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Option - The Chemistry of Art Chapter 1 - Pigments Throughout History Sources of Pigments in Early History

Colour Pigments Black Charcoal burnt wood, animal bones, soot White Gypsum CaSO4.2H2O

Chalk - CaCO3 Kaolin - Al2Si2O5(OH)4

Red Ochre crushed rocks Cinnabar HgS

Yellow/Brown Orpiment As2S3 Limonite FeO(OH)H2O

Green Malachite Cu2(CO3)(OH)2 Grey Galena PbS

Cerussite PbCO3 In early history, the pigments used were the minerals readily available in coloured earth and soft rocks. Insolubility of Pigments Pigments need to be insoluble in most substances for several reasons:

1. To ensure that the pigment particles remain on the painting after the medium has dried 2. To ensure that the pigments are not easily removed from the surface of the painting when

exposed to rain 3. To ensure pigments used in cosmetics do not get washed away by perspiration

Prehistoric Use of Pigments Self Decoration Aborigines:

Body art used red and yellow ochres, white clay and charcoal as pigments Pigments were applied with fingers or brushed to produce symbols representing history and

myths Mainly used for ceremonial purposes

Egyptians:

Accentuated the eye as the main facial feature (symbol of the sun god Ra) Cosmetics used include: kohl as eyeliner, green malachite as eye shadow, white lead to

whiten the skin, red ochres on lips and cheeks and dyes as nail polish Raw pigments were prepared by grinding on a stone slab and applied in powder form Some cosmetics contained metallic poisons e.g. lead, mercury and antimony

Ancient Greeks:

White lead was used to whiten faces and red lead was used as blush on cheeks Prehistoric Use of Pigments Preparation of the Dead for Burial The Egyptians had a reverence for the passage of the dead into the afterlife and therefore used pigments for the dead in much the same way as they did for the living.

The main colours used included white, black, red, yellow and ultramarine blue and due to the lack of sunlight in the tombs, the colours remained vivid for long periods of time

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Egyptian tombs (sarcophagi) were pained inside and out with representations of the dead and the life they lived

Mummified bodies were also decorated with geometric designers to ensure the continued survival of the owner

Cosmetic boxes of pigments were often buried with the dead in tombs to be used in the afterlife e.g. powdered gypsum, orpiment and haematite

Prehistoric Use of Pigments Cave Drawings The Aboriginal cave drawings are considered to be amongst the earliest known paintings.

Few pigments were available, with the main ones being: chalk, ash, charcoal, limonite, kaolin and red, yellow and brown ochres

Binders used included egg yolk, tree resin, blood, urine and saliva Traditional techniques were used to apply paints e.g. use of sticks and barks as well as

body parts such as fingers and spraying with the mouth Cave paintings were used to tell stories and pass on information e.g. tales of the

Dreamtime Health Risks of Heavy Metal Pigments Ancient cultures including the Egyptians and the Romans often used heavy metals in the preparation of cosmetic products.

Antimony - Sb2S3 was used as a kohl, mascara or eye shadow. It is known to be carcinogenic and can cause nausea and vomiting in the short term

Copper - used in eye make-up in the form of malachite or as a kohl in black copper oxide form. It can cause anaemia, liver and kidney damage and intestinal irritation

Lead - Galena (PbS) was crushed up and used in black eye makeup. White lead (PbCO3) was also used as a face powder. Use may result in damage to the nervous system, mental retardation and even death

Mercury - often used in cinnabar (HgS) for lipsticks due to its vibrant red colour. Mercury poisoning can result in numbness, tunnel vision and brain damage

Structure of Paintings Wood and canvas (the support) are normally unsuitable to paint on directly as they are too rough and absorbent and will therefore be prepared through layers of ground or by priming. Gypsum or chalk would be applied as a warm liquid and then later scrapped off as a white layer. On the ground, paint layers were added, which consisted of pigments mixed with a binding medium or oil. Paint layers were then protected by a cover of varnish.

Medieval Painting An example of a medieval painting is Madonna and Child with Saints Jerome, John the Baptist, Bernardino and Bartholomew. The painter was Sano Di Pietro and it is estimated that the art work was completed in Sienna, Italy 550 years ago.

