as compressive strength, shear strength, tensile strength depending on the direction of the forces involved. Ultimate strength is an engineering measure of the maximum load a part of a specific material and geometry can withstand.

Brittleness, in technical usage, is the tendency of a material to fracture with very little or no detectable deformation beforehand. Thus in technical terms, a material can be both brittle and strong. In everyday usage “brittleness” usually refers to the tendency to fracture under a small amount of force, which exhibits both brittleness and a lack of strength (in the technical sense). For perfectly brittle materials, yield strength and ultimate strength are the same, because they do not experience detectable plastic deformation. The opposite of brittleness is ductility.

The toughness of a material is the maximum amount of energy it can absorb before fracturing, which is different from the amount of force that can be applied. Toughness tends to be small for brittle materials, because elastic and plastic deformations allow materials to absorb large amounts of energy.

Hardness increases with decreasing particle size. This is known as the Hall-Petch relationship. However, below a critical grain-size, hardness decreases with decreasing grain size. This is known as the inverse Hall-Petch effect.

Hardness of a material to deformation is dependent on its microdurability or small-scale shear modulus in any direction, not to any rigidity or stiffness properties such as its bulk modulus or Young’s modulus. Stiffness is often confused for hardness. Some materials are stiffer than diamond (e.g. osmium) but are not harder, and are prone to spalling and flaking in squamose or acicular habits.

 

Puckorius Scaling Index (PSI)

The Puckorius Scaling Index (PSI) uses slightly different parameters to quantify the relationship between the saturation state of the water and the amount of limescale deposited.

Other indices

Other indices include the Larson-Skold Index, the Stiff-Davis Index, and the Oddo-Tomson Index.

Regional information

Hard water in Australia

Analysis of water hardness in major Australian cities by the Australian Water Association shows a range from very soft (Melbourne) to very hard (Adelaide). Total Hardness levels of calcium carbonate in ppm are: Canberra: 40; Melbourne: 10–26; Sydney: 39.4–60.1; Perth: 29–226; Brisbane: 100; Adelaide: 134–148;  Hobart: 5.8–34.4; Darwin: 31.

Hard water in Canada

Prairie provinces (mainly Saskatchewan and Manitoba) contain high quantities of calcium and magnesium, often as dolomite, which are readily soluble in the groundwater that contains high concentrations of trapped carbon dioxide from the last glaciation. In these parts of Canada, the total hardness in ppm of calcium carbonate equivalent frequently exceed 200 ppm, if groundwater is the only source of potable water. The west coast, by contrast, has unusually soft water, derived mainly from mountain lakes fed by glaciers and snowmelt.

Some typical values are: Montreal 116 ppm, Calgary 165 ppm, Regina 202 ppm, Saskatoon 160-180 ppm, Winnipeg 77 ppm, Toronto 121 ppm, Vancouver < 3 ppm, Charlottetown, PEI 140–150 ppm, Waterloo Region 400 ppm, Guelph 460 ppm.

Hard water in England and Wales

Information from the British Drinking Water Inspectorate shows that drinking water in England is generally considered to be ‘very hard’, with most areas of England, particularly east of a line between the Severn and Tees estuaries, exhibiting above 200 ppm for the calcium carbonate equivalent. Wales, Devon, Cornwall and parts of North-West England are softer water areas, and range from 0 to 200 ppm. In the brewing industry in England and Wales, water is often deliberately hardened with gypsum in the process of Burtonisation.

Generally water is mostly hard in urban areas of England where soft water sources are unavailable. A number of cities built water supply sources in the 18th century as the industrial revolution and urban population burgeoned. Manchester was a notable such city in North West England and its wealthy corporation built a number of reservoirs at Thirlmere and Haweswater in the Lake District to the north. There is no exposure to limestone or chalk in their headwaters and consequently the water quality in Manchester is rated as ‘very soft’. Similarly, tap water in Birmingham is also soft as it is sourced from the Elan Valley Reservoirs in Wales.

Hard water in the United States

More than 85% of American homes have hard water. The softest waters occur in parts of the New England, South Atlantic-Gulf, Pacific Northwest, and Hawaii regions. Moderately hard waters are common in many of the rivers of the Tennessee, Great Lakes, and Alaska regions. Hard and very hard waters are found in some of the streams in most of the regions throughout the country. The hardest waters (greater than 1,000 ppm) are in streams in Texas, New Mexico, Kansas, Arizona, and southern California.

 

Methods for determination of total hardness

 

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?

 

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!

·                     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

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

·                     www.youtube.com/watch?v=DFQPnHkQlZM

·                     http://chemteacher.chemeddl.org/services/chemteacher/index.php?option=com_content&view=article&id=101

·                     www.youtube.com/watch?v=zQjIr_jgqtE

 

Prepared by PhD Falfushynska H.