The materials to prepare students for practical lessons of inorganic chemistry

Lesson № 22

Theme: General characteristics of d-elements.

Types of chemical reactions with their participation

 

Common characteristic of d-elements

The d-block transition metals have great importance in our lives. They are building blocks for life and are found directly in the center of the periodic table. The d-block simply means that the elements’ d–orbitals are the last to get occupied according to the building-up principle. The transition metals give off electrons from their outer s orbital, but most can lose a multiple number of d orbital electrons. Because of this many of the d-block metals have multiple oxidatioumbers. A good example is copper which has two common oxidation states +1 and +2. This causes d-block metals to make great catalysts.

Transition metals, for the most part, are good conductors. They are also malleable, ductile, lustrous, and sliver-white in color. An exception to this would be copper, which is brownish red in color. Metals have another great characteristic, they easily mix. This is because all the d-block metals have about the same atomic size. This allows them to replace one another easily in a crystal lattice. When two or more metals mix, or replace one another, we call the new metal an alloy. Brass is a good example of an alloy, which comes from copper and zinc combined. These elements and alloys are fundamental for the existence of life, and also for its progression through time. The d-block metals, and some of it’s key alloys, shaped the Bronze Age, Iron Age, and most importantly the steel age. Now with the booming of technology and the aerospace industry, metals with high conductivity and large strength to weight ratios are at top demand. Without these precious, durable, and sometimes highly valued metals, life simply would not exist.

Transition metals are found everywhere on Earth in various amounts. Most are not found in a pure substance, but rather in compounds buried in the Earth’s crust. This means that we must extract the metal from the compound in one of two ways. One process is pyrometallurgical which is when you use extremely high temperatures. The other is hydrometallurgical if you used aqueous solution.

Sometimes it takes only one step but many times multiple steps are mandatory. For example, iron is found abundantly in two dominant ores in the Earth’s crust: hematite Fe₂ O₃ and magnetite Fe₃ O_4. The ore is put into a blast furnace with some coke and limestone. The limestone then decomposes to form calcium oxide and carbon dioxide. The calcium oxide helps remove the nonmetal oxide and amphoteric impurities from the ore. The mixture is all liquid so the denser molten iron floats on the bottom. The mixture on top is simply drawn off and you are left with pig iron. Pig iron is almost pure iron. It is contaminated with small amounts of carbon and silicon.

Some metals that are rare can be sold at extremely high prices, such as gold. Other metals are found right in front to you. That computer is full of transition metals. It has to have metals in it to send electrical currents. How about your chair, it has metal ball-bearings in the wheels. Or the pictures hanging on the wall, they are hanging by nails, which are made from metal alloys. Almost everything around you is made from transition metals. Titanium is a relatively new transition metal that is in high demand due to its light weight, great strength, and high temperature and corrosion resistance. It is used to make airplane bodies and engines. Other temperature resistant metals are used to make blast furnaces and high temperature technology that can withstand extreme temperature changes.

Transition metals have always been on Earth. They have helped humans evolve through time. When humans learned how to make bronze from copper and tin they started the Bronze Age. Then came the Iron Age when higher processing temperatures became available. With higher temperatures came iron reduction.Finally the age of industry, and with it, the demand for steel.

In today’s society transition metals are in their highest demand ever. Steel is used to make bridges, buildings, and even works of art. Almost all of the skyscrapers have steel skeletons. Steel caot only be used independently; it can be mixed with other compounds or elements, such as carbon to give certain effects. If you add less than .15% carbon the alloy is ductile like iron wire. If the percentage is between .15 – .25% the alloy is much stronger. This alloy is used to make cables, chains, and nails. If the percentage is between .20 – .60% the alloy is mostly used for girders, rails, and structural purposes. If the percentage is .61 – 1.5% it is considered high-carbon steel. This is used to make knives, razors, cutting tools, and drill bits. As you can tell it takes only small changes in the concentration of ingredients to make large changes in the characteristics of the alloy.

Metals are also the key ingredient in automobiles because of their strength, durability, and extreme resistance to heat and fire. Metals are used to make bicycles, electrical toothbrushes, wires, refrigerators, and anything else that has metal parts. Anything that needs electricity has metal components because metals are electrical conductors. Battery casings, scissors, and microwaves are a few more examples of objects that are made from metals.

The main problem with transition metals is their readiness to oxidize. When they oxidize the metals corrode and become brittle. This is easily overcome by simply covering up, so they don’t come in contact with oxygen. Iron is a good example of this because it is used to make car bodies. If they didn’t paint cars they would all rust and the iron car body would fall apart. Coincidently when your paint scratches off rust forms there and the rust will eventually become brittle and fall off. Not all metals form oxides and become brittle. For instance, Titanium is corrosive resistant because it forms a protective skin when the exterior is exposed to oxygen. When the exterior is exposed oxides are formed. Once this occurs no further oxidizing takes place because oxygen caot get past the already formed oxides. The last way to stop the metals from oxidation is by making alloys. Alloys of chromium, for example, have a higher corrosion resistance than that of most alloys; they are given the name, stainless (i.e. stainless steel).

Transition metals are used as catalysts in many ways. We use metal surfaces with oxides to make ammonia. This is the most economical way to produce ammonia, and is highly used in fertilizers. The metal surface can adsorb elements and compounds into itself. Once this occurs bonds break between elements so they adsorb into the metal. Since the elements can move around they end up colliding together with enough energy to form a bond between each other and break the adsorption bond. It is in this fashion that ammonia in produced. This is not the only way metals can be used as a catalyst. Many times transition metals can be used to simply speed up a reaction. This is used because it is often economically cheaper to add some metal rather than waste time waiting for the reaction to occur. An example of this would be the use of a vanadium oxidizing catalyst in the process the making sulfuric acid.

We also use transition elements in many other ways. They are the key to making different colored paints, photo reactive eye glasses, and mercury thermometers. Titanium is used to detect underwater sound. Barium titanate is piezoelectric, which means that it generates an electrical charge when it is mechanically distorted. When the sound wave hits the compound it mechanically vibrates generating an electrical signal. Some like iron, cobalt, and nickel produce a magnetic effect called ferromagnetism and are used for permanent magnets and magnetic devices. Ferromagnetism is an effect similar toparamagnetism which occurs when an element has an unpaired electron. The electrons in paramagnetism are spinning randomly causing a much weaker effect. With ferromagnetism, the unpaired electrons have aligned spins forming domains that survive even after the applied field is turned off. For this reason ferromagnetic materials are used in coating cassette tapes, computer disks, and other devices that use magnetic codes and signals. We also use them to turn sunlight into electricity. Things like copper indium diselenide (CIS), and cadmium telluride (CdTe) can be found in photovoltaic solar cells used to convert solar energy into heat energy and eventually it is converted into electrical energy.

Transition metals are also found in our bodies. Humans excrete about 1 mg of iron every day and must constantly have approximately three grams of iron in their bodies. The iron is mostly used as hemoglobin, which transports oxygen to the brain and muscles. Iron deficiency, or anemia, occurs when your body doesn’t have enough iron and causes one to become chronically tired. Cobalt is another transition metal our bodies need. It is a component of vitamin B₁₂ which humans need in their diet.

The transition elements have hundreds of responsibilities. They are key elements in life and evolution. Without iron, oxygen wouldn’t make it to the brain and life would not exist. The bronze, iron, and steel ages would never have happened leaving us in the Stone Age. Transition metals have become of utmost importance due to our every growing population and economy. Their demand will continue as long as life as we know it continues.

d – is called transitional elements. They are listed in the periodic system in the long periods between s – and p – elements. A characteristic feature of the transition elements is that their atoms are filled orbitals is external, as in the s – and p – elements, and the penultimate (d – items) layer.

The transition elements occupy the central part of the periodic table, bridging the gap between the active s-block metals of groups 1A and 2A on the left and the p-block metals, semimetals, and nonmetals of groups 3A–8A on the right

Each d subshell consists of five orbitals and can accommodate 10 electrons, so each transition series consists of 10 elements. The first series extends from scandium through zinc and includes many familiar metals, such as chromium, iron, and copper. The second series runs from yttrium through cadmium, and the third series runs from lanthanum through mercury. In addition, there is a fourth transition series made up of actinium through the recently discovered and as yet unnamed element 112.

Tucked into the periodic table between lanthanum (atomic number 57) and hafnium (atomic number 72) are the lanthanides. In this series of 14 metallic elements, the seven 4 f orbitals are progressively filled. Following actinium (atomic number 89) is a second series of 14 elements, the actinides, in which the 5f subshell is progressively filled. The lanthanides and actinides together comprise the f-block elements, or inner transition elements.

With a few minor exceptions, the electronic structure of transition metal atoms can be written as [ ]ns²(n-1)d^m, where the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals, the situation is reversed such that the s electrons have higher energy. Consequently, an ion such as Fe²⁺ has no s electrons: it has the electronic configuration [Ar]3d⁶ as compared with the configuration of the atom, [Ar]4s²3d⁶.

The elements of groups 3–12 are now generally recognized as transition metals, although the elements La-Lu and Ac-Lr and Group 12 attract different definitions from different authors.

1.       Many chemistry textbooks and printed periodic tables classify La and Ac as Group 3 elements and transition metals, since their atomic ground-state configurations are s²d¹ like Sc and Y. The elements Ce-Lu are considered as the “lanthanide” series (or “lanthanoid” according to IUPAC) and Th-Lr as the “actinide” series. The two series together are classified as f-block elements, or (in older sources) as “inner transition elements”.

2.       Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides. This classification is based on similaritites in chemical behaviour, and defines 15 elements in each of the two series even though they correspond to the filling of an f subshell which can only contain 14 electrons.

3.       A third classification defines the f-block elements as La-Yb and Ac-No, while placing Lu and Lr in Group 3. This is based on the aufbau principle (or Madelung rule) for filling electron subshells, in which 4f is filled before 5d (and 5f before 6d), so that the f subshell is actually full at Yb (and No) while Lu (and Lr) has an [ ]s²f¹⁴d¹ configuration. However La and Ac are exceptions to the Aufbau principle with electron configuration [ ]s²d¹ (not [ ]s²f¹ as the aufbau principle predicts) so it is not clear from atomic electron configurations whether La or Lu (Ac or Lr) should be considered a transition metal.

The transition metals iron and copper have been known since antiquity and have played an important role in the development of civilization. Iron, the main constituent of steel, is still important as a structural material. Worldwide production of steel amounts to some 800 million tons per year. Iewer technologies, other transition elements are useful. For example, the strong, lightweight metal titanium is a major component in modern jet aircraft. Transition metals are also used as heterogeneous catalysts in automobile catalytic converters and in the industrial synthesis of essential chemicals such as sulfuric acid, nitric acid, and ammonia.

The role of the transition elements in living systems is equally important. Iron is present in biomolecules such as hemoglobin, which transports oxygen from our lungs to other parts of the body. Cobalt is an essential component of vitamin B_12. Nickel, copper, and zinc are vital constituents of many enzymes, the large protein molecules that catalyze biochemical reactions.

Electron Configurations

Look at the electron configurations of potassium and calcium, the s-block elements immediately preceding the first transition series. These atoms have 4s valence electrons, but no d electrons.

The filling of the 3d subshell begins at atomic number 21 (scandium) and continues until the subshell is completely filled at atomic number 30 (zinc).

The valence electrons are generally considered to be those in the outermost shell because they are the ones that are involved in chemical bonding. For transition elements, however, both the (n-1)d and the ns electrons are involved in bonding and are considered valence electrons.

The filling of the 3d subshell generally proceeds according to Hund’s rule with one electron adding to each of the five 3d orbitals before a second electron adds to any one of them. There are just two exceptions to the  expected regular filling pattern, chromium and copper:

Electron configurations depend on both orbital energies and electron–electron repulsions. Consequently, it’s not always possible to predict configurations when two valence subshells have similar energies. It’s often found, however, that exceptions from the expected orbital filling pattern result in either half-filled or completely filled subshells. In the case of chromium, for example, the 3d and 4s subshells have similar energies. It’s evidently advantageous to shift one electron from the 4s to the 3d subshell, which decreases electron–electron repulsions and gives two halffilled subshells. Because each valence electron is in a separate orbital, the electron–electron repulsion that would otherwise occur between the two 4s electrons in the expected configuration is eliminated. Asimilar shift of one electron from 4s to 3d in copper gives a completely filled 3d subshell and a half-filled 4s subshell.

For transition metal cations, the valence s orbital is vacant, and all the valence electrons occupy the d orbitals. Iron, for example, which forms 2+ and 3+ cations, has the following valence electron configurations:

When a neutral atom loses one or more electrons, the remaining electrons are less shielded, and the effective nuclear charge (Z_eff) increases. Consequently, the remaining electrons are more strongly attracted to the nucleus, and their orbital energies decrease. It turns out that the 3d orbitals experience a steeper drop in energy with increasing Z_eff than does the 4s orbital, making the 3d orbitals in cations lower in energy than the 4s orbital. As a result, all the valence electrons occupy the 3d orbitals, and the 4s orbital is vacant.

Ieutral molecules and complex anions, the metal atom usually has a positive oxidation state. It therefore has a partial positive charge and a higher than that of the neutral atom. As a result, the 3d orbitals are again lower in energy than the 4s orbital, and so all the metal’s valence electrons occupy the d orbitals. The metal atom in bothVCl₄ and Mn O₄^2-, for example, has the valence configuration 3d¹. Electron configurations for atoms for transition elements are summarized in Table.

Electron shells filled in violation of Madelung’s rule

 

Period 4 after [Ar]			Period 5 after [Kr]			Period 6 after [Xe]			Period 7 after [Rn]
Scandium	21	4s2 3d1	Yttrium	39	5s2 4d1	Lutetium	71	6s2 4f14 5d1	Lawrencium	103	7s2 4f14 7p1?
Titanium	22	4s2 3d2	Zirconium	40	5s2 4d2	Hafnium	72	6s2 4f14 5d2	Rutherfordium	104	7s2 4f14 6d2?
Vanadium	23	4s2 3d3	Niobium	41	5s1 4d4	Tantalu	73	6s2 4f14 5d3	Dubnium	105	7s2 4f14 6d3?
Chromium	24	4s1 3d5	Molybdenum	42	5s1 4d5	Tungsten	74	6s2 4f14 5d4	Seaborgium	106	7s2 4f14 6d4?
Manganese	25	4s2 3d5	Technetium	43	5s2 4d5	Rheniu	75	6s2 4f14 5d5	Bohrium	107	7s2 4f14 6d5?
Iron	26	4s2 3d6	Ruthenium	44	5s1 4d7	Osmium	76	6s2 4f14 5d6	Hassium	108	7s2 4f14 6d6?
Cobalt	27	4s2 3d7	Rhodium	45	5s1 4d8	Iridium	77	6s2 4f14 5d7	Meitnerium	109	7s2 4f14 6d7?
Nickel *[8]	28	4s2 3d8 or 4s1 3d9	Palladium	46	4d10	Platinum	78	6s1 4f14 5d9	Darmstadtium	110	7s2 4f14 6d8?
Copper	29	4s1 3d10	Silver	47	5s14d10	Gold	79	6s1 4f14 5d10	Roentgenium	111	7s1 4f14 6d10?
Zinc	30	4s2 3d10	Cadmium	48	5s2 4d10	Mercury	80	6s2 4f14 5d10	Copernicium	112	7s2 4f14 6d10?