In this painting, a layer of linen was glued to the wood. Next, a layer of gesso grosso (CaSO4) was mixed in with animal glue, followed by finer layers of gesso sottile. Finer layers are highly polished to a smooth surface and fine black powder is added to fill in any holes. Charcoal was used in the underdrawing. The gold leaf was adhered to the surface of the pain through oil, varnish or garlic juice. Mordant gilding was used in order to apply it over the paint. Pigments used included vermilion (red), ultramarine (blue), gypsum (white), Naples yellow and orpiment (yellow) and malachite (green). Obtaining Colour For paintings, colours are obtained through spreading several layers of paint and superimposing them, allowing the creation of complex colour effects. Pigments themselves are rarely mixed. The same pigments were also used in the production of coloured glass. The oxides were finely powdered and added to the class mixture before melting e.g. cobalt oxide produces blue and manganese oxide produces a purple, both in accordance to temperature and pressure. Flashing involves coating clear glass with a thin layer of coloured glass to produce a lighter colour. Staining involves painting glass with silver nitrate before firing it in an oven. According to the number of times the glass is stained, a range of tones can be achieved. Structure of Paints Paint is formed through the colouring matter (the pigment) and a liquid (medium) which carries the pigment, allowing it to be spread over a surface.

A pigment is a substance in particulate form which is substantially insoluble in a medium, but which can be dispersed in this medium to modify its colour

The binder is the material used in paint that causes pigment particles to adhere to one another to the support. This provides a protective film, improving hardness and resistance to water

Examples of binders include:

Egg tempera - formed when egg yolk is mixed with a pigment to form a long lasting paint, used since the 14th century

Linseed oil - mixed with pigments in varying proportions to form oil paints, goes brownish yellow with age and is generally acidic

Acrylics - these are used in the majority of synthetic paints used today as they are durable, flexible and suitable for use on paper, panels and canvas

The Egyptians were advanced in using dyes, which are different to pigments in that the colouring matter is dissolved in solution. This provided new pigments for painting e.g. vegetable dyes were dissolved in water and fixed onto an insoluble white powder, which could be dried and extracted. Lake pigments are manufactured by precipitating a dye with an inert binder. The binder used must be inert and insoluble and must be white or very neutral. Madder lake is a natural pink-red dyestuff formed form the root of the madder plant rubia tinctorium. Lakes are fugitive in that they will lighten in a relatively short time when exposed to light because the dyes are unstable. Increasing Range of Pigments The very first Aboriginal cave paintings only used red, yellow and black pigments as these were easily obtained in the forms of ochres and charcoal. As time progresses, new minerals are available due to improvement in extraction methods e.g. terra verte (green earth) was the first green pigment used and was only used after the iron silicate and clay mixture could be extracted.

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The Egyptians and the Romans had better extraction technologies and therefore had a wider array of pigments available. Better purification techniques mean pigments are more easily isolated from ores. More recently, there has been a shift from natural to synthetic pigments due to further advancements in technology. More and more organic pigments (carbon based) are being used. Colours and the Periodic Table

Metals are present in each pigment and most of the metals involved are transition metals. The metals present in the pigment often have the same colour or colour range.

Chapter 2 - Structure of the Atom and Colours Flame Colours A flame test reveals that many metallic elements emit light of a distinctive colour when their compounds are heated in a flame or when their vapours are exposed to electrical discharge.

