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The Chemistry of Art Part 1: The Earliest Uses of Colour Sources of Early Pigments Early humans used pigments from natural sources such as ochre, clay and charred wood, which were mixed with a medium, including egg yolk, tree and plant gums. Colour Black White Red Brown Yellow Source Charcoal Burnt wood, animal bones, lamp soot Pyrolusite Manganese(IV) dioxide Kaolin Hydrated aluminium silicate Chalk Calcium carbonate Gypsum Calcium sulphate dihydrate Haematite Anhydrous iron(III) oxide Burnt yellow ochre Limonite Partly hydrated Fe2O3 iron(III) oxide hydroxide Goethite Hydrated iron(III) oxide Chemical Composition C MnO2 Al