 

Transition elements are placed in groups of the periodic table designated 1B–8B because their valence electron configurations are similar to those of analogous elements in the main groups A–8A. Thus, copper in group 1B ([Ar] 3d¹⁰ 4s¹) and zinc in group 2B ([Ar] 3d¹⁰ 4s²) have valence electron configurations similar to those of potassium in group 1A ([Ar] 4s¹) and calcium in group 2A ([Ar] 4s²). Similarly, scandium in group 3B ([Ar] 3d¹ 4s²) through iron in group 8B ([Ar] 3d⁶ 4s²) have the same number of valence electrons as the p-block elements aluminum in group 3A ([Ar] 3s² 3p¹) through argon in group 8A ([Ar] 3s² 3p⁶). Cobalt ([Ar] 3d⁷ 4s²) and nickel ([Ar] 3d⁸4s²) are also assigned to group 8B although there are no main-group elements with 9 or 10 valence electrons.

Except zinc all the transition metals complex ions are colourless due to presence of unpaired electrons. The colour of ions can be explained on the basis of “Crystal field theory”. According to this theory ,the    bonding between ligands and a metal ion is electrostatic. The ligands surrounding the metal ion and  create an electrostatic field around its d-orbitals. This field split ‘5’ degenerated d-orbitals in to two  sets of different energies.

Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.

·  charge transfer transitions. An electron may jump from a predominantly ligand orbital to a predominantly metal orbital, giving rise to a ligand-to-metal charge-transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. For example, the colour of chromate, dichromate and permanganate ions is due to LMCT transitions. Another example is that mercuric iodide, HgI₂, is red because of a LMCT transition.

A metal-to ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily reduced.

·  d–d transitions. An electron jumps from one d-orbital to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using crystal field theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on Tanabe-Sugano diagrams.

In centrosymmetric complexes, such as octahedral complexes, d–d transitions are forbidden by the Laporte rule and only occur because of vibronic coupling in which a molecular vibration occurs together with a d-d transition. Tetrahedral complexes have somewhat more intense colour because mixing d and p orbitals is possible when there is no centre of symmetry, so transitions are not pure d-d transitions. The molar absorptivity (ε) of bands caused by d-d transitions are relatively low, roughly in the range 5-500 M⁻¹cm⁻¹ (where M = mol dm⁻³).^[13] Some d–d transitions are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese(II), which has a d⁵ configuration in which all five electron has parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. In fact many compounds of manganese(II) appear almost colourless. The spectrum of [Mn(H₂O)₆]²⁺ shows a maximum molar absorptivity of about 0.04 M⁻¹cm⁻¹ in the visible spectrum.

(i) A high energy pair (eg) of d_x2-d_y2 and d_z2.

(ii) A low energy trio (t2g) of d_xy, d_yz, d_zx.

In many cases difference of energy between two sets ‘eg‘ and ‘t2g‘ is equivalent to a wavelength in the    visible region. Thus absorbing visible light, an electron may be able to move from lower energy set t2g to    higher energy set eg. In doing so, some of the component wavelength of white light is removed, so the    remaining component wavelength of light reflected or transmitted shows the colour.

For example: Cu+2 (blue), V+3 (green), Co+2 (pink), Fe+2 (green), Fe+3 (yellow), Ti+3 (purple).

From left to right, aqueous solutions of: Co(NO₃)₂ (red); K₂Cr₂O₇ (orange); K₂CrO₄ (yellow); NiCl₂ (turquoise); CuSO₄ (blue); KMnO₄ (purple)

The formation of coloured compounds

Some common examples

The diagrams show aproximate colours for some common transition metal complex ions.

You will find these and others discussed if you follow links to individual metals from the transition metal menu (link at the bottom of the page).

Alternatively, you could explore the complex ions menu (follow the link in the help box which has just disappeared off the top of the screen).

The origin of colour in the transition metal ions

When white light passes through a solution of one of these ions, or is reflected off it, some colours in the light are absorbed. The colour you see is how your eye perceives what is left.

Attaching ligands to a metal ion has an effect on the energies of the d orbitals. Light is absorbed as electrons move between one d orbital and another. This is explained in detail on another page.

Note: You will find a detailed explanation of the origin of colour in complex ions and the factors which cause it to change by following this link. That page is on the part of the site dealing with complex ions.

Use the BACK button on your browser if you want to return to this page again.

Catalytic activity

Transition metals and their compounds are often good catalysts. A few of the more obvious cases are mentioned below, but you will find catalysis explored in detail elsewhere on the site (follow the link after the examples).

Transition metals and their compounds function as catalysts either because of their ability to change oxidation state or, in the case of the metals, to adsorb other substances on to their surface and activate them in the process. All this is expored in the main catalysis section.

Transition metals as catalysts

Iron in the Haber Process

The Haber Process combines hydrogen and nitrogen to make ammonia using an iron catalyst.

Contact Process

This reaction is part of the production of sulphuric acid. The reaction between sulphur dioxide and oxygen to form sulphur trioxide is catalysed by vanadium (V) oxide, V₂O₅. This is a heterogeneous system.

 

Hydrogenation (Reduction)

Raney nickel, which is very finely divided nickel powder, is a catalyst for the addition of hydrogen across C=C and C=C bonds. Finely divided platinum, and finely divided paladium also catalyse this sort of reaction, but are a bit expensive. This is a heterogeneous system.

The margarine industy uses the above type of reaction to convert unsaturated vegetable oils into higher metal point saturated fats – i.e. margarine.

Nickel in the hydrogenation of C=C bonds

This reaction is at the heart of the manufacture of margarine from vegetable oils.

However, the simplest example is the reaction between ethene and hydrogen in the presence of a nickel catalyst.

Decomposition

Manganese dioxide catalyses the decomposition of hydrogen peroxide. Lots of other things catalyse this reaction too, but MnO₂ is one of the best. This is a heterogeneous system.

Transition metal compounds as catalysts

Vanadium(V) oxide in the Contact Process

At the heart of the Contact Process is a reaction which converts sulphur dioxide into sulphur trioxide. Sulphur dioxide gas is passed together with air (as a source of oxygen) over a solid vanadium(V) oxide catalyst.

Iron ions in the reaction between persulphate ions and iodide ions

Persulphate ions (peroxodisulphate ions), S₂O₈^2-, are very powerful oxidising agents. Iodide ions are very easily oxidised to iodine. And yet the reaction between them in solution in water is very slow.

The reaction is catalysed by the presence of either iron(II) or iron(III) ions.

On more advanced courses, you may come across the following examples of transition metals acting as catalysts.

Catalytic Convertors

The pollutants NO and CO are removed from car exhaust fumes by passing the exhaust gases over platinum and rhodium to produce nitrogen and carbon dioxide. This is a heterogeneous system.

Decomposition

Cobalt (II) catalyses the decompositon of the chlorate(I) anion, in two steps, as it is easily converted to cobalt (III). This is a heterogeneous system.

Ziegler-Natta Polymerisation

Titanium tetrachloride in the presence of triethylaluminium polymerises ethene, CH₂=CH₂. This is a heterogeneous system. A similar system using vanadyl halides exists in the co-polymerisation of styrene, butadiene and either dicyclopentadiene or 1,4-hexadiene to form synthetic rubber.

Many other examples of transiton metals and their compounds acting as catalysts can be found in text books. Yes, text books. Do try reading them sometimes as they are very helpful.

The d-orbitals of a free transition metal atom or ion are degenerate (all have the same energy.) However, when transition metals form coordination complexes, the d-orbitals of the metal interact with the electron cloud of theligands in such a manner that the d-orbitals become non-degenerate (not all having the same energy.) The way in which the orbitals are split into different energy levels is dependent on the geometry of the complex. Crystal field theory can be used to predict the energies of the different d-orbitals, and how the d-electrons of a transition metal are distributed among them. When the d-level is not completely filled, it is possible to promote and electron from a lower energy d-orbital to a higher energy d-orbital by absorption of a photon of electromagnetic radiation having an appropriate energy. Electromagnetic radiations in the visible region of the spectrum often possess the appropriate energy for such transitions.

Seeing Color

The sensors in our eyes detect only those wavelengths in the visible portion of the electromagnetic spectrum.

Although visible light appears “white”, it is made up of a series of colors. White light consists of three primary colors (red, yellow and blue). These primary colors can be mixed to make three secondary colors (orange, green and violet).

Transition Metal Complexes

When light passes through a solution containing transition metal complexes, we see those wavelengths of light that are transmitted. The solutions of most octahedral Cu (II) complexes are blue. The visible spectrum for an aqueous solution of Cu (II), [Cu(H₂O₆]²⁺, shows that the absorption band spans the red-orange-yellow portion of the spectrum and green, blue and violet are transmitted.

The absorption band corresponds to the energy required to excite an electron from the t_2g level to the e_g level.

 

Recall, the energy possessed by a light wave is inversely proportional to its wavelength. The Cu(II) solution transmits relatively high energy waves and absorbs the low energy wavelengths. This indicates that the band gap between the two levels is relatively small for this ion in aqueous solution

d-Orbital Splitting

The magnitude of the splitting of the d-orbitals in a transition metal complex depends on three things:

• the geometry of the complex

• the oxidation state of the metal

• the nature of the ligands

 

Properties of Transition Elements

There are some trends in the properties of the transition elements shown in Table 1. and try to understand them in terms of electron configurations.

Metallic Properties

All the transition elements are metals. Like the metals of groups 1A and 2A, the transition metals are malleable, ductile, lustrous, and good conductors of heat and electricity. Silver has the highest electrical conductivity of any element at room temperature, with copper a close second. The transition metals are harder, have higher melting and boiling points, and are more dense than the group 1A and 2A metals, largely because the sharing of d, as well as s, electrons results in stronger bonding.

From left to right across the first transition series, melting points increase from 1541°C for Sc to a maximum of 1910°C for V in group 5B, then decrease to 420°C for Zn (Table 1, Figure.2). The second- and third-series elements exhibit a similar maximum in melting point, but at group 6B: 2623°C for Mo and 3422°C for W, the metal with the highest melting point. The melting points increase as the number of unpaired d electrons available for bonding increases and then decrease as the d electrons pair up and become less available for bonding. Zinc (3d¹⁰ 4s²) in which all the d and s electrons are paired, has a relatively low melting point (420°C) and mercury (4f¹⁴ 5d¹⁰ 6s²) is a liquid at room temperature (mp –39°C).

Figure 2. Relative melting points of the transition elements. Melting points reach a maximum value in the middle of each series.

Atomic Radii and Densities

Atomic radii are given in Figure 3. from left to right across a transition series, the atomic radii decrease, at first markedly and then more gradually after group 6 B. Toward the end of each series, the radii increase again. The decrease in radii with increasing atomic number occurs because the added d electrons only partially shield the added nuclear charge. As a result, the effective nuclear charge Z_eff increases. With increasing Z_eff, the electrons are more strongly attracted to the nucleus, and atomic size decreases. The upturn in radii toward the end of each series is probably due to more effective shielding and to increasing electron–electron repulsion as double occupation of the d orbitals is completed. In contrast to the large variation in radii for main-group elements, all transition metal atoms have quite similar radii, which accounts for their ability to blend together in forming alloys such as brass (mostly copper and zinc).

Figure 3. Atomic radii (in pm) of the transition elements

The radii decrease with increasing atomic number and then increase again toward the end of each transition series. Note that the second- and third-series transition elements have nearly identical radii.

The atomic radii of the second- and third-series transition elements from group 4B on are nearly identical, though we would expect an increase in size on adding an entire principal quantum shell of electrons. The small sizes of the thirdseries atoms are associated with what is called the lanthanide contraction, the general decrease in atomic radii of the f-block lanthanide elements between the second and third transition series (Figure.4).

Figure 4. Atomic radii (in pm) of the lanthanide elements. The radii generally decrease with increasing atomic number

The lanthanide contraction is due to the increase in effective nuclear charge with increasing atomic number as the 4f subshell is filled. By the end of the lanthanides, the size decrease due to a larger Z_eff almost exactly compensates for the expected size increase due to an added quantum shell of electrons. Consequently, atoms of the third transition series have radii very similar to those of the second transition series.

The densities of the transition metals are inversely related to their atomic radii (Figure.5). The densities initially increase from left to right across each transition series and then decrease toward the end of each series. Because the second and third-series elements have nearly the same atomic volume, the much heavier third-series elements have unusually high densities: 22.6 g/cm³ for osmium and 22.5 g/cm³ for iridium, the most dense elements.

Figure 5. Relative densities of the transition metals. Density initially increases across each series and then decreases.

 

Ionization Energies and Oxidation Potentials

Ionization energies generally increase from left to right across a transition series, though there are some irregularities, as indicated in Table.1 for the atoms of the first transition series. The general trend correlates with an increase in effective nuclear charge and a decrease in atomic radius. Table.2 lists standard potentials E° for oxidation of first-series transition metals. Note that these potentials are the negative of the corresponding standard reduction potentials. Except for copper, all the E° values are positive, which means that the solid metal is oxidized to its aqueous cation more readily than H₂ gas is oxidized to H⁺ (aq).

Table. Standard Potentials for Oxidation of First-Series Transition Metals

In other words, the first-series metals, except for copper, are stronger reducing agents than H₂ gas and can therefore be oxidized by the H⁺ ion in acids like HCl that lack an oxidizing anion.

Oxidation of copper requires a stronger oxidizing agent, such as HNO₃.

The standard potential for the oxidation of a metal is a composite property that depends on ΔG° for the sublimation of the metal, the ionization energies of the metal atom, and ΔG° for the hydration of the metal ion.

Nevertheless, the general trend in the E° values shown in Table.2 correlates with the general trend in the ionization energies in Table.1. The ease of oxidation of the metal decreases as the ionization energies increase across the transition series from Sc to Zn. (Only Mn and Zn deviate from the trend of decreasing E° values.) Thus, the so-called early transition metals, those on the left side of the d block (Sc through Mn), are oxidized most easily and are the strongest reducing agents.

 

Properties of Transition Metals

·         The Properties of Transition Metals are largely dependent on the electronic configuration of the electrons in the outer shell and in the penultimate outer shell.

·         The transition elements readily form alloys with themselves and with other elements (e.g. a copper-tin alloy is used for mirrors, brass is a copper-zinc alloy). Tungsten, is used to make tools and filaments in light bulbs.

·         The atomic size is fairly constant since the electrons in the outer most shells have similar environments.

·         The low ionisation potentials mean that the elements show variable valency states by loss of electrons from the s and 3d orbitals.

·         The elements in this group can have different oxidation states which makes them useful as catalysts.

·         Compounds of the transition elements can be paramagnetic (i.e. attracted by a magnetic field) or diamagnetic (i.e. not attracted by a magnetic field). Paramagnetism in the transition elements is caused by the presence of unpaired electrons in the d sub-orbital. Diamagnetism is characteristic of compounds where all the electrons are paired in the d sub-orbitals.

Physical Properties of Transition Metals

Apart from Copper, the transition metals are all white lustrous metals. They vary widely in abundance (e.g. Iron, Fe, and Titanium, Ti, are plentiful, Scandium, Sc, is rare). They have high melting points and high densities. This suggests that the electrons which enter the d orbitals are being used to bind the atoms together in the crystal lattice.

The transition elements form Complex Ions.

Examples of these compound ions include

(a) Ferrocyanide Ion, Fe(CN)6(4 -) in Potassium Ferrocyanide, K4Fe(CN)6, and,

(b) Chromate Ion, CrO4(2 -), in Barium Chromate, BaCrO4.

Most compounds of the transition metal are coloured. There are variations in colour for compounds of the same valency, and with different valency (oxidation) states. For example,

A Coordination Complex is a compound in which molecules or ions form coordinate bonds to a central metal atom or ion. The complex may contain positive ions, negative ions or neutral molecules. The formation of such coordination complexes is typical behaviour of transition metals.

Diamagnetism is magnetic characteristic of those transition metals where all the electrons in the d sub-orbitals of the atoms are paired, and results in these elements not attracted by a magnetic field.

The Sub-Group Ib Elements are the transition metals Copper, Silver and Gold.

The Group III Elements in the periodic table are divided into two Sub-Groups.