Name of Element Cation Flame Colour Sodium Na+ Yellow

Potassium K+ Lilac Calcium Ca2+ Orange-red Barium Ba2+ Apple-green

Strontium Sr2+ Red Copper Cu2+ Green-blue

Excitation of Electrons According to the Bohr Model, electrons can move between energy levels by absorbing or emitting photons. A photon is an energy packet of electromagnetic radiation and its energy will equal the difference in energy between two orbits or energy levels. Electrons are promoted or excited into higher energy levels when it is exposed to energy (e.g. heat). As the electron drops down to lower energy levels, it emits photons characteristic of the element and thus produces the colour we see. The energy of an atom is not continuous but quantitised i.e. it exists only in fixed amounts. The photon frequencies absorbed or emitted by an atom are fixed by the differences between energy levels of the orbit of electrons. Any photon that has a frequency that lies between two orbits is unable to excite the electron. Line Spectra Atoms give off visible light and other forms of electromagnetic radiation when heated or otherwise energetically excited. However, this is not a continuous distribution of wavelengths (i.e. the whole spectrum of visible light) but only consists of a few. Spectral lines are the discrete lines which represent the frequencies of light emitted when a gaseous sample of an element is dissociated and the atoms are excited. The absorption spectrum is produced when atoms of a cool gas absorb photons of certain wavelengths from white light and become excited from lower to higher energy levels. When photons from the light pass through the gas, some of them can interact with the atoms and bump an electron to a higher energy level (provided that they have just the right frequency). Photons at those frequencies are thus absorbed by the gas and the rest pass through the sample. It therefore appears as black lines against a bright background. The emission spectrum is formed when atoms that have been excited to higher energy levels emit photons characteristic of the element as they return to lower energy levels. In the visible part of the electromagnetic spectrum, it is seen as coloured lines on a black background. Emission and absorption spectra for the same substance have lines at identical positions.

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The reflectance spectrum is formed when atoms and molecules absorb and reflect energy at wavelengths related to their atomic structures. When a substance absorbs certain wavelengths of light in the visible region, the colour of the substance is determined by the wavelengths of visible light that remain. The substance exhibits the colour complementary to those absorbed. Bohr Model of the Hydrogen Atom In 1913, Bohr attempted to explain the hydrogen spectrum and proposed that in a hydrogen atom, the electron moves around the nucleus in a circular orbit without radiating energy. He also put forward the idea that only by emitting or absorbing a photon can an electron move from one electron shell to another. Part of his proposal was the belief that the hydrogen atom had only certain allowable energy levels or stationary states. Each state was associated with a fixed circular orbit (electron shell). In his work, Bohr used Rutherford's model of the atom, which involved electrons orbiting the nucleus, as well as Max Planck's theory that energy is quantised i.e. it exists as photons. The Principal Quantum Number Bohr calculated a set of allowed energies, with each value corresponding to a circular path of different radii. Each energy level corresponds to a principal quantum number (n), they will have the values 1, 2, 3, 4 etc. The term shell can also be used to describe the energy levels. The energy level closest to the nucleus is denoted as n = 1 and so on. The lower the quantum number, the smaller is the radius of the orbit and the lower is the energy of the atom. Energy Levels and Colours Every line in the hydrogen spectrum represents a transition from one energy level to another. When an electron jumps from one energy level to a lower energy level, a photon of radiation is emitted. The energy of this photon is exactly equal to the difference in energy between the two levels of this atom. For some of these emissions, the energy released is in the visible spectrum region and so can be seen on the visible line spectrum. When the electron is in a higher level, the atom is said to be in an excited state. When the electron is in its lowest energy state, the atom is said to be in its ground state. Elements other than hydrogen also show atomic line emission spectra when energetically excited electrons fall from higher energy levels to lower ones. Merits and Limitations of the Bohr Model Merits

1) Bohr's model attempted to explain why electrons did not spiral into the nucleus (principle of classical mechanics that explained how particles on a curved path emit EMR and so would eventually lose energy)

2) Attempted to explain how emission and absorption spectra for different elements had a different and complex energy structure

3) Was able to predict the line spectrum of hydrogen with reasonable precision Limitations

1) The model does not make it possible to calculate the wavelengths of the spectral lines of all the other atoms

2) It only works well for atoms with one electron in their outer shell but not for others 3) Does not explain why some spectral line are more intense than others 4) Does not explain hyperfine lines (e.g. Na doublet) 5) Could not account for the Zeeman effect where the emission spectrum shows a splitting of

spectral lines when a gas is excited in a magnetic field 6) The Bohr model is a mixture of classical and quantum physics

The Sodium Doublet In the case of sodium, the 3p orbital splits into the two different energy levels of 3p3/2 and 3p1/2 due to the magnetic fields caused by the electrons' spins and their orbits (Zeeman effect). As these two energy levels have a minute difference, there are two hyperfine lines on the emission spectra known as the sodium doublet. The two lines are 0.597 nm apart, with the line at 589.0 nm having twice the intensity of the line at 589.6 nm due to the energy levels. Analysis of Pigments Infrared absorption spectroscopy identifies the presence of organic functional groups (e.g. in binders and organic pigments) and allows the molecules to be identified. It is both quantitative and qualitative but it is destructive in that you have a take a sample and create a solution. The tested material has to be soluble in a solvent. IR radiation sets up vibrations in the atoms of molecules. The wavelength at which a vibration occurs is characteristic of the elements forming the bond. These vibrations can then be measured by passing infrared light through the sample and recording at which wavelengths the absorptions occur. Ultraviolet absorption spectroscopy can be used to determine purity or concentration of the solution of a substance that absorbs UV light as it can cause specific fluorescence in materials depending on composition and age. The absorbing materials can then be compared to developed calibration tools to make measurements. It is a quantitative tool.