The Sub-Group IIIa Elements, which are also called the main group elements, are Boron, Aluminium, Gallium, Indium and Thallium. The main group is essentially a group of weak metals, which have two s-electrons and one p-electron in their valence shells. Unlike the elements found in Groups I and in Group II, the Group III elements have more than one oxidation state, being the first members of the p block.

The Sub-Group IIIb Elements are the transition metals Scandium, Yttrium, Lanthanum, which is generally classified with the Lanthanoids and Actinium, which is generally classified with the Actinoids.

The Group IV Elements in the periodic table are divided into two Sub-Groups.

The Sub-Group IVa Elements, which are also called the main group elements, are Carbon, Silicon, Germanium, Tin, and Lead. These elements have two s electrons and two p electrons in their valence shells. The p electrons give rise to a +2 oxidation state while the s and p electrons give rise to a +4 oxidation state. The +4 oxidation state is formed due to an s electron being promoted to a vacant p orbital when the atom attains an excited state.

The Sub-Group IVb Elements consists of the three elements Titanium, Zirconium, and Hafnium. These three elements are generally classified with the transition elements.

The Group V Elements in the periodic table are divided into two Sub-Groups.

The Sub-Group Va Elements, which are also called the main group elements, are Nitrogen, Phosphorus, Arsenic, Antimony, and Bismuth. There are five electrons in the valence shells of these elements (i.e. two s orbital electrons and three p orbital). Thus, the principal oxidation state of these elements are +3 and +5 for the main group elements.

The Sub-Group Vb Elements are Vanadium, Niobium, and Tantalum. These three elements are generally classified with the transition elements.

The Group VI Elements in the periodic table are divided into two Sub-Groups.

The Sub-Group VIa Elements, which are also called the main group elements, are Oxygen, Sulphur, Selenium, Tellurium, and Polonium. These elements have two electrons short of the electronic configuration of a Nobel Gas. Thus, the oxidation states +2, +4, and +6 are found in the elements of this group.

The Sub-Group VIb Elements are Chromium, Molybdenum and Tungsten. These three elements are generally classified with the transition elements.

The Group VII Elements in the periodic table are divided into two Sub-Groups.

The Sub-Group VIIa Elements, which are also known as the Halogens (i.e. the main group elements), are Fluorine, Chlorine, Bromine, Iodine and Astatine. These elements are non-metals and are too reactive to occur on their own iature. Thus, they are usually found combined with other elements in a salt.

The Sub-Group VIIb Elements are the metallic transition elements Manganese, Technetium, and Rhenium. These three elements are generally classified with the transition elements.

The Group VIII Elements in the periodic table are a group of nine transition elements consisting of the platinum metals. This group consists of three triads of elements Iron, Ruthenium, Osmium, Cobalt, Rhodium, Iridium, and Nickel, Palladium Platinum. These elements are generally classified with the transition elements.

Inner Transition Metals

The Inner Transition Metals are the series of elements

(a) from Cerium (Atomic Number 58) to Lutetium (Atomic Number 71), which are called the Lanthanoids, and

(b) from Thorium (Atomic Number 90) to Lawrencium (Atomic Number 103), which are called the Actinoids.

The inner transition elements are found between Group 2 and the transition elements in the fifth row of the periodic table. The transition elements are also known as the f-block elements.

These two series make up the f-block elements in the periodic table, and their chemical properties of the elements derive from the filling of the f atomic sub-orbitals. The electronic configuration of these elements are characterised as having full outer orbitals and full second outermost orbitals, while the second outermost orbitals are incompletely filled. Thus, in the case of the first inner transition metals series, the electronic configuration of the outermost and second outermost orbitals is 4s2 3d10, while the third outermost orbitals (i.e. the 4f level) are incompletely filled.

The Main Group Elements in the periodic table can be classified as either being metals or non-metals. The vertical columns in the table are called groups. Elements within the same group all have the same number of electrons in their outer shell. Group 1 are known as Alkali metals. Group II are known as the Alkaline earth metals and the group of elements between group II and III are known as the transition metals.

Paramagnetism in the transition elements is caused by the presence of unpaired electrons in the d sub-orbital, and results in these elements being attracted by a magnetic field. Families

 

·        Coinage Metals

The “coinage metals” are copper, silver, gold, and roentgenium. These elements are used for much more than just coins, and many other elements besides these are made into coins. Furthermore, roentgenium is radioactive with a half-life of 3.6 seconds, making it useless for commercial applications. Consequently, the “coinage metals” are more appropriately called Group 11 (IB) elements.

Copper, silver, and gold, although relatively rare (copper), rare (silver), or extremely rare (gold), are among the longest-known and most familiar elements. They are soft, shiny, dense metals resistant to corrosion and very good conductors of electricity. Roentgenium, a recently-discovered synthetic element, is so short-lived that its physical and chemical properties are ill-defined.

Copper is by far the most heavily used of these elements due to its electrical properties, its commonness (contrasted to silver and gold) and the attractiveness of its alloys brass and bronze. Until aluminum became commonplace, copper was second only to iron in production among the metals.

They are easy to identify when found because copper (reddish) and gold (yellow) are the only two colored metals that people are likely to encounter. Silver is the shiniest of metals, and it is usually found in the presence of copper or gold and gives an obvious contrast. They are often found uncombined.

Because of their softness they are easily struck as coins, and their comparative rarity and attractiveness, along with their resistance to corrosion make them compact stores of wealth. They are too soft to have structural value, but copper alloys with such elements as zinc and tin to form harder brasses and bronzes. Brass and bronze were essential in the earliest metal tools; without them, civilization as we know would be impossible. Gold and silver, due to their attractiveness and their resistance to oxidation, have been used heavily in jewelry and other ornamental works. Gold, although extremely expensive, is so malleable that at modest cost a small amount can be pounded into a foil of extreme thinness that allows it to be used as a covering of some architectural objects; a little gold goes a long way.

Copper oxidizes with some difficulty to the +1 state in halides and an oxide and to the +2 state in salts such as copper sulfate CuSO4. Soluble copper compounds are easily identified by their distinctive blue-green color. Silver oxidizes to the +1 state in such substances as silver nitrate AgNO3 and silver sulfide Ag2S, the latter the typical blackening of silver. Gold oxidizes to the +1 and +3 state with great difficulty.

These elements are poor (copper) to extremely-poor (gold) reducing agents and their compounds are very good oxidizing agents. Copper ions oxidize most metals:

The reaction is even stronger with either silver or gold. In effect a solution of one of these metals’ salts plates most other metals.

Silver is most electrically conductive metal, followed by copper then gold. This makes copper a favorite material for electrical wires. Gold-tipped wires are employed in situations that need electrical precision (like high-quality audio) because gold will not tarnish (the tarnish of copper is much less conductive).

·        Zinc Family

The Zinc Family is Group 12 (IIB) and consists of zinc, cadmium, mercury, and copernicum.

Zinc, cadmium, and mercury are metals with low melting points for metals. This is because they have an especially stable electron configuration. Mercury is so poor at forming metallic bonds that it is liquid at room temperature.

Zinc and cadmium are soft metals that easily oxidize to the +2 oxidation state. Neither of these two metals appears uncombined iature. Zinc is heavily used in alloys with copper to create a harder metal known as brass; as a coating for iron (the process is called “galvanizing”), it oxidizes to form a protective layer of zinc oxide (ZnO) that protects the iron from oxidation, also known as rust. Zinc oxide is much safer than lead oxide and is often used in white paint. Since 1982, zinc has been the main metal used in American pennies. It is now used in new organ pipes.

Cadmium forms two substances, cadmium yellow (cadmium sulfide, CdS) and cadmium red (cadmium selenide, CdSe) that appeared in paints. These paints had strong colors that many of the great artists of the Impressionist periods cherished in their paintings. But these substances are very poisonous, and painters who used them often died young and crippled. Modern painters ordinarily use different paints that do not use these two poisonous chemicals.

Mercury, in contrast, is a shiny liquid at room temperature and oxidizes with some difficulty. It conducts electricity well. Because it is liquid it is an unusual metal—but it is a metal. It has been used in thermometers (but not so often after it has been identified as a dangerous poison) because it expands with heat and in switches where it can flow into a closed space to close a circuit. Mercury oxidizes to the +2 state in mercuric chloride (HgCl₂); in some strange compounds, two mercury atoms share an electron and offer their “spare” electrons to form substances in the +1 state, such as mercurous chloride (Hg₂Cl₂).

Zinc is an essential trace element for living things; it has some germicidal properties and is toxic (poisonous) in large quantities. Zinc pennies should never be swallowed. Cadmium, mercury, and their compounds are very dangerous poisons. Although mercury is attractive and has remarkable properties, it should be used with extreme care, and only by workers who have appropriate knowledge of its hazards.

The artificial Element 112 named copernicum in 2010 is probably part of this group in its properties, but it is extremely difficult to produce and too unstable to have a well-defined chemistry. Few atoms of this element have ever been made.

The elements of Groups 8, 9, and 10 are in two distinct groups: the common elements iron, cobalt, and nickel of the upper row of transition metals and the platinum metals of the second and third rows, and the far-scarcer platinum metals of the two lower rows of transition elements.

·        Iron, cobalt, and nickel

These elements are fairly-good reducing agents — so good that they rarely appear uncombined in nature. Iron is by far the most common of these. One of the most common elements in the universe, it is the heaviest metal that forms iormal fusion in stars (but only the largest stars). Once a star begins to produce iron in its core, that star is doomed in short order to a violent explosion that destroys the star and scatters its matter, including all of the elements that it has formed in fusion.

Uncombined iron, cobalt, and nickel — but especially iron — are to be found in meteors, solid objects that strike the earth. Iron is by far the most common of the transition elements, and one of the most useful. It’s hard to count all the uses of iron, the metal most used (whether pure or in alloys) in almost all machines. Giant “glass box” skyscrapers depend upon iron bars within their concrete “skeletons” to give them strength and stability. The rails of railroads are long iron bars. Concrete highways and airstrips have iron re-enforcing bars to give them the strength to hold heavy vehicles. The vehicles themselves are largely iron and a harder material known as steel, an alloy of iron, carbon, and often metals other than iron.

Iron is the cheapest of all structural metals. With some skill of an artisan known as a blacksmith it can be worked into many useful objects such as horseshoes, nails, plows, chains, pails, ladders, and many tools. In foundries, iron and steel are shaped in far greater quantities into such objects as furniture and parts of aircraft, ships, motor vehicles, and appliances.

Iron has one fault as a structural material: it rusts easily. In the presence of water (especially salt water) it corrodes into oxides:

Fe(s) + 1/2 O₂(g) → FeO(s) 2 Fe(s) + 3/2 O₂(g) → Fe₂O₃(s)

and a mixed oxide known as hematite

2 Fe(s) + 3/2 O₂(g) → Fe₂O₃(s)

one of the most common ores of iron. Iron oxides are mildly alkaline, so iron resists attacks by alkalis; acids attack it. For example,

Fe (s) + H₂SO₄(l) → Fe₂+(aq) + SO₄2-(aq) + H₂(g)

Even a comparatively weak acid, like phosphoric acid, can attack iron oxide. This is the “naval jelly” reaction that removes rust from iron:

FeO(s) + H₃PO₄(l) → Fe₂+(aq) + HPO₄(aq)-2 + H₂O(l)

A great advance of humanity, the beginning of the Iron Age, began when people found that they could separate iron from oxygen by burning it with carbon (usually charcoal) which can reduce iron oxides to iron:

Fe₃O₄(s) + 4 C(s) → 3 Fe(s) + 4 CO(g)

Much of existing economic activity depends upon the extraction of iron ore, the reduction of iron ore to iron, the strengthening of iron to steel, the creation of iron and steel objects, and the various practices used in protecting iron from corrosion.

Important as that activity is, our lives would be impossible without an important compound of iron known as hemoglobin which carries oxygen through the bloodstream to cells where the cells can use the oxygen to release energy from food also delivered to cells through the bloodstream.

The earth itself has a hot, dense core of largely iron and nickel. At the temperatures characteristic of the Earth’s core the iron and nickel form a giant natural magnet that creates a magnetic field that goes beyond the Earth itself into the atmosphere. That magnetic field drives off much dangerous radiation that would kill life on the Earth’s surface if it reached the Earth’s surface.

Cobalt and nickel are both far scarcer than iron and not as extensively used in commerce as iron, although they have specialized uses.

·        Platinum Family

The Platinum group metals are ruthenium, rhodium, palladium, osmium, iridium, and platinum. These elements are found in the second two rows of Groups 8/9/10 (IIIB).

Unlike their lighter counterparts in Groups 8, 9, and 10 of these elements are resistant to corrosion and tarnish. They serve as catalysts for many chemical reactions, speeding up the reaction without being consumed by it.

Palladium, osmium, and the other platinum group metals absorb hydrogen when powdered.

Rhodium is used in catalytic converters—metallic structures found inside vehicles. Catalytic converters convert nitric oxides (which are toxic pollutants) into elemental nitrogen and oxygen (both of which make up breathable air):

 That reaction would not occur without rhodium to serve as a catalyst.

 

OXIDATION STATES OF TRANSITION ELEMENTS

The transition elements differ from most main-group metals in that they exhibit a variety of oxidation states. Sodium, magnesium, and aluminum, for example, have a single oxidation state equal to their periodic group number (Na⁺, Mg²⁺, and Al³⁺), but the transition elements frequently have oxidation states less than their group number. For example, manganese in group 7B shows oxidation states of +2 in Mn²⁺ (aq), +3 in Mn(OH)₃ (s), +4 in MnO₂ (s), +6 in MnO₄^2- (aq) (manganate ion), and +7 in MnO₄^– (permanganate ion). Figure.6 summarizes the common oxidation states for elements of the first transition series, with the most frequently encountered ones indicated in red.

Common oxidation states for first-series transition elements. The states encountered most frequently are shown in red. The highest oxidation state for the group 3B–7B metals is their periodic group number, but the group 8B transition metals have a maximum oxidation state less than their group number. Most transition elements have more than one common oxidation state.

All the first-series transition elements except scandium form a 2+ cation, corresponding to loss of the two 4s valence electrons. Because the 3d and 4s orbitals have similar energies, loss of a 3d electron is also possible, yielding 3+ cations such as V³⁺ (aq), Cr³⁺(aq), and Fe³⁺ (aq). Additional energy is required to remove the third electron, but this is more than compensated for by the larger (more negative) ΔG° of hydration of the more highly charged 3+ cation. Still higher oxidation states result from loss or sharing of additional d electrons. In their highest oxidation states, the transition elements are combined with the most electronegative elements (F and O): for example, VF₅ (l) and V₂O₅(s) for vanadium in group 5B; CrO₄^2- (chromate ion) and Cr₂O₇^2- (dichromate ion) for chromium in group 6B; MnO₄^– for manganese in group 7B.

Note in Figure.6 that the highest oxidation state for the group 3B–7B metals is the group number, corresponding to loss or sharing of all the valence s and d electrons. For the group 8B transition metals, though, loss or sharing of all the valence electrons is energetically prohibitive because of the increasing value of Z_eff. Consequently, only lower oxidation states are accessible for these transition metals – for example, +6 in FeO₄^2- and +3 in Co³⁺. Even these species have a great tendency to be reduced to still lower oxidation states. For example, the aqueous Co³⁺ ion oxidizes water to O₂ gas and is thereby reduced to Co²⁺.

In general, ions that have the transition metal in a high oxidation state tend to be good oxidizing agents – for example, Cr₂O₇^2-, MnO₄^– and FeO₄^2-. Conversely, early transition metal ions with the metal in a low oxidation state are good reducing agents – for example, V²⁺ and Cr²⁺ Divalent ions of the later transition metals on the right side of the d block, such as Co²⁺, Ni²⁺, Cu²⁺ and Zn²⁺ are poor reducing agents because of the larger value of Z_eff. In fact, zinc has only one oxidation state (+2).

The elements of the second and third transition series also exhibit a variety of oxidation states. In general, the stability of the higher oxidation states increases down a periodic group. In group 8B, for example, the oxide of iron with the highest oxidation state is iron(III) oxide, Fe₂O₃. Ruthenium and osmium, though, form volatile tetroxides, RuO₄ and OsO₄ in which the metals have an oxidation state of +8.

Sockets d – elements usually painted. This is the transition of electrons to the high free energy level, which is due to absorption of visible light.

Varying Oxidation States

The following chart describes the most common oxidation states of the period 3 elements.

This diagram brings up a few concepts illustrating the stable states for specific elements.

Example 1.What makes zinc stable as Zn²⁺? What makes scandium stable as Sc³⁺?

Example 2. Why is iron almost always Fe²⁺ or Fe³⁺?

Example 3. Write manganese oxides in a few different oxidation states. Which ones are possible and/or reasonable?

Solution 1. Zinc has the neutral configuration [Ar]4s²3d¹⁰. Losing 2 electrons does not alter the complete d orbital. Neutral scandium is written as [Ar]4s²3d¹. Losing 3 electrons brings the configuration to the noble state with valence 3p⁶.

Solution 2. Iron is written as [Ar]4s²3d⁶. Losing 2 electrons from the s-orbital (3d⁶) or 2 s- and 1 d-orbital (3d⁵) electron are fairly stable oxidation states.

Solution 3. Although Mn⁺² is the most stable ion for manganese, the d-orbital can be made to remove 0 to 7 electrons. Compounds of manganese therefore range from Mn(0) as Mn_(s), Mn(II) as MnO, Mn(II,III) as Mn₃O₄, Mn(IV) as MnO₂, or manganese dioxide, Mn(VII) in the permanganate ion MnO₄^–, and so on.

 

Oxidative stress and transition metals. Metals such as iron, copper, chromium, vanadium, and cobalt are capable of redox cycling in which a single electron may be accepted or donated by the metal. This action catalyzes production of reactive radicals and reactive oxygen species. The presence of such metals in biological systems in an uncomplexed form (not in a protein or other protective metal complex) can significantly increase the level of oxidative stress. These metals are thought to induce Fenton reactions and the Haber-Weiss reaction, in which hydroxyl radical is generated from hydrogen peroxide. The hydroxyl radical then can modify amino acids. For example meta-tyrosine and ortho-tyrosine form by hydroxylation of phenylalanine. Other reactions include lipid peroxidation and oxidation of nucleobases. Metal catalyzed oxidations also lead to irreversible modification of R (Arg), K (Lys), P (Pro) and T (Thr) Excessive oxidative-damage leads to protein degradation or aggregation.

The reaction of transition metals with proteins oxidated by Reactive Oxygen Species or Reactive Nitrogen Species can yield reactive products that accumulate and contribute to aging and disease. For example, in Alzheimer’s patients, peroxidized lipids and proteins accumulate in lysosomes of the brain cells.

The Haber–Weiss reaction generates •OH (hydroxyl radicals) from H₂O₂ (hydrogen peroxide) and superoxide (•O₂^–). This reaction can occur in cells and is therefore a possible source foroxidative stress. The reaction is very slow, but is catalyzed by iron. The first step of the catalytic cycle involves reduction of ferric ion to ferrous:

Fe³⁺ + •O₂⁻ → Fe²⁺ + O₂

The second step is the Fenton reaction:

Fe²⁺ + H₂O₂ → Fe³⁺ + OH⁻ + •OH

Net reaction:

•O₂^– + H₂O₂ → •OH + OH^– + O₂

The reaction is named after Fritz Haber and his student Joseph Joshua Weiss.

 

The Fenton reaction has importance in biology because it involves the creation of free radicals by chemicals that are present in vivo. Transition-metal ions such as iron and copper donate or accept free electrons via intracellular reactions and help in creating free radicals. Although most intracellular iron is in ferrous (+2 ion) form, superoxide ions can convert it to the ferric (+3) form to take part in Fenton reaction. Since superoxide ions and transition metals act in a synergistic manner in the creation of free radical damage, Iron supplementation must not be done in patients with any active infections or in general any diseases.

H.J.H Fenton discovered in 1894 that several metals have a special oxygen transfer properties which improve the use of hydrogen peroxide. Actually, some metals have a strong catalytic power to generate highly reactive hydroxyl radicals (.OH). Since this discovery, the iron catalyzed hydrogen peroxide has been called Fenton’s reaction. Nowadays, the Fenton’s reaction is used to treat a large variety of water pollution such as phenols, formaldehyde, BTEX, pesticides, rubber chemicals and so on.

Applications of the Fenton’s reaction:

This process may be use to wastewater, contaminated soils and sludges with the following actions:

· Organic pollutant destruction

· Toxicity reduction

· Biodegradability improvement

· BOD/COD removal

· Odor and color removal

· Destruction of resin in radioactive contaminated sludge

How does the Fenton’s reaction work ?

After addition of the iron and the hydrogen peroxide, they are going to react together to generate some hydroxyl radicals as it shows in the following equations:

Fe²⁺ + H₂O₂ —-> Fe³⁺ + .OH + OH^–

Fe³⁺ + H₂O₂ —-> Fe²⁺ + .OOH + H⁺

The typical range for the iron dose is 1 part of Fe per 5-25 parts of H₂O₂.

After that the hydroxyl radicals are going to react with the pollutants to oxidize its. Actually the hydroxyl radicals can react according 4 kinds of reactions with the pollutants:

· Addition: ^.OH + C₆H₆ —-> (OH)C₆H₆

· Hydrogen Abstraction: OH + CH₃OH —-> CH₂OH + H₂O

· Electron Transfer: ^.OH + [Fe(CN)₆]^4- —-> [Fe(CN)₆]^3- + OH^–

· Radical Interaction: ^.OH + ^.OH —-> H₂O₂

During the Fenton’s reaction all the parameters are adjusted to promote the two first kind of reaction between the pollutant and the hydroxyl radicals.

Requirements of the reaction:

·  pH adjustment to 3-5 : if the pH is too high the iron precipitate in Fe(OH)₃ and will decompose the H₂O₂ to oxygen. Basically, the optimal pH occurs between 3 and 6. It’s really important to pay attention to the double pH drop due to the addition of Iron and H₂O₂ as you can see in the following chart. Indeed, FeSO₄ catalyst which contains residual H₂SO₄ and the H₂O₂ addition is responsible for the fragmentation of organic material into organic acids.

· addition of the iron catalyst as a solution of FeSO₄

· Adding slowly the H₂O₂ : in order to control the increasing of the pH and the temperature during the reaction it’s better to complete the reaction step by step with a continuous adjustment.

 

FIGURE. Generation of ^•OH by Fenton reaction (in red); ^•O₂ˉ in the mitochondria, peroxisomes and glyoxysomes and by Mehler reaction in chloroplast (in green), singlet oxygen in chloroplast (in dark green), and H₂O₂ by SOD, photorespiration, fatty acid oxidation or other reactions. SOD acts as the first line of defense converting ^•O₂ˉ into H₂O₂ (in yellow). CAT (in grey), POX (in pink), GPX (in dark blue), and APX (in orange) then detoxify H₂O₂. In contrast to CAT, APX requires ASC, POX requires phenolic compounds and/or ASC, and GPX requires GSH as electron donor substrate. In the removal of H₂O₂ through the ascorbate-glutathione cycle (in orange), ASC and GSH participate of the cyclic transfer of reducing equivalents. This cycle uses NADPH as reducing power. ^•OH may be removed by GSH (in blue), and the GSSG formed is regenerated via GR. Although the pathways of generation and scavenging in the different cell compartments are separate, H₂O₂ can easily diffuse through membranes and antioxidants such as GSH and ASC can be transported between the different compartments. Non-enzymatic pathways are indicated by dotted lines. Abbreviations: APX, ascorbate peroxidase; ASC, ascorbate; AH2, oxidizable substrate; DHA, dehydroascorbate; DHAR, dehydroascorbate reductase; GPX, glutathione peroxidase; POX, non-specific peroxidase; GR, glutathione reductase; GSH, reduced glutathione; GSSG, oxidized glutathione; hydrogen peroxide (H₂O₂); hydroxyl radical (^•OH); MDHA, monodehydroascorbate; MDHAR, monodehydroascorbate reductase SOD, superoxide dismutase; superoxide radical (^•O₂ˉ).

 

Practical skills

The electronic structures of transition metals

What is a transition metal?

The terms transition metal (or element) and d block element are sometimes used as if they mean the same thing. They don’t – there’s a subtle difference between the two terms.

We’ll explore d block elements first:

d block elements

You will remember that when you are building the Periodic Table and working out where to put the electrons using the Aufbau Principle, something odd happens after argon.

At argon, the 3s and 3p levels are full, but rather than fill up the 3d levels next, the 4s level fills instead to give potassium and then calcium.

Only after that do the 3d levels fill.

Note:  If you aren’t sure about atomic orbitals and electronic structures, you really need to follow this link before you go on. It takes you to a page explaining atomic orbitals and then on to other pages about electronic structures.

If you do follow the link, use the BACK button on your browser (or the History file or Go menu) to return quickly to this page.

The elements in the Periodic Table which correspond to the d levels filling are called d block elements. The first row of these is shown in the shortened form of the Periodic Table below.

The electronic structures of the d block elements shown are:

Sc      [Ar] 3d¹4s²

Ti      [Ar] 3d²4s²

V       [Ar] 3d³4s²

Cr     [Ar] 3d⁵4s¹

Mn     [Ar] 3d⁵4s²

Fe      [Ar] 3d⁶4s²

Co     [Ar] 3d⁷4s²

Ni      [Ar] 3d⁸4s²

Cu     [Ar] 3d¹⁰4s¹

Zn     [Ar] 3d¹⁰4s²

You will notice that the pattern of filling isn’t entirely tidy! It is broken at both chromium and copper.

Note:  This is something that you are just going to have to accept. There is no simple explanation for it which is usable at this level. Any simple explanation which is given is faulty!

People sometimes say that a half-filled d level as in chromium (with one electron in each orbital) is stable, and so it is – sometimes! But you then have to look at why it is stable. The obvious explanation is that chromium takes up this structure because separating the electrons minimises the repulsions between them – otherwise it would take up some quite different structure.

But you only have to look at the electronic configuration of tungsten (W) to see that this apparently simple explanation doesn’t always work. Tungsten has the same number of outer electrons as chromium, but its outer structure is different – 5d⁴6s². Again the electron repulsions must be minimised – otherwise it wouldn’t take up this configuration. But in this case, it isn’t true that the half-filled state is the most stable – it doesn’t seem very reasonable, but it’s a fact! The real explanation is going to be much more difficult than it seems at first sight.

Neither can you use the statement that a full d level (for example, in the copper case) is stable, unless you can come up with a proper explanation of why that is. You can’t assume that looking nice and tidy is a good enough reason!

If you can’t explain something properly, it is much better just to accept it than to make up faulty explanations which sound OK on the surface but don’t stand up to scrutiny!

 

Transition metals

Not all d block elements count as transition metals! There are discrepancies between the various UK-based syllabuses, but the majority use the definition:

A transition metal is one which forms one or more stable ions which have incompletely filled d orbitals.

Note:  The most recent IUPAC definition includes the possibility of the element itself having incomplete d orbitals as well. This is unlikely to be a big problem (it only really arises with scandium), but it would pay you to learn the version your syllabus wants. Both versions of the definition are currently in use in various UK-based syllabuses.

If you are working towards a UK-based exam and haven’t got a copy of your syllabus, follow this link to find out how to get one. Use the BACK button on your browser to return quickly to this page.

On the basis of the definition outlined above, scandium and zinc don’t count as transition metals – even though they are members of the d block.

Scandium has the electronic structure [Ar] 3d¹4s². When it forms ions, it always loses the 3 outer electrons and ends up with an argon structure. The Sc³⁺ ion has no d electrons and so doesn’t meet the definition.

Zinc has the electronic structure [Ar] 3d¹⁰4s². When it forms ions, it always loses the two 4s electrons to give a 2+ ion with the electronic structure [Ar] 3d¹⁰. The zinc ion has full d levels and doesn’t meet the definition either.

By contrast, copper, [Ar] 3d¹⁰4s¹, forms two ions. In the Cu⁺ ion the electronic structure is [Ar] 3d¹⁰. However, the more common Cu²⁺ ion has the structure [Ar] 3d⁹.

Copper is definitely a transition metal because the Cu²⁺ ion has an incomplete d level.

Transition metal ions

Here you are faced with one of the most irritating facts in chemistry at this level! When you work out the electronic structures of the first transition series (from scandium to zinc) using the Aufbau Principle, you do it on the basis that the 3d orbitals have a higher energy than the 4s orbital.

That means that you work on the assumption that the 3d electrons are added after the 4s ones.

However, in all the chemistry of the transition elements, the 4s orbital behaves as the outermost, highest energy orbital. When these metals form ions, the 4s electrons are always lost first.

You must remember this: When d-block elements form ions, the 4s electrons are lost first.

Note:  The problem here is that the Aufbau Principle can only really be used as a way of working out the electronic structures of most atoms. It is a simple way of doing that, although it fails with some, like chromium or copper, of course, and you have to learn these.

There is, however, a flaw in the theory behind it which produces problems like this. Why are the apparently higher energy 3d electrons not the ones to get lost when the metal ionises?

I have written a detailed explanation of this on another page called the order of filling 3d and 4s orbitals. If you are a teacher or a very confident student then you might like to follow this link.

If you aren’t so confident, I suggest that you ignore it. Make sure that you can work out the structures of these atoms using the Aufbau Principle on the assumption that the 3d orbitals fill after the 4s, and learn that when the atoms ionise, the 4s electrons are always lost first. Just ignore the contradictions between these two ideas!

 

To write the electronic structure for Co²⁺:

Co                        [Ar] 3d⁷4s²

Co²⁺                                   [Ar] 3d⁷

The 2+ ion is formed by the loss of the two 4s electrons.

To write the electronic structure for V³⁺:

V       [Ar] 3d³4s²

V³⁺ [Ar] 3d²

The 4s electrons are lost first followed by one of the 3d electrons.

Note:  You will find more examples of writing the electronic structures for d block ions, by following this link.

Use the BACK button on your browser to return quickly to this page.

Variable oxidation state (number)

One of the key features of transition metal chemistry is the wide range of oxidation states (oxidatioumbers) that the metals can show.

Note:  If you aren’t sure about oxidation states, you really need to follow this link before you go on.

Use the BACK button on your browser to return quickly to this page.

It would be wrong, though, to give the impression that onlytransition metals can have variable oxidation states. For example, elements like sulphur or nitrogen or chlorine have a very wide range of oxidation states in their compounds – and these obviously aren’t transition metals.

However, this variability is less common in metals apart from the transition elements. Of the familiar metals from the main groups of the Periodic Table, only lead and tin show variable oxidation state to any extent.

Examples of variable oxidation states in the transition metals

Iron

Iron has two common oxidation states (+2 and +3) in, for example, Fe²⁺ and Fe³⁺. It also has a less common +6 oxidation state in the ferrate(VI) ion, FeO₄^2-.

Manganese

Manganese has a very wide range of oxidation states in its compounds. For example:

 

Other examples

You will find the above examples and others looked at in detail if you explore the chemistry of individual metals from the transition metal menu. There is a link to this menu at the bottom of the page.

Explaining the variable oxidation states in the transition metals

We’ll look at the formation of simple ions like Fe²⁺ and Fe³⁺.

When a metal forms an ionic compound, the formula of the compound produced depends on the energetics of the process. On the whole, the compound formed is the one in which most energy is released. The more energy released, the more stable the compound.