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Type of Analysis Source of Radiation Type of Detector UV-Vis Tungsten lamp or deuterium

discharge tube A photomultiplier tube which detects and amplifies light

IR Heated ceramic e.g. a SiC rod A thermocouple In double beam spectrophotometers, the radiation is split into two beams. One is passed through the sample while the other is passed through the reference, usually the solvent used to dissolve the substance being investigated. The detector measures the radiation passing through and a comparison of the two beams allows the absorption of radiation to be determined. Infrared reflectography is a non-destructive form of analysis that supplies information about the whole painting. An IR reflectance lamp is shone on the surface of a painting and the reflected light is detected. The IR radiation is able to penetrate the pigments and reflect off the white ground. It is particularly useful for detecting under-drawings in a painting where the artist has used graphite pencil, charcoal or black inks. Ultraviolet reflectography uses the fact that higher energy photons in the UV region of the spectrum are strongly absorbed in these varnishes and binding media. The absorption of UV photons causes chemical reactions, sometimes resulting in the emission of photons in the visible region i.e. fluorescence. It can provide additional information in identifying certain pigments and varnishes. Reflectance Spectra and the Effect of Infrared and Ultraviolet Light Reflectance spectra of pigments can be used to determine their presence in different paintings. A pigment can be studied by shining white light on the surface of the painting and then examining the spectrum of light that is reflected. As the reflectance spectrum is the compliment of the absorption spectrum, it can be compared to known spectra for identification.

Infra-red radiation is known to change the colour of zinc oxide from white when cold to yellow when hot

It also changes red CuO2, malachite and verdigras permanently to black copper oxide UV light has less of an effect and just causes pigments to fluoresce ZnO fluoresces a yellow colour while malachite (containing copper) fluoresces a dirty-

mauve colour Laser Microspectral Analysis Laser microspectral analysis involves concentrating a powerful pulsed laser on a small sample of the pigment to be identified. The material is then vaporised (i.e. it is destructive) and the vapour is fed through a gap between two electrodes that sparks and excites the atoms and ions, producing an emission spectrum as electrons are de-excited. The obtained emission spectra consist of lines corresponding to the elements evaporated from the sample surface. This is useful for determining trace elements in solid and liquid samples. It is important as it has high sensitivity and requires minimal preparation. It is particular useful in analysing the elemental composition of pigments used in restoring paintings. A microscopic amount of a colour of the painting is vaporised using laser energy and from the analysis, a synthetic substitute for this colour can be prepared. Another use is in the identification of the validity of an artwork (i.e. whether a painting is from the era it is thought to have been painted in).

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Chapter 3 - The Periodic Table and Electron Distribution Pauli Exclusion Principle The Pauli exclusion principle states that no two electrons in the same atom can have the same set of four quantum numbers. Each electron has its own 4 quantum numbers referring to the level of the shell (n), sublevel or sub-shell (l), orbital (ml) and spin (ms). Orbits and Orbitals An orbit refers to the energy level or shell in which an electron can be found. An orbit with quantum number n can contain up to 2n2 electrons. An orbital is different and can contain zero, one or two electrons. If two electrons are present they have opposing spins. Sub-Shells Experimental and theoretical evidence suggests that apart from the n=1 shell, each principle shell consists of a number of energy sublevels or sub-shells with slightly different energies. The number of sub-shells in any shell is equal to the quantum number of a shell e.g. the n=2 shell has two sub-shells. More specifically, they are one or more orbitals with the same set of n and e quantum numbers. Ground States and Excitation When an electron is in the orbital closest to the nucleus, it is in its ground state or lowest energy level. Electrons are only able to move to a higher energy level when they absorb a photon whose energy equals the difference in energy between the two stationary states. Hund's Rule Hund's Rule states that: if two or more orbitals of the same energy are available (same subshell), then an electron will slot into each orbital until all orbitals are half filled with electrons of the same spin before any orbital receives a second electron. This produces a configuration of low energy. If two electrons were placed in the same orbital when there were empty orbitals available, their mutual repulsion would result in a higher potential energy.