There are several energy terms to think about, but the key ones are:

·         The amount of energy needed to ionise the metal (the sum of the various ionisation energies)

·         The amount of energy released when the compound forms. This will either be lattice enthalpy if you are thinking about solids, or the hydration enthalpies of the ions if you are thinking about solutions.

The more highly charged the ion, the more electrons you have to remove and the more ionisation energy you will have to provide.

But off-setting this, the more highly charged the ion, the more energy is released either as lattice enthalpy or the hydration enthalpy of the metal ion.

Note:  What I am talking about here in a general way are Born-Haber cycles. You will find these covered in theenergetics section of Chemguide, or my chemistry calculations book.

Thinking about a typical non-transition metal (calcium)

Calcium chloride is CaCl₂. Why is that?

If you tried to make CaCl, (containing a Ca⁺ ion), the overall process is slightly exothermic.

By making a Ca²⁺ ion instead, you have to supply more ionisation energy, but you get out lots more lattice energy. There is much more attraction between chloride ions and Ca²⁺ ions than there is if you only have a 1+ ion. The overall process is very exothermic.

Because the formation of CaCl₂ releases much more energy than making CaCl, then CaCl₂ is more stable – and so forms instead.

What about CaCl₃? This time you have to remove yet another electron from calcium.

The first two come from the 4s level. The third one comes from the 3p. That is much closer to the nucleus and therefore much more difficult to remove. There is a large jump in ionisation energy between the second and third electron removed.

Although there will be a gain in lattice enthalpy, it isn’t anything like enough to compensate for the extra ionisation energy, and the overall process is very endothermic.

It definitely isn’t energetically sensible to make CaCl₃!

Thinking about a typical transition metal (iron)

Here are the changes in the electronic structure of iron to make the 2+ or the 3+ ion.

Fe [Ar] 3d⁶4s²

Fe²⁺ [Ar] 3d⁶

Fe³⁺ [Ar] 3d⁵

The 4s orbital and the 3d orbitals have very similar energies. There isn’t a huge jump in the amount of energy you need to remove the third electron compared with the first and second.

The figures for the first three ionisation energies (in kJ mol^-1) for iron compared with those of calcium are:

There is an increase in ionisation energy as you take more electrons off an atom because you have the same number of protons attracting fewer electrons. However, there is much less increase when you take the third electron from iron than from calcium.

In the iron case, the extra ionisation energy is compensated more or less by the extra lattice enthalpy or hydration enthalpy evolved when the 3+ compound is made.

The net effect of all this is that the overall enthalpy change isn’t vastly different whether you make, say, FeCl₂ or FeCl₃. That means that it isn’t too difficult to convert between the two compounds.

USES OF THE TRANSITION METALS IN MEDICINE

The uses of the transition metals are wide and varied. They appear in almost every facet of our day to day life, from electric cabling to decorative door handles to corrosion-resistant alloys. Hence the transition metals are extremely important elements.

SCANDIUM – Sc

Scandium is a chemical element with symbol Sc and atomic number 21. A silvery-white metallic transition metal, it has historically been sometimes classified as a rare earth element, together with yttrium and the lanthanoids. It was discovered in 1879 by spectral analysis of the minerals euxenite and gadolinite from Scandinavia.

Scandium is present in most of the deposits of rare earth and uranium compounds, but it is extracted from these ores in only a few mines worldwide. Because of the low availability and the difficulties in the preparation of metallic scandium, which was first done in 1937, it took until the 1970s before applications for scandium were developed. The positive effects of scandium on aluminium alloys were discovered in the 1970s, and its use in such alloys remains its only major application. The global trade of the pure metal is around a hundred pounds a year on average.

The properties of scandium compounds are intermediate between those of aluminium and yttrium. A diagonal relationship exists between the behavior of magnesium and scandium, just as there is between beryllium and aluminium. In the chemical compounds of the elements shown as group 3, above, the predominant oxidation state is +3.

Applications

The addition of scandium to aluminium limits the excessive grain growth that occurs in the heat-affected zone of welded aluminium components. This has two beneficial effects: the precipitated Al3Sc forms smaller crystals than are formed in other aluminium alloys and the volume of precipitate-free zones that normally exist at the grain boundaries of age-hardening aluminium alloys is reduced. Both of these effects increase the usefulness of the alloy. However, titanium alloys, which are similar in lightness and strength, are cheaper and much more widely used.

The main application of scandium by weight is in aluminium-scandium alloys for minor aerospace industry components. These alloys contain between 0.1% and 0.5% of scandium. They were used in the Russian military aircraft, specifically the MiG-21 and MiG-29.

Some items of sports equipment, which rely on high performance materials, have been made with scandium-aluminium alloys, including baseball bats, and bicycle frames and components. Lacrosse sticks are also made with scandium. The American firearm manufacturing company Smith & Wesson produces revolvers with frames composed of scandium alloy and cylinders of titanium or carbon steel.

Dentists use Erbium, chromium: yttrium-scandium-gallium garnet (Er,Cr:YSGG) lasers for cavity preparation and in endodontics.

Approximately 20 kg (as Sc₂O₃) of scandium is used annually in the United States to make high-intensity discharge lamps.Scandium iodide, along with sodium iodide, when added to a modified form of mercury-vapor lamp, produces a form of metal halide lamp. This lamp is a white light source with high color rendering index that sufficiently resembles sunlight to allow good color-reproduction with TV cameras. About 80 kg of scandium is used in metal halide lamps/light bulbs globally per year. The first scandium-based metal halide lamps were patented by General Electric and initially made in North America, although they are now produced in all major industrialized countries. The radioactive isotope 46Sc is used in oil refineries as a tracing agent. Scandium triflate is a catalytic Lewis acid used in organic chemistry.

Scandium forms 25 ppm of the earth’s crust (cf the better known transition metals cobalt at 29 ppm and copper at 68 ppm). Scandium is very widely but thinly distributed (in over 800 minerals). The only mineral containing appreciable amounts of scandium is the rare thortveitite, Sc₂Si₂O₇, found in Norway, which contains 35-40% Sc₂O₃. But since the metal scandium has no significant uses, this lack of a good ore is no real problem. Scandium can be obtained as a by-product from the processing of uranium ores which contain only about 0.02% Sc₂O₃, but as large amounts of the uranium ore have to be processed anyway, this is a major source of Scandium. Scandium does have some uses – the iodide is added to mercury vapour lamps which the TV industry use because of the efficient light output and the similarity to the solar spectrum, while trace amount make aluminium alloys very strong (but they become four times as expensive). Isotopes of scandium are also used in crude oil analysis.

Elemental scandium is considered non-toxic and little animal testing of scandium compounds has been done. The median lethal dose (LD50) levels for scandium(III) chloride for rats have been determined as 4 mg/kg for intraperitoneal and 755 mg/kg for oral administration. In the light of these results compounds of scandium should be handled as compounds of moderate toxicity.

TITANIUM – Ti

Titanium is a chemical element with the symbol Ti and atomic number 22. It is a lustrous transition metal with a silver color, low density and high strength. It is highly resistant to corrosion in sea water, aqua regia and chlorine.

Titanium was discovered in Cornwall, Great Britain, by William Gregor in 1791 and named by Martin Heinrich Klaproth for the Titans of Greek mythology. The element occurs within a number of mineral deposits, principally rutile and ilmenite, which are widely distributed in the Earth’s crust and lithosphere, and it is found in almost all living things, rocks, water bodies, and soils. The metal is extracted from its principal mineral ores via the Kroll process or the Hunter process. Its most common compound, titanium dioxide, is a popular photocatalyst and is used in the manufacture of white pigments. Other compounds include titanium tetrachloride (TiCl₄), a component of smoke screens and catalysts; and titanium trichloride (TiCl₃), which is used as a catalyst in the production of polypropylene.

Titanium can be alloyed with iron, aluminium, vanadium, molybdenum, among other elements, to produce strong lightweight alloys for aerospace (jet engines, missiles, and spacecraft), military, industrial process (chemicals and petro-chemicals, desalination plants, pulp, and paper), automotive, agri-food, medical prostheses, orthopedic implants, dental and endodontic instruments and files, dental implants, sporting goods, jewelry, mobile phones, and other applications.

The two most useful properties of the metal are corrosion resistance and the highest strength-to-weight ratio of any metal. In its unalloyed condition, titanium is as strong as some steels, but 45% lighter. There are two allotropic forms and five naturally occurring isotopes of this element, 46Ti through 50Ti, with 48Ti being the most abundant (73.8%). Titanium’s properties are chemically and physically similar to zirconium, as both of them have the same number of valence electrons and are in the same group in the periodic table.

Physical properties

A metallic element, titanium is recognized for its high strength-to-weight ratio. It is a strong metal with low density that is quite ductile (especially in an oxygen-free environment), lustrous, and metallic-white in color. The relatively high melting point (more than 1,650 °C or 3,000 °F) makes it useful as a refractory metal. It is paramagnetic and has fairly low electrical and thermal conductivity.

Commercial (99.2% pure) grades of titanium have ultimate tensile strength of about 63,000 psi (434 MPa), equal to that of common, low-grade steel alloys, but are 45% lighter. Titanium is 60% more dense than aluminium, but more than twice as strong as the most commonly used 6061-T6 aluminium alloy. Certain titanium alloys (e.g., Beta C) achieve tensile strengths of over 200,000 psi (1,400 MPa). However, titanium loses strength when heated above 430 °C (806 °F).

Titanium is fairly hard (although not as hard as some grades of heat-treated steel), non-magnetic and a poor conductor of heat and electricity. Machining requires precautions, as the material will soften and gall if sharp tools and proper cooling methods are not used. Like those made from steel, titanium structures have a fatigue limit which guarantees longevity in some applications. Titanium alloys have lower specific stiffnesses than in many other structural materials such as aluminium alloys and carbon fiber.

The metal is a dimorphic allotrope whose hexagonal alpha form changes into a body-centered cubic (lattice) β form at 882 °C (1,620 °F). The specific heat of the alpha form increases dramatically as it is heated to this transition temperature but then falls and remains fairly constant for the β form regardless of temperature. Similar to zirconium and hafnium, an additional omega phase exists, which is thermodynamically stable at high pressures, but is metastable at ambient pressures. This phase is usually hexagonal (ideal) or trigonal (distorted) and can be viewed as being due to a soft longitudinal acoustic phonon of the β phase causing collapse of (111) planes of atoms.

Applications

Titanium is used in steel as an alloying element (ferro-titanium) to reduce grain size and as a deoxidizer, and in stainless steel to reduce carbon content. Titanium is often alloyed with aluminium (to refine grain size), vanadium, copper (to harden), iron, manganese, molybdenum, and with other metals. Applications for titanium mill products (sheet, plate, bar, wire, forgings, castings) can be found in industrial, aerospace, recreational, and emerging markets. Powdered titanium is used in pyrotechnics as a source of bright-burning particles.

Pigments, additives and coatings

About 95% of titanium ore extracted from the Earth is destined for refinement into titanium dioxide (TiO₂), an intensely white permanent pigment used in paints, paper, toothpaste, and plastics. It is also used in cement, in gemstones, as an optical opacifier in paper, and a strengthening agent in graphite composite fishing rods and golf clubs.

TiO₂ powder is chemically inert, resists fading in sunlight, and is very opaque: this allows it to impart a pure and brilliant white color to the brown or gray chemicals that form the majority of household plastics. Iature, this compound is found in the minerals anatase, brookite, and rutile. Paint made with titanium dioxide does well in severe temperatures, and stands up to marine environments. Pure titanium dioxide has a very high index of refraction and an optical dispersion higher than diamond. In addition to being a very important pigment, titanium dioxide is also used in sunscreens due to its ability to protect skin by itself. Recently, titanium oxide has been put to use in air purifiers (as a filter coating), or in film used to coat windows on buildings so that when titanium oxide becomes exposed to UV light (either solar or artificial) and moisture in the air, reactive redox species like hydroxyl radicals are produced so that they can purify the air or keep window surfaces clean.

Aerospace and marine

Due to their high tensile strength to density ratio, high corrosion resistance, fatigue resistance, high crack resistance, and ability to withstand moderately high temperatures without creeping, titanium alloys are used in aircraft, armor plating, naval ships, spacecraft, and missiles. For these applications titanium alloyed with aluminium, zirconium, nickel, vanadium, and other elements is used for a variety of components including critical structural parts, fire walls, landing gear, exhaust ducts (helicopters), and hydraulic systems. In fact, about two thirds of all titanium metal produced is used in aircraft engines and frames. The SR-71 “Blackbird” was one of the first aircraft to make extensive use of titanium within its structure, paving the way for its use in modern military and commercial aircraft. An estimated 59 metric tons (130,000 pounds) are used in the Boeing 777, 45 in the Boeing 747, 18 in the Boeing 737, 32 in the Airbus A340, 18 in the Airbus A330, and 12 in the Airbus A320. The Airbus A380 may use 77 metric tons, including about 11 tons in the engines. In engine applications, titanium is used for rotors, compressor blades, hydraulic system components, and nacelles. The titanium 6AL-4V alloy accounts for almost 50% of all alloys used in aircraft applications.

Due to its high corrosion resistance to sea water, titanium is used to make propeller shafts and rigging and in the heat exchangers of desalination plants; in heater-chillers for salt water aquariums, fishing line and leader, and for divers’ knives. Titanium is used to manufacture the housings and other components of ocean-deployed surveillance and monitoring devices for scientific and military use. The former Soviet Union developed techniques for making submarines with hulls of titanium alloys. Techniques were developed in the Soviet Union to forge titanium in huge vacuum tubes.

Industrial

Welded titanium pipe and process equipment (heat exchangers, tanks, process vessels, valves) are used in the chemical and petrochemical industries primarily for corrosion resistance. Specific alloys are used in downhole and nickel hydrometallurgy applications due to their high strength (e. g.: titanium Beta C alloy), corrosion resistance, or combination of both. The pulp and paper industry uses titanium in process equipment exposed to corrosive media such as sodium hypochlorite or wet chlorine gas (in the bleachery). Other applications include: ultrasonic welding, wave soldering, and sputtering targets.

Titanium tetrachloride (TiCl₄), a colorless liquid, is important as an intermediate in the process of making TiO₂ and is also used to produce the Ziegler-Natta catalyst, and is used to iridize glass and because it fumes strongly in moist air it is also used to make smoke screens.

Consumer and architectural

Titanium metal is used in automotive applications, particularly in automobile or motorcycle racing, where weight reduction is critical while maintaining high strength and rigidity. The metal is generally too expensive to make it marketable to the general consumer market, other than high-end products, particularly for the racing/performance market. Late model Corvettes have been available with titanium exhausts.

Titanium is used in many sporting goods: tennis rackets, golf clubs, lacrosse stick shafts; cricket, hockey, lacrosse, and football helmet grills; and bicycle frames and components. Although not a mainstream material for bicycle production, titanium bikes have been used by race teams and adventure cyclists.[65] Titanium alloys are also used in spectacle frames. This results in a rather expensive, but highly durable and long lasting frame which is light in weight and causes no skin allergies. Many backpackers use titanium equipment, including cookware, eating utensils, lanterns, and tent stakes. Though slightly more expensive than traditional steel or aluminium alternatives, these titanium products can be significantly lighter without compromising strength. Titanium is also favored for use by farriers, since it is lighter and more durable than steel when formed into horseshoes.

Titanium has occasionally been used in architectural applications: the 40 m (131 foot) memorial to Yuri Gagarin, the first man to travel in space, in Moscow, is made of titanium for the metal’s attractive color and association with rocketry. The Guggenheim Museum Bilbao and the Cerritos Millennium Library were the first buildings in Europe and North America, respectively, to be sheathed in titanium panels. Other construction uses of titanium sheathing include the Frederic C. Hamilton Building in Denver, Colorado and the 107 m (350 foot) Monument to the Conquerors of Space in Moscow.

Because of its superior strength and light weight when compared to other metals traditionally used in firearms (steel, stainless steel, and aluminium), and advances in metalworking techniques, the use of titanium has become more widespread in the manufacture of firearms. Primary uses include pistol frames and revolver cylinders. For these same reasons, it is also used in the body of laptop computers (for example, in Apple’s PowerBook line).

Some upmarket categories of tools made to be lightweight and corrosion-resistant, such as shovels and flashlights, are made of titanium or titanium alloys as well.

Jewelry

Because of its durability, titanium has become more popular for designer jewelry (particularly, titanium rings). Its inertness makes it a good choice for those with allergies or those who will be wearing the jewelry in environments such as swimming pools. Titanium is also alloyed with gold to produce an alloy that can be marketed as 24-carat gold, as the 1% of alloyed Ti is insufficient to require a lesser mark. The resulting alloy is roughly the hardness of 14-carat gold and thus is more durable than a pure 24-carat gold item would be.

Titanium’s durability, light weight, dent- and corrosion resistance makes it useful in the production of watch cases. Some artists work with titanium to produce artworks such as sculptures, decorative objects and furniture.

The inertness and ability to be attractively colored makes titanium a popular metal for use in body piercing. Titanium may be anodized to produce various colors, which varies the thickness of the surface oxide layer and causes interference fringes.

Medical

Because it is biocompatible (it is non-toxic and is not rejected by the body), titanium is used in a gamut of medical applications including surgical implements and implants, such as hip balls and sockets (joint replacement) that can stay in place for up to 20 years. The titanium is often alloyed with about 4% aluminium or 6% Al and 4% vanadium.

Titanium has the inherent ability to osseointegrate, enabling use in dental implants that can remain in place for over 30 years. This property is also useful for orthopedic implant applications. These benefit from titanium’s lower modulus of elasticity (Young’s modulus) to more closely match that of the bone that such devices are intended to repair. As a result, skeletal loads are more evenly shared between bone and implant, leading to a lower incidence of bone degradation due to stress shielding and periprosthetic bone fractures, which occur at the boundaries of orthopedic implants. However, titanium alloys’ stiffness is still more than twice that of bone, so adjacent bone bears a greatly reduced load and may deteriorate.

Since titanium is non-ferromagnetic, patients with titanium implants can be safely examined with magnetic resonance imaging (convenient for long-term implants). Preparing titanium for implantation in the body involves subjecting it to a high-temperature plasma arc which removes the surface atoms, exposing fresh titanium that is instantly oxidized.

Titanium is also used for the surgical instruments used in image-guided surgery, as well as wheelchairs, crutches, and any other products where high strength and low weight are desirable.

Titanium forms 0.63% (i.e. 6320 PPM) of the earth’s crust and is a very abundant element, 9th of all elements in fact, and second commonest of the transition elements. Titanium has a very low density for a metal, 57% of that of steel, and has good mechanical strength. When alloyed with small quantities of metals such as Al and Sn, the product has the highest strength to weight ratio of any of the engineering metals. Since about 1950 titanium has found uses in the manufacture of gas turbine engines which led to increased production and improvements in the production techniques. Its major uses are still in the aircraft industry, for airframes and engines, but chemical processing and marine equipment is also an important use. TiO₂ is the most widely used white pigment in the paint industry, over 3 millions tonnes being used annually. Titanium alloys are also used in hip replacement units and bone implants, due to the unreactive nature of the alloys. From the chemist’s point of view, titanium is important for its use as a component of the Ziegler Natta catalysts used in the polymerisation of alkenes. (Many transition metals find uses as catalysts.) World production of Titanium is about 100,000 tonnes p.a.

VANADIUM – V

Vanadium is a chemical element with the symbol V and atomic number 23. It is a hard, silvery gray, ductile and malleable transition metal. The element is found only in chemically combined form iature, but once isolated artificially, the formation of an oxide layer stabilizes the free metal somewhat against further oxidation.

Andrés Manuel del Río discovered compounds of vanadium in 1801 by analyzing a new lead-bearing mineral he called “brown lead,” and presumed its qualities were due to the presence of a new element, which he named erythronium (Greek for “red”) since, upon heating, most of its salts turned from their initial color to red. Four years later, however, he was (erroneously) convinced by other scientists that erythronium was identical to chromium. Chlorides of vanadium were generated in 1830 by Nils Gabriel Sefström who thereby proved that a new element was involved, which he named “vanadium” after the Germanic goddess of beauty and fertility, Vanadís (Freyja). Both names were attributed to the wide range of colors found in vanadium compounds. Del Rio’s lead mineral was later renamed vanadinite for its vanadium content. In 1867 Henry Enfield Roscoe obtained the pure element.

Vanadium occurs naturally in about 65 different minerals and in fossil fuel deposits. It is produced in China and Russia from steel smelter slag; other countries produce it either from the flue dust of heavy oil, or as a byproduct of uranium mining. It is mainly used to produce specialty steel alloys such as high speed tool steels. The most important industrial vanadium compound, vanadium pentoxide, is used as a catalyst for the production of sulfuric acid.

Large amounts of vanadium ions are found in a few organisms, possibly as a toxin. The oxide and some other salts of vanadium have moderate toxicity. Particularly in the ocean, vanadium is used by some life forms as an active center of enzymes, such as the vanadium bromoperoxidase of some ocean algae. Vanadium is probably a micronutrient in mammals, including humans, but its precise role in this regard is unknown.

Characteristics

Vanadium is a hard, ductile, silver-gray metal. Some sources describe vanadium as “soft”, perhaps because it is ductile, malleable and not brittle. Vanadium is harder than most metals and steels (see Hardnesses of the elements (data page) and iron). It has good resistance to corrosion and it is stable against alkalis, sulfuric and hydrochloric acids. It is oxidized in air at about 933 K (660 °C, 1220 °F), although an oxide layer forms even at room temperature.

Isotopes

Naturally occurring vanadium is composed of one stable isotope 51V and one radioactive isotope 50V. The latter has a half-life of 1.5×1017 years and a natural abundance 0.25%. 51V has a nuclear spin of 7/2 which is useful for NMR spectroscopy. 24 artificial radioisotopes have been characterized, ranging in mass number from 40 to 65. The most stable of these isotopes are 49V with a half-life of 330 days, and 48V with a half-life of 16.0 days. All of the remaining radioactive isotopes have half-lives shorter than an hour, most of which are below 10 seconds. At least 4 isotopes have metastable excited states. Electron capture is the main decay mode for isotopes lighter than the 51V. For the heavier ones, the most common mode is beta decay. The electron capture reactions lead to the formation of element 22 (titanium) isotopes, while for beta decay, it leads to element 24 (chromium) isotopes.

Compounds

The chemistry of vanadium is noteworthy for the accessibility of the four adjacent oxidation states 2-5. In aqueous solution, vanadium forms metal aquo complexes the colours are lilac [V(H₂O)₆]²⁺, green [V(H₂O)₆]³⁺, blue [VO(H₂O)₅]²⁺, yellow VO^3-. Vanadium (II) compounds are reducing agents, and vanadium(V) compounds are oxidizing agents.

Ammonium vanadate(V) (NH₄VO₃) can be successively reduced with elemental zinc to obtain the different colors of vanadium in these four oxidation states. Lower oxidation states occur in compounds such as V(CO)₆, [V(CO)₆]− and substituted derivatives.

The most commercially important compound is vanadium pentoxide. It is used as a catalyst for the production of sulfuric acid. This compound oxidizes sulfur dioxide (SO₂) to the trioxide (SO₃). In this redox reaction, sulfur is oxidized from +4 to +6, and vanadium is reduced from +5 to +4:

V₂O₅ + SO₂ → 2 VO₂ + SO₃

The catalyst is regenerated by oxidation with air:

2 VO₂ + O₂ → V₂O₅

Similar oxidations are used in the production of maleic anhydride, phthalic anhydride, and several other bulk organic compounds.

Applications

Alloys

Approximately 85% of vanadium produced is used as ferrovanadium or as a steel additive. The considerable increase of strength in steel containing small amounts of vanadium was discovered in the beginning of the 20th century. Vanadium forms stable nitrides and carbides, resulting in a significant increase in the strength of the steel. From that time on vanadium steel was used for applications in axles, bicycle frames, crankshafts, gears, and other critical components. There are two groups of vanadium containing steel alloy groups. Vanadium high-carbon steel alloys contain 0.15% to 0.25% vanadium and high speed tool steels (HSS) have a vanadium content of 1% to 5%. For high speed tool steels, a hardness above HRC 60 can be achieved. HSS steel is used in surgical instruments and tools. Some powder metallurgic alloys can contain up to 18% percent vanadium. The high content of vanadium carbides in those alloys increases the wear resistivity significantly. One application for those alloys are tools and knives.

Vanadium stabilizes the beta form of titanium and increases the strength and temperature stability of titanium. Mixed with aluminium in titanium alloys it is used in jet engines, high-speed airframes and dental implants. One of the common alloys is Titanium 6AL-4V, a titanium alloy with 6% aluminium and 4% vanadium.

Other uses

Vanadium is compatible with iron and titanium, therefore vanadium foil is used in cladding titanium to steel. The moderate thermal neutron-capture cross-section and the short half-life of the isotopes produced by neutron capture makes vanadium a suitable material for the inner structure of a fusion reactor. Several vanadium alloys show superconducting behavior. The first A15 phase superconductor was a vanadium compound, V3Si, which was discovered in 1952. Vanadium-gallium tape is used in superconducting magnets (17.5 teslas or 175,000 gauss). The structure of the superconducting A15 phase of V3Ga is similar to that of the more common Nb₃Sn and Nb₃Ti.

The most common oxide of vanadium, vanadium pentoxide V₂O₅, is used as a catalyst in manufacturing sulfuric acid by the contact process and as an oxidizer in maleic anhydride production. Vanadium pentoxide is also used in making ceramics. Another oxide of vanadium, vanadium dioxide VO₂, is used in the production of glass coatings, which blocks infrared radiation (and not visible light) at a specific temperature. Vanadium oxide can be used to induce color centers in corundum to create simulated alexandrite jewelry, although alexandrite iature is a chrysoberyl. The possibility to use vanadium redox couples in both half-cells, thereby eliminating the problem of cross contamination by diffusion of ions across the membrane is the advantage of vanadium redox rechargeable batteries. Vanadate can be used for protecting steel against rust and corrosion by electrochemical conversion coating. Lithium vanadium oxide has been proposed for use as a high energy density anode for lithium ion batteries, at 745 Wh/L when paired with a lithium cobalt oxide cathode. It has been proposed by some researchers that a small amount, 40 to 270 ppm, of vanadium in Wootz steel and Damascus steel, significantly improves the strength of the material, although it is unclear what the source of the vanadium was.

Biological role

Vanadium plays a very limited role in biology, and is more important in marine environments than terrestrial ones.

Vanadoenzymes

A number of species of marine algae produce vanadium-containing vanadium bromoperoxidase as well as the closely related chloroperoxidase and iodoperoxidases. The bromoperoxidase produce an estimated 1–2 million tons of bromoform and 56,000 tons of bromomethane annually. Most naturally occurring organobromine compounds, accounting arise by the action of this enzyme. They catalyse the following reaction (R-H is hydrocarbon substrate):

R-H + Br- + H2O2 → R-Br + H₂O + OH-

A vanadium nitrogenase is used by some nitrogen-fixing micro-organisms, such as Azotobacter. In this role vanadium replaces more common molybdenum or iron, and gives the nitrogenase slightly different properties.

Vanadium is essential to ascidians and tunicates, where it is stored in the highly acidified vacuoles of certain blood cell types, designated vanadocytes. Vanabins (vanadium binding proteins) have been identified in the cytoplasm of such cells. The concentration of vanadium in these ascidians’ blood is up to ten million times higher than the concentration of vanadium in the seawater around them, the seawater contains 1 to 2 µg/l. The function of this vanadium concentration system, and these vanadium-containing proteins, is still unknown but the vanadocytes are later deposited just under the outer surface of the tunic where their presence may deter predation.

Fungi

Several species of macrofungi, namely Amanita muscaria and related species, accumulate vanadium (up to 500 mg/kg in dry weight). Vanadium is present in the coordination complex amavadin, in fungal fruit-bodies. However, the biological importance of the accumulation process is unknown. Toxin functions or peroxidase enzyme functions have been suggested.

Mammals and birds

Deficiencies in vanadium result in reduced growth and impaired reproduction in rats and chickens. Vanadium is a relatively controversial dietary supplement, used primarily for increasing insulin sensitivity and body-building. Whether it works for the latter purpose has not been proven; some evidence suggests that athletes who take it are merely experiencing a placebo effect. Vanadyl sulfate may improve glucose control in people with type 2 diabetes. Decavanadate and oxovanadates appear to play a role in a variety of biochemical processes, such as those relating to oxidative stress.

Safety

All vanadium compounds should be considered toxic. Tetravalent VOSO₄ has been reported to be over 5 times more toxic than trivalent V₂O₃. The Occupational Safety and Health Administration (OSHA) has set an exposure limit of 0.05 mg/m3 for vanadium pentoxide dust and 0.1 mg/m3 for vanadium pentoxide fumes in workplace air for an 8-hour workday, 40-hour work week. The National Institute for Occupational Safety and Health (NIOSH) has recommended that 35 mg/m3 of vanadium be considered immediately dangerous to life and health. This is the exposure level of a chemical that is likely to cause permanent health problems or death.

Vanadium compounds are poorly absorbed through the gastrointestinal system. Inhalation exposures to vanadium and vanadium compounds result primarily in adverse effects on the respiratory system. Quantitative data are, however, insufficient to derive a subchronic or chronic inhalation reference dose. Other effects have been reported after oral or inhalation exposures on blood parameters, on liver, on neurological development in rats, and other organs.

There is little evidence that vanadium or vanadium compounds are reproductive toxins or teratogens. Vanadium pentoxide was reported to be carcinogenic in male rats and male and female mice by inhalation in an NTP study, although the interpretation of the results has recently been disputed. Vanadium has not been classified as to carcinogenicity by the United States Environmental Protection Agency.

Vanadium traces in diesel fuels present a corrosion hazard; it is the main fuel component influencing high temperature corrosion. During combustion, it oxidizes and reacts with sodium and sulfur, yielding vanadate compounds with melting points down to 530 °C, which attack the passivation layer on steel, rendering it susceptible to corrosion. The solid vanadium compounds also cause abrasion of engine components.

Vanadium represents about 136 PPM of the earth’s crust and is the 19th most abundant element. Over 60 different minerals contain vanadium. It is often recovered as a by product from other processes (such as uranium production) rather than directly produced. An unusual source is as a by-product from Venezuelan and Canadian crude oil. Vanadium has several important uses. It is added to some steels because it reacts with carbon to form V₄C₃ and disperses to produce a fine grained steel with resistance to wear and is stronger at high temperatures. These steels are widely used for producing springs and high speed tools. Vanadium(V) oxide is the catalyst in the production of SO₃ from SO₂ in the manufacture of sulphuric acid – the Contact process. (More tonnes of H₂SO₄ are manufactured annually than other compound.) Annual production of vanadium exceeds 35,000 tonnes p.a.

CHROMIUM – Cr

Chromium is a chemical element which has the symbol Cr and atomic number 24. It is the first element in Group 6. It is a steely-gray, lustrous, hard and brittle metal which takes a high polish, resists tarnishing, and has a high melting point. The name of the element is derived from the Greek word “chrōma” (χρώμα), meaning colour, because many of its compounds are intensely coloured.