The S, P, D and F Blocks Information about energy levels and orbital shells can be summarised:

The lowest energy sub-shell within a shell is called a s-subshell and contains one orbital i.e. an s-orbital. As there is only one orbital, the maximum number of electrons the sub-shell can accommodate is two

The second lowest energy sub-shell within a shell is the p-subshell, which contains three identical orbitals called p-orbitals. As there are three orbitals, the maximum number of electrons accommodated is six

The third lowest energy subshell within a shell is called a d-subshell and contains five orbitals called d-orbitals. As there are five orbitals, the maximum number of electrons the shell can hold is ten

The fourth lowest energy subshell is a f-subshell and contains seven orbitals called f-orbitals. As there are seven orbitals in a f-subshell, the maximum number of electrons held is 14

To determine the order in which sub-shells are filled, place the electrons into orbitals starting with the lowest energy orbital first. Secondly, place a maximum of two electrons in each orbital but where more than one orbital of the same energy is available, place one electron in each orbital before pairing electrons up. The atomic number of sodium is 11, meaning there are 11 electrons. The first two electrons are located in the 1s orbital. The next two are located in the 2s orbital and the next six in the three 2p orbitals in the 2p subshell. The 11th electron is located in the 3s orbital. The electron configuration of sodium can therefore be written as 1s22s22p63s1. Electron Configuration and the Periodic Table

When elements are listed in order of increasing atomic number, similar outer shell or energy level electron configurations are observed to recur at regular intervals. In the periodic table, elements with similar outer shell electron configurations appear in the same vertical group. For example, the group IA elements all have an s1 outermost or valence shell electron configuration.

Electronegativity Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to itself. The greater an atom's electronegativity, the greater its ability to attract electrons to itself. It is expressed on a relative scale with the most electronegative element fluorine assigned a value of 4.0. Electronegativity depends on several factors:

The number of protons in the nucleus The distance from the nucleus The amount of screening by inner electrons

Electronegativity increases across periods from left to right i.e. from metals to non metals. This is because there is an increase in nuclear charge across the periods. Although an extra electron is added for each element, this does not fully shield the effect of the increased number of protons. Electronegativity decreases going down groups of the periodic table. This occurs because in moving down a group, the size of the atom is increased and the attractive force of the shielded nucleus is therefore reduced. Ionisation Energy Ionisation energy is the amount of energy required to remove the outermost electron from a mole of gaseous atoms or ions. The first ionisation energy is the energy required to remove the first electron from a neutral atom. Strength of ionisation energy will be determined by the same factors affecting electronegativity. The elements that show the highest ionisation energies are helium, neon and argon. As these are the noble gases, this indicates that the electron configurations of these elements are very stable.

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On the other hand, the alkali metals (e.g. lithium, sodium and potassium) have very low ionisation energies. This is because these elements have only one electron in their valence or outer shell. Overall, this means that ionisation energy increases moving from left to right across a period. There is a gradual decrease in ionisation energy going down a group. This occurs because the outer electrons are further from the nucleus as atomic number increases. Although the charge on the nucleus increases, the number of electrons shielding the outer electrons from the nuclear charge also increases.

Successive Ionisation Energies When successive electrons are removed from an atom, they are removed from their orbitals in the reverse order to which they were filled. The second ionisation energy (l2) is the energy required to remove the second electron and so forth. Successive ionisation energies of atoms have larger values because the electrons are being removed from progressively larger positive charges. The greater the charge on the ion from which the electron is removed, the greater the energy required to remove the electron. Another factor that can contribute to an increase in ionisation energy is the removal of an electron from a shell or energy level closer to the nucleus.