Chromium oxide was used by the Chinese in the Qin dynasty over 2,000 years ago to coat metal weapons found with the Terracotta Army. Chromium was discovered as an element after it came to the attention of the western world in the red crystalline mineral crocoite (lead(II) chromate), discovered in 1761 and initially used as a pigment. Louis Nicolas Vauquelin first isolated chromium metal from this mineral in 1797. Since Vauquelin’s first production of metallic chromium, small amounts of native (free) chromium metal have been discovered in rare minerals, but these are not used commercially. Instead, nearly all chromium is commercially extracted from the single commercially viable ore chromite, which is iron chromium oxide (FeCr₂O₄). Chromite is also now the chief source of chromium for chromium pigments.

Chromium metal and ferrochromium alloy are commercially produced from chromite by silicothermic or aluminothermic reactions, or by roasting and leaching processes. Chromium metal has proven of high value due to its high corrosion resistance and hardness. A major development was the discovery that steel could be made highly resistant to corrosion and discoloration by adding metallic chromium to form stainless steel. This application, along with chrome plating (electroplating with chromium) currently comprise 85% of the commercial use for the element, with applications for chromium compounds forming the remainder.

Trivalent chromium (Cr(III)) ion is possibly required in trace amounts for sugar and lipid metabolism, although the issue remains in debate. In larger amounts and in different forms, chromium can be toxic and carcinogenic. The most prominent example of toxic chromium is hexavalent chromium (Cr(VI)). Abandoned chromium production sites often require environmental cleanup.

Physical

Chromium is remarkable for its magnetic properties: it is the only elemental solid which shows antiferromagnetic ordering at room temperature (and below). Above 38 °C, it transforms into a paramagnetic state.

Passivation

Chromium metal left standing in air is passivated by oxygen, forming a thin protective oxide surface layer. This layer is a spinel structure only a few atoms thick. It is very dense, and prevents the diffusion of oxygen into the underlying material. This barrier is in contrast to iron or plain carbon steels, where the oxygen migrates into the underlying material and causes rusting. The passivation can be enhanced by short contact with oxidizing acids like nitric acid. Passivated chromium is stable against acids. The opposite effect can be achieved by treatment with a strong reducing agent that destroys the protective oxide layer on the metal. Chromium metal treated in this way readily dissolves in weak acids.

Chromium, unlike metals such as iron and nickel, does not suffer from hydrogen embrittlement. However, it does suffer from nitrogen embrittlement, reacting with nitrogen from air and forming brittle nitrides at the high temperatures necessary to work the metal parts.

Applications

Metallurgy

The strengthening effect of forming stable metal carbides at the grain boundaries and the strong increase in corrosion resistance made chromium an important alloying material for steel. The high-speed tool steels contain between 3 and 5% chromium. Stainless steel, the main corrosion-proof metal alloy, is formed when chromium is added to iron in sufficient concentrations, usually above 11%. For its formation, ferrochromium is added to the molten iron. Also nickel-based alloys increase in strength due to the formation of discrete, stable metal carbide particles at the grain boundaries. For example, Inconel 718 contains 18.6% chromium. Because of the excellent high-temperature properties of these nickel superalloys, they are used in jet engines and gas turbines in lieu of common structural materials.

The relative high hardness and corrosion resistance of unalloyed chromium makes it a good surface coating, being still the most “popular” metal coating with unparalleled combined durability. A thin layer of chromium is deposited on pretreated metallic surfaces by electroplating techniques. There are two deposition methods: Thin, below 1 µm thickness, layers are deposited by chrome plating, and are used for decorative surfaces. If wear-resistant surfaces are needed then thicker chromium layers are deposited. Both methods normally use acidic chromate or dichromate solutions. To prevent the energy-consuming change in oxidation state, the use of chromium(III) sulfate is under development, but for most applications, the established process is used.

In the chromate conversion coating process, the strong oxidative properties of chromates are used to deposit a protective oxide layer on metals like aluminium, zinc and cadmium. This passivation and the self-healing properties by the chromate stored in the chromate conversion coating, which is able to migrate to local defects, are the benefits of this coating method. Because of environmental and health regulations on chromates, alternative coating method are under development.

Anodizing of aluminium is another electrochemical process, which does not lead to the deposition of chromium, but uses chromic acid as electrolyte in the solution. During anodization, an oxide layer is formed on the aluminium. The use of chromic acid, instead of the normally used sulfuric acid, leads to a slight difference of these oxide layers. The high toxicity of Cr(VI) compounds, used in the established chromium electroplating process, and the strengthening of safety and environmental regulations demand a search for substitutes for chromium or at least a change to less toxic chromium(III) compounds.

Dye and pigment

The mineral crocoite (lead chromate PbCrO₄) was used as a yellow pigment shortly after its discovery. After a synthesis method became available starting from the more abundant chromite, chrome yellow was, together with cadmium yellow, one of the most used yellow pigments. The pigment does not photodegrade, but it tends to darken due to the formation of chromium(III) oxide. It has a strong color, and was used for school buses in the US and for Postal Service (for example Deutsche Post) in Europe. The use of chrome yellow declined due to environmental and safety concerns and was replaced by organic pigments or alternatives free from lead and chromium. Other pigments based on chromium are, for example, the bright red pigment chrome red, which is a basic lead chromate (PbCrO₄·Pb(OH)₂). A very important chromate pigment, which was used widely in metal primer formulations, was zinc chromate, now replaced by zinc phosphate. A wash primer was formulated to replace the dangerous practice of pretreating aluminium aircraft bodies with a phosphoric acid solution. This used zinc tetroxychromate dispersed in a solution of polyvinyl butyral. An 8% solution of phosphoric acid in solvent was added just before application. It was found that an easily oxidized alcohol was an essential ingredient. A thin layer of about 10–15 µm was applied, which turned from yellow to dark green when it was cured. There is still a question as to the correct mechanism. Chrome green is a mixture of Prussian blue and chrome yellow, while the chrome oxide green is chromium(III) oxide.

Chromium oxides are also used as a green color in glassmaking and as a glaze in ceramics. Green chromium oxide is extremely light-fast and as such is used in cladding coatings. It is also the main ingredient in IR reflecting paints, used by the armed forces, to paint vehicles, to give them the same IR reflectance as green leaves.

Synthetic ruby and the first laser

Natural rubies are corundum (aluminum oxide) crystals that are colored red (the rarest type) due to chromium (III) ions (other colors of corundum gems are termed sapphires). A red-colored artificial ruby may also be achieved by doping chromium(III) into artificial corundum crystals, thus making chromium a requirement for making synthetic rubies. Such a synthetic ruby crystal was the basis for the first laser, produced in 1960, which relied on stimulated emission of light from the chromium atoms in such a crystal.

Wood preservative

Because of their toxicity, chromium(VI) salts are used for the preservation of wood. For example, chromated copper arsenate (CCA) is used in timber treatment to protect wood from decay fungi, wood attacking insects, including termites, and marine borers. The formulations contain chromium based on the oxide CrO₃ between 35.3% and 65.5%. In the United States, 65,300 metric tons of CCA solution have been used in 1996.

Tanning

Chromium(III) salts, especially chrome alum and chromium(III) sulfate, are used in the tanning of leather. The chromium(III) stabilizes the leather by cross linking the collagen fibers. Chromium tanned leather can contain between 4 and 5% of chromium, which is tightly bound to the proteins. Although the form of chromium used for tanning is not the toxic hexavalent variety, there remains interest in management of chromium in the tanning industry such as recovery and reuse, direct/indirect recycling, use of less chromium or “chrome-less” tanning are practiced to better manage chromium in tanning.

Refractory material

The high heat resistivity and high melting point makes chromite and chromium(III) oxide a material for high temperature refractory applications, like blast furnaces, cement kilns, molds for the firing of bricks and as foundry sands for the casting of metals. In these applications, the refractory materials are made from mixtures of chromite and magnesite. The use is declining because of the environmental regulations due to the possibility of the formation of chromium(VI).

Catalysts

Several chromium compounds are used as catalysts for processing hydrocarbons. For example the Phillips catalysts for the production of polyethylene are mixtures of chromium and silicon dioxide or mixtures of chromium and titanium and aluminium oxide. Fe-Cr mixed oxides are employed as high-temperature catalysts for the water gas shift reaction. Copper chromite is a useful hydrogenation catalyst.

Other use

Chromium(IV) oxide (CrO₂) is a magnetic compound. Its ideal shape anisotropy, which imparts high coercivity and remnant magnetization, made it a compound superior to the γ-Fe2O3. Chromium(IV) oxide is used to manufacture magnetic tape used in high-performance audio tape and standard audio cassettes. Chromates can prevent corrosion of steel under wet conditions, and therefore chromates are added to drilling muds.

Chromic acid is a powerful oxidizing agent and is a useful compound for cleaning laboratory glassware of any trace of organic compounds. It is prepared in situ by dissolving potassium dichromate in concentrated sulfuric acid, which is then used to wash the apparatus. Sodium dichromate is sometimes used because of its higher solubility (50 g/L versus 200 g/L respectively). The use of dichromate cleaning solutions is now phased out due to the high toxicity and environmental concerns. Modern cleaning solutions are highly effective and chromium free. Potassium dichromate is a chemical reagent,used as a titrating agent. It is also used as a mordant (i.e., a fixing agent) for dyes in fabric.

Biological role

Trivalent chromium (Cr(III) or Cr³⁺) occurs in trace amounts in foods and waters, and appears to be benign. In contrast, hexavalent chromium (Cr(VI) or Cr⁶+) is very toxic and mutagenic when inhaled. Cr(VI) has not been established as a carcinogen when in solution, although it may cause allergic contact dermatitis (ACD).

Chromium deficiency, involving a lack of Cr(III) in the body, or perhaps some complex of it, such as glucose tolerance factor is controversial, or is at least extremely rare. Chromium has no verified biological role and has been classified by some as not essential for mammals. However, other reviews have regarded it as an essential trace element in humans.

Watch glass holds two grams of pure chromium (III) picolinate powder manufactured by NUTRITION 21 as a dietary supplement. It is a reddish crystalline powder, 12.5% Cr(III) by weight, very sparingly soluble in water.

Chromium deficiency has been attributed to only three people on long-term parenteral nutrition, which is when a patient is fed a liquid diet through intravenous drips for long periods of time.

Although no biological role for chromium has ever been demonstrated, dietary supplements for chromium include chromium(III) picolinate, chromium(III) polynicotinate, and related materials. The benefit of those supplements is questioned by some studies. The use of chromium-containing dietary supplements is controversial, owing to the absence of any verified biological role, the expense of these supplements, and the complex effects of their use. The popular dietary supplement chromium picolinate complex generates chromosome damage in hamster cells (due to the picolinate ligand). In the United States the dietary guidelines for daily chromium uptake were lowered in 2001 from 50–200 µg for an adult to 35 µg (adult male) and to 25 µg (adult female).

No comprehensive, reliable database of chromium content of food currently exists. Data reported prior to 1980 is unreliable due to analytical error. Chromium content of food varies widely due to differences in soil mineral content, growing season, plant cultivar, and contamination during processing. In addition, large amounts of chromium (and nickel) leech into food cooked in stainless steel.

Precautions

Water insoluble chromium(III) compounds and chromium metal are not considered a health hazard, while the toxicity and carcinogenic properties of chromium(VI) have been known for a long time. Because of the specific transport mechanisms, only limited amounts of chromium(III) enter the cells. Several in vitro studies indicated that high concentrations of chromium(III) in the cell can lead to DNA damage. Acute oral toxicity ranges between 1.5 and 3.3 mg/kg. The proposed beneficial effects of chromium(III) and the use as dietary supplements yielded some controversial results, but recent reviews suggest that moderate uptake of chromium(III) through dietary supplements poses no risk.

Cr(VI). The acute oral toxicity for chromium(VI) ranges between 50 and 150 µg/kg. In the body, chromium(VI) is reduced by several mechanisms to chromium(III) already in the blood before it enters the cells. The chromium(III) is excreted from the body, whereas the chromate ion is transferred into the cell by a transport mechanism, by which also sulfate and phosphate ions enter the cell. The acute toxicity of chromium(VI) is due to its strong oxidational properties. After it reaches the blood stream, it damages the kidneys, the liver and blood cells through oxidation reactions. Hemolysis, renal and liver failure are the results of these damages. Aggressive dialysis can improve the situation.

The carcinogenity of chromate dust is known for a long time, and in 1890 the first publication described the elevated cancer risk of workers in a chromate dye company. Three mechanisms have been proposed to describe the genotoxicity of chromium(VI). The first mechanism includes highly reactive hydroxyl radicals and other reactive radicals which are by products of the reduction of chromium(VI) to chromium(III). The second process includes the direct binding of chromium(V), produced by reduction in the cell, and chromium(IV) compounds to the DNA. The last mechanism attributed the genotoxicity to the binding to the DNA of the end product of the chromium(III) reduction.

Chromium salts (chromates) are also the cause of allergic reactions in some people. Chromates are often used to manufacture, amongst other things, leather products, paints, cement, mortar and anti-corrosives. Contact with products containing chromates can lead to allergic contact dermatitis and irritant dermatitis, resulting in ulceration of the skin, sometimes referred to as “chrome ulcers”. This condition is often found in workers that have been exposed to strong chromate solutions in electroplating, tanning and chrome-producing manufacturers.

Environmental issues

As chromium compounds were used in dyes and paints and the tanning of leather, these compounds are often found in soil and groundwater at abandoned industrial sites, now needing environmental cleanup and remediation per the treatment of brownfield land. Primer paint containing hexavalent chromium is still widely used for aerospace and automobile refinishing applications.

In 2010, the Environmental Working Group studied the drinking water in 35 American cities. The study was the first nationwide analysis measuring the presence of the chemical in U.S. water systems. The study found measurable hexavalent chromium in the tap water of 31 of the cities sampled, with Norman, Oklahoma, at the top of list; 25 cities had levels that exceeded California’s proposed limit. Note: Concentrations of Cr(VI) in US municipal drinking water supplies reported by EWG are within likely, natural background levels for the areas tested and not necessarily indicative of industrial pollution (CalEPA Fact Sheet), as asserted by EWG. This factor was not taken into consideration in their report.

Occurrence

Chromium is present at a concentration of 122 PPM of the earth, comparable to vanadium and chlorine (126 PPM). The principle ore of chromium is chromite, FeCr₂O₄, produced mainly in the old USSR, southern Africa where 96% of known reserves are found, and the Philippines. The main use of chromium is in the production of non-ferous alloys, pure chromium having a low ductility at normal temperatures. It also finds many applications in electroplating which is both a decorative and protective process. About 9.5 million tonnes are produced annually.

MANGANESE – Mn

Manganese is a chemical element, designated by the symbol Mn. It has the atomic number 25. It is found as a free element iature (often in combination with iron), and in many minerals. Manganese is a metal with important industrial metal alloy uses, particularly in stainless steels.

Historically, manganese is named for various black minerals (such as pyrolusite) from the same region of Magnesia in Greece which gave names to similar-sounding magnesium, Mg, and magnetite, an ore of the element iron, Fe. By the mid-18th century, Swedish chemist Carl Wilhelm Scheele had used pyrolusite to produce chlorine. Scheele and others were aware that pyrolusite (now known to be manganese dioxide) contained a new element, but they were not able to isolate it. Johan Gottlieb Gahn was the first to isolate an impure sample of manganese metal in 1774, by reducing the dioxide with carbon.

Manganese phosphating is used as a treatment for rust and corrosion prevention on steel. Depending on their oxidation state, manganese ions have various colors and are used industrially as pigments. The permanganates of alkali and alkaline earth metals are powerful oxidizers. Manganese dioxide is used as the cathode (electron acceptor) material in zinc-carbon and alkaline batteries.