Chapter 4 - Transition Metals The Transition Elements A transition metal is one which forms one or more stable ions which have incompletely filled d-orbitals. Transition elements occupy the d-block of the periodic table. They occur in the fourth, fifth and sixth rows of the periodic table and result from the filling of the d-subshell of the third, fourth and fifth shells. Each d-subshell contains five orbitals and as a result, each transition series consists of 10 elements. Zinc has the electron 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 therefore is not a transition metal. When d-block elements form ions, the 4s electrons are lost first. Oxidation States of Transition Elements

Oxidation states refers to the number of electrons gained or lost by an atom or ion. Transition metals differ from most main-group metals in that they exhibit a variety of oxidation states. This is because they can lose electrons from both the 3d and 4s sub-shells, which have similar energies. The +2 oxidation state, which occurs commonly for nearly all transition metals, is due to the loss of two 4s electrons. Oxidation states above +2 result from the additional loss of 3d electrons. Maximum oxidation states are equal to the total number of 4s and 3d electrons in the atom. E.g. from scandium through to manganese, the maximum OS

increases from +3 to +7. Oxidation States and Colour Changes The colour of transition metal compounds are due to the absorption of photons of light by electrons. The small energy differences between the d-orbitals are similar to the energies of photons of visible light. Thus the absorption of photons of appropriate frequency can result in an electron being excited from a lower to a higher energy orbital. Transition metals in different oxidation states often exhibit different colours e.g. when yellow VO2+ is combined with zinc in dilute HCl, it is reduced to blue VO2+, then green VO3+ and finally violet V2+. The colour of transition metal compounds will often give a clue to the transition metal involved e.g. most copper(II) compounds are blue while nickel compounds are usually green. Species of Transition Metals as Oxidising Agents Transition metal elements in their highest oxidation state will combine with highly electronegative oxygen and fluorine e.g. chromium forms the chromate and dichromate ions. Complex ions such as CrO42- are strong oxidising agents because they contain many oxygen atoms. These are readily able to absorb more electrons because most oxygen atoms only have six in their outer shell.

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In general, transition metal ions in which the metal has a high oxidation state tend to be strong oxidising agents e.g. +6 in Cr2O72-. The strength of an oxidising agent depends on the ease with which the compound will accept electrons (and hence be reduced).

Chapter 5 - Coloured Complex Ions Hydrated Ions A hydrated ion is an ion in which a specific number of water molecules is associated with each formula unit and is formed when an ionic substance dissolves in water, with the ions disassociating and being surrounded by water. E.g. Cu2+ forms blue Cu(OH2)42+ in aqueous solution. Hydrated ions are examples of complex ions or coordination complexes. Complex Ions and Ligands In a complex ion, a central metal ion is surrounded by a group of anions or molecules called ligands. The number of ligand atoms bonded to the central atom is the coordination number. Examples of ligands include anions, such as Cl- and CN-, and polar molecules, including H2O and NH3. Compounds that contain complex ions are known as coordination compounds.

Due to the definition of Lewis acids and bases, ligands act as Lewis bases in complex ions. The coordinate covalent bond forming between the ligand and the metal ion is therefore an acid/base interaction. As all ligands are bases, they must have at least one unshared pair of electrons that can be used to form a coordinate covalent bond to a metal ion. The charge on the metal complex is equal to the charge on the metal ion (its oxidation state) plus the sum of the charges on the ligands. Chelated Ligands Chelated ligands are ligands that have more than one donor atom. Ligands with two or more donor atoms tend to form rings around the complex ion. Chelate comes from the Greek word meaning 'crab's claw', with these ligands surrounding the metal ion like a crab claw.

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Bidentate ligands have two lone pairs of elections, both of which can bond to the central metal ion. This includes the oxalate or ethanedioate ion (-OOC-COO-) or ethylenediamine. Quadridentate ligands have four lone pairs, all of which bond to the central metal ion. An example is haemoglobin, which consists of an iron (II) ion surrounded by a complicated molecule known as haem. Haem is a hollow ring of carbon and hydrogen atoms, at the centre of which are 4 nitrogen atoms with lone pairs. A globin molecule and a water molecule then attach onto the spaces above and below the iron ion. A hexadentate ligand has 6 lone pairs of electrons. The best example is EDTA (ethylenediaminetetraacetic acid) which is most commonly used as a negative ion (EDTA-4). The EDTA ion entirely wraps up a metal ion using all of its lone pair of elections. This allows for medical uses e.g. heavy metal poisoning.

Colours of Transition Metal Complexes Different colours are seen in metal complexes because the energies of the d-orbitals are affected by both the metal ions and its surrounding ligand groups. As there are separations in the energy gaps between d-orbitals, the frequencies of the photons absorbed, and therefore the colour, is affected. Importance of Models Models are important in explaining:

How metal ligands form Why certain geometrical shapes are preferred Why these compounds are brightly coloured and Why some ligands can chelate

Valence Bond Theory An extension of the Valence Bond Theory explained that bonding in the formation of complexes depends on the orbitals available, the tendency of the ions or groups to share a pair of elections, the number of molecules placed around a central ion and the geometry assumed by the ligands. It assumes that all coordinate bonds are entirely covalent. Example 1: Mn2+ + 4Cl- MnCl42-

The Mn2+ has the electron configuration of 1s22s22p63s23p6. One pair of electrons on each chloride ion can form a covalent bond with the 4s and each of the 4p orbitals on the manganese ion. This therefore

forms the sp3 complex, which is tetrahedral in shape.

As Cl- is a weak field ligand, there was no forced pairing. As the complex ion has unpaired electrons, it is paramagnetic i.e. it is attracted in a magnetic field. Valence bond theory postulates that the electrons from the ligands are donated into the hybridised 4s and 4p orbitals. Example 2: Fe3+ + 6CN- Fe(CN)63-

The Fe3+ ion has the electron configuration 1s22s22p63s23p6. In this instance, the extra pairs of electrons of the cyanide ions occupy the d, s and p orbitals of the metal ion and

the Fe3+ is surrounded by six cyanide ions. As CN is a strong field ligand, it creates forced pairing. Two unpaired electrons enter into orbitals already with electrons, leaving two pairs and one unpaired. This creates empty d-orbitals for incoming elections, allowing for the formation of hybridised d2sp3 type orbitals. As there is one unpaired electron a 3d orbital, the ion is very weakly paramagnetic. As only the inner d orbitals are used, it is an inner orbital complex. Example 3: Fe3+ + 6H2O Fe(H2O)63+

As water is a weak field ligand, there is no forced pairing. Instead, the 6 pairs of electrons have to enter into the next 6 available orbitals, thus leading to sp3d2

hybridisation. Again, it is paramagnetic but it forms an outer orbital complex because outer d-orbitals are involved. Crystal Field Theory The Crystal Field Theory assumes that the bonds between the ligand and the metal ion are completely ionic, with the ligand and metal ion being infinitesimally small, non-polarisable point charges. A crucial aspect of the theory uses the fact that as an approaching ligand nears a charged particles, disturbances are caused in the electron clouds of both the ion and the ligand. This affects the orientation and energies of the d-election orbitals and produces a new set of unequally spaced energy states. In a free atom, the d-electrons are said to be degenerate as they is an absence of an electric field. If a field is applied in particular directions, the degenerate levels are separated into two or more sets of levels, with the magnitude of the separation depending on the arrangement of the ligands. In the octahedral case, two orbitals are concentrated in regions of space pointing directly at the bonding electrons of the ligands while the others are concentrated in regions away from the bonding electrons. These electrons would prefer to be as far as possible from the electrons of the ligands, leading to a separation into two sets: 3 at a lower energy and 2 at a higher energy. If the transition metal ion has electrons in the d-orbitals, the electrons will enter the split d-orbitals so that the lowest energy levels fill first. They will only pair when there is an energy advantage.

The ammonia ligands in the cobalt complex produce a strong field, resulting in a large separation between the energies of the orbitals. Less energy is used in overcoming the repulsion of another electron in two of the lower orbitals than would be needed to promote two electrons to the higher orbitals. This is a low-spin complex.

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The water ligands in the cobalt complex produce a weak field and therefore the separation between the energy levels is less. The electrons remain unpaired, since the energy used to promote two electrons to the higher levels is less than would be needed to overcome the repulsion of another electron in two lower energy levels. This is therefore known as a high-spin complex. The crystal field theory can therefore explain the colours of transition metals. Electrons in the lower d levels may absorb photons and move to higher d levels (d-d transitions). E.g. the violet colour of the Ti(H2O)63+ ion occurs when the single d-electron absorbs a photon and moves to one of the higher d-levels. The light transmitted by the solution will be white minus the green photons absorbed i.e. violet.