In biology, manganese(II) ions function as cofactors for a large variety of enzymes with many functions. Manganese enzymes are particularly essential in detoxification of superoxide free radicals in organisms that must deal with elemental oxygen. Manganese also functions in the oxygen-evolving complex of photosynthetic plants. The element is a required trace mineral for all known living organisms. In larger amounts, and apparently with far greater activity by inhalation, it can cause a poisoning syndrome in mammals, with neurological damage which is sometimes irreversible.

Manganese is the 12th most abundant element, and the third most abundant transition element at 0.106% (1060 PPM) of the earth’s crust. Over 300 minerals contain manganese, of which 12 are important commercially. More than 10¹² tonnes of “manganese nodules” litter the ocean floor, with 10⁷ tonnes deposited annually. This is due to weathering of rocks which produces colloidal particles of Mo, Fe, and other metal oxides, which agglomerate into lumps in the sea. At 15-30% Mn, it is the presence of other metals which makes these nodules an attractive source of Mn. Of its many uses, one is dominant – 95% of all the manganese ores are used in the production of steel, mainly in the form of ferromanganese. Ores equivalent to 10 million tonnes of MN are produced annually. In the steel furnace, the Mn forms MnS which joins the slag, preventing the formation of FeS which induces brittleness. It also combines with oxygen to prevent the formation of bubbles in the solid metal. Finally it increases the hardness of steel, and is used where resistance to wear and mechanical shock is important e.g. excavators, dredgers, rail crossings etc. It is also used in the production of non ferrous alloys. “Manganin” (84% Cu, 12% Mn, 4% Ni) is used in electrical instruments because the temperature coefficient of its resistivity is almost zero. Many Mn compounds find uses as well. You may have heard of potassium permanganate, KMnO₄ for example.(Interestingly, although KMnO₄ is very wellknown, the Mn⁷⁺ oxidation state it displays is actually very rare indeed.) MnO₂ finds extensive use in dry batteries, and is also used in the brick and glass industry, and in the production of MnFe₂O₄to make magnets of televisions.

Applications

Manganese has no satisfactory substitute in its major applications, which are related to metallurgical alloy use. In minor applications, (e.g., manganese phosphating), zinc and sometimes vanadium are viable substitutes.

Steel

Manganese is essential to iron and steel production by virtue of its sulfur-fixing, deoxidizing, and alloying properties. Steelmaking including its ironmaking component, has accounted for most manganese demand, presently in the range of 85% to 90% of the total demand. Among a variety of other uses, manganese is a key component of low-cost stainless steel formulations.

Small amounts of manganese improve the workability of steel at high temperatures, because it forms a high-melting sulfide and therefore prevents the formation of a liquid iron sulfide at the grain boundaries. If the manganese content reaches 4%, the embrittlement of the steel becomes a dominant feature. The embrittlement decreases at higher manganese concentrations and reaches an acceptable level at 8%. Steel containing 8 to 15% of manganese can have a high tensile strength of up to 863 MPa. Steel with 12% manganese was used for British steel helmets. This steel composition was discovered in 1882 by Robert Hadfield and is still known as Hadfield steel.

Aluminium alloys

The second large application for manganese is as alloying agent for aluminium. Aluminium with a manganese content of roughly 1.5% has an increased resistance against corrosion due to the formation of grains absorbing impurities which would lead to galvanic corrosion. The corrosion-resistant aluminium alloys 3004 and 3104 with a manganese content of 0.8 to 1.5% are the alloys used for most of the beverage cans. Before year 2000, more than 1.6 million tonnes have been used of those alloys; with a content of 1% manganese, this amount would need 16,000 tonnes of manganese.

Other uses

Methylcyclopentadienyl manganese tricarbonyl is used as an additive in unleaded gasoline to boost octane rating and reduce engine knocking. The manganese in this unusual organometallic compound is in the +1 oxidation state.

Manganese(IV) oxide (manganese dioxide, MnO2) is used as a reagent in organic chemistry for the oxidation of benzylic alcohols (i.e. adjacent to an aromatic ring). Manganese dioxide has been used since antiquity to oxidatively neutralize the greenish tinge in glass caused by trace amounts of iron contamination. MnO₂ is also used in the manufacture of oxygen and chlorine, and in drying black paints. In some preparations, it is a brown pigment that can be used to make paint and is a constituent of natural umber.

Manganese(IV) oxide was used in the original type of dry cell battery as an electron acceptor from zinc, and is the blackish material found when opening carbon–zinc type flashlight cells. The manganese dioxide is reduced to the manganese oxide-hydroxide MnO(OH) during discharging, preventing the formation of hydrogen at the anode of the battery.

MnO₂ + H₂O + -e → MnO(OH) + OH−

The same material also functions iewer alkaline batteries (usually battery cells), which use the same basic reaction, but a different electrolyte mixture. In 2002, more than 230,000 tons of manganese dioxide was used for this purpose.

The metal is very occasionally used in coins; until 2000, the only United States coin to use manganese was the “wartime” nickel from 1942 to 1945. An alloy of 75% copper and 25% nickel was traditionally used for the production of nickel coins. However, because of shortage of nickel metal during the war, it was substituted by more available silver and manganese, thus resulting in an alloy of 56% copper, 35% silver and 9% manganese. Since 2000, dollar coins, for example the Sacagawea dollar and the Presidential $1 coins, are made from a brass containing 7% of manganese with a pure copper core. In both cases of nickel and dollar, the use of manganese in the coin was to duplicate the electromagnetic properties of a previous identically sized and valued coin, for vending purposes. In the case of the later U.S. dollar coins, the manganese alloy was an attempt to duplicate properties of the copper/nickel alloy used in the previous Susan B. Anthony dollar.

Manganese compounds have been used as pigments and for the coloring of ceramics and glass. The brown color of ceramic is sometimes based on manganese compounds. In the glass industry, manganese compounds are used for two effects. Manganese(III) reacts with iron(II) to induce a strong green color in glass by forming less-colored iron(III) and slightly pink manganese(II), compensating for the residual color of the iron(III). Larger amounts of manganese are used to produce pink colored glass.

Biological role

Manganese is an essential trace nutrient in all known forms of life. The classes of enzymes that have manganese cofactors are very broad, and include oxidoreductases, transferases, hydrolases, lyases, isomerases, ligases, lectins, and integrins. The reverse transcriptases of many retroviruses (though not lentiviruses such as HIV) contain manganese. The best-known manganese-containing polypeptides may be arginase, the diphtheria toxin, and Mn-containing superoxide dismutase (Mn-SOD).

Mn-SOD is the type of SOD present in eukaryotic mitochondria, and also in most bacteria (this fact is in keeping with the bacterial-origin theory of mitochondria). The Mn-SOD enzyme is probably one of the most ancient, for nearly all organisms living in the presence of oxygen use it to deal with the toxic effects of superoxide, formed from the 1-electron reduction of dioxygen. Exceptions include a few kinds of bacteria, such as Lactobacillus plantarum and related lactobacilli, which use a different nonenzymatic mechanism, involving manganese (Mn²⁺) ions complexed with polyphosphate directly for this task, indicating how this function possibly evolved in aerobic life.

The human body contains about 12 mg of manganese, which is stored mainly in the bones; in the tissue, it is mostly concentrated in the liver and kidneys. In the human brain, the manganese is bound to manganese metalloproteins, most notably glutamine synthetase in astrocytes.

Manganese is also important in photosynthetic oxygen evolution in chloroplasts in plants. The oxygen-evolving complex (OEC) is a part of photosystem II contained in the thylakoid membranes of chloroplasts; it is responsible for the terminal photooxidation of water during the light reactions of photosynthesis, and has a metalloenzyme core containing four atoms of manganese. For this reason, most broad-spectrum plant fertilizers contain manganese.

Precautions

Manganese compounds are less toxic than those of other widespread metals, such as nickel and copper. However, exposure to manganese dusts and fumes should not exceed the ceiling value of 5 mg/m3 even for short periods because of its toxicity level. Manganese poisoning has been linked to impaired motor skills and cognitive disorders.

The permanganate exhibits a higher toxicity than the manganese(II) compounds. The fatal dose is about 10 g, and several fatal intoxications have occurred. The strong oxidative effect leads to necrosis of the mucous membrane. For example, the esophagus is affected if the permanganate is swallowed. Only a limited amount is absorbed by the intestines, but this small amount shows severe effects on the kidneys and on the liver.

In 2005, a study suggested a possible link between manganese inhalation and central nervous system toxicity in rats.

Manganese exposure in United States is regulated by Occupational Safety and Health Administration.

Generally, exposure to ambient Mn air concentrations in excess of 5 μg Mn/m3 can lead to Mn-induced symptoms. Increased ferroportin protein expression in human embryonic kidney (HEK293) cells is associated with decreased intracellular Mn concentration and attenuated cytotoxicity, characterized by the reversal of Mn-reduced glutamate uptake and diminished lactate dehydrogenase leakage.

Environmental health concerns

Manganese in drinking water

Waterborne manganese has a greater bioavailability than dietary manganese. According to results from a 2010 study, higher levels of exposure to manganese in drinking water are associated with increased intellectual impairment and reduced intelligence quotients in school-age children. It is hypothesized that long-term exposure to the naturally occurring manganese in shower water puts up to 8.7 million Americans at risk.

Manganese in gasoline

Methylcyclopentadienyl manganese tricarbonyl (MMT) is a gasoline additive used to replace lead compounds for unleaded gasolines, to improve the octane number in low octane number petrol distillates. It functions as an antiknock agent by the action of the carbonyl groups. Fuels containing manganese tend to form manganese carbides, which damage exhaust valves. The need to use lead or manganese compounds is merely historic, as the availability of reformation processes which create high-octane rating fuels increased. The use of such fuels directly or in mixture with non-reformed distillates is universal in developed countries (EU, Japan, etc.). In USA the imperative to provide the lowest possible price per volume on motor fuels (low fuel taxation rate) and lax legislation of fuel content (before 2000) caused refineries to use MMT. Compared to 1953, levels of manganese in air have dropped. Many racing competitions specifically ban manganese compounds in racing fuel (cart, minibike). MMT contains 24.4–25.2% manganese. There is strong correlation between elevated atmospheric manganese concentrations and automobile traffic density.

Role ieurological disorders

Manganism

Manganese overexposure is most frequently associated with manganism, a rare neurological disorder associated with excessive manganese ingestion or inhalation. Historically, persons employed in the production or processing of manganese alloys have been at risk for developing manganism; however, current health and safety regulations protect workers in developed nations. The disorder was first described in 1837 by British academic John Couper, who studied two patients who were manganese grinders.

Manganism is a biphasic disorder. In its early stages, an intoxicated person may experience depression, mood swings, compulsive behaviors, and psychosis. Early neurological symptoms give way to late-stage manganism, which resembles Parkinson’s disease. Symptoms include weakness, monotone and slowed speech, an expressionless face, tremor, forward-leaning gait, inability to walk backwards without falling, rigidity, and general problems with dexterity, gait and balance. Unlike Parkinson’s disease, manganism is not associated with loss of smell and patients are typically unresponsive to treatment with L-DOPA. Symptoms of late-stage manganism become more severe over time even if the source of exposure is removed and brain manganese levels return to normal.

Childhood developmental disorders

Several recent studies attempt to examine the effects of chronic low-dose manganese overexposure on development in children. The earliest study of this kind was conducted in the Chinese province of Shanxi. Drinking water there had been contaminated through improper sewage irrigation and contained 240–350 µg Mn/L. Although WMn concentrations at or below 300 µg Mn/L are considered safe by the US EPA and 400 µg Mn/L are considered safe by the World Health Organization, the 92 children sampled (between 11 and 13 years of age) from this province displayed lower performance on tests of manual dexterity and rapidity, short-term memory, and visual identification when compared to children from an uncontaminated area. More recently, a study of 10-year-old children in Bangladesh showed a relationship between WMn concentration in well water and diminished IQ scores. A third study conducted in Quebec examined school children between the ages of 6 and 15 living in homes that received water from a well containing 610 µg Mn/L; controls lived in homes that received water from a 160 µg Mn/L well. Children in the experimental group showed increased hyperactive and oppositional behaviors.

Neurodegenerative disease

Chronic low-dose manganese intoxication is strongly implicated in a number of neurodegenerative disorders, including Alzheimer’s disease, Parkinson’s disease, and amyotrophic lateral sclerosis. It may also play a role in the development of multiple sclerosis, restless leg syndrome, and Huntington’s disease. A protein called DMT1 is the major transporter involved in manganese absorption from the intestine, and may be the major transporter of manganese across the blood–brain barrier. DMT1 also transports inhaled manganese across the nasal epithelium. The putative mechanism of action is that manganese overexposure and/or dysregulation lead to oxidative stress, mitochondrial dysfunction, glutamate-mediated excitoxicity, and aggregation of proteins.

IRON – Fe

Iron is the 4th most abundant element at 6.2% (62000 PPM) after oxygen, silicon, and aluminium, making it the commonest transition metal, and it is one of the most important elements. It is produced on a vast scale in the blast furnace process. Iron finds use in an enormous variety of steels, annual world production being about 700 million tonnes i.e. it is very important economically, and of course its main use is in building things. Iron is the catalyst used in the Haber process to manufacture NH₃. (More moles of NH₃ are manufactured than any other compound and only H₂SO₄ is made on a bigger tonnes scale, and both compounds have many uses.)

COBALT – Co

Cobalt is present at 29 PPM i.e. 0.0029% of the earth’s crust, is only the thirtieth most abundant element and apart from Sc at 25 PPM, is the rarest of the first row transition elements. Only a few important ores exist although over 200 cobalt containing minerals are known. About 33,000 tonnes are produced annually. About 30% of this is used to produce chemicals for the ceramic and paint industries. Many blue paints contain cobalt compounds. Cobalt compounds are also used as catalysts in the OXO process for the formation of aldehydes, and hydrogenation and dehydrogenation processes are equally important. Another 30% of the cobalt is used in the production of high-temperature alloys in the construction of gas turbines. Many magnetic alloys containing Co are also important e.g. Alnico steel.

NICKEL – Ni

Nickel is the seventh most abundant transition metal and 22nd most abundant element at 99 PPM Over a quarter of the world production of Ni comes from the Sudbury Basin in Canada, a giant meteor crater measuring 17 by 37 miles. Total production of the metal is around 750,000 tonnes. Most is used in the production of alloys, both ferrous and non ferrous. Stainless steel contains up to 8% Ni and it is also used in Alnico steel for strong magnets. Small amounts are used as catalysts for hydrogenation of unsaturated vegetable oils. Nickel is well known as a useful catalyst for hydrogenation and other reactions in the lab and industry, often as the very finely divided form called Raney Nickel.

COPPER – Cu

Copper is present at 68 PPM in the earth’s crust and has several important ores, as well as being found native. About 8 million tonnes are produced annually. An important metal, copper’s major use is as a conductor of electricity, but it is also used in the formation of coins (money), as well as brass (Cu-Zn), bronze (Cu with Sn) and special alloys like Monel (Ni-Cu)

ZINC – Zn

Zinc is more abundant than copper at 76 PPM, and about 6 million tonnes are produced each year. About 35-40% of this is used as an anti-corrosion coating – e.g. galvanized steel, which is simply steel that has been zinc plated to prevent it rusting. Apart from brasses, several special alloys of zinc are used for diecasting, including the majority of pressure diecastings. Large quantities are used in the manufacture of batteries, and as a roof cladding.

 

 

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

9. http://en.wikipedia.org/wiki

The following website shows the reaction of d-elements. It’s cool stuff! Check it out!

www.youtube.com/watch?v=_dIafW3yrEc

www.youtube.com/watch?v=8cyghkL-q0g

www.youtube.com/watch?v=6b_-oydvqEA

www.youtube.com/watch?v=fYGBMWU9btM

 

Prepared by PhD Falfushynska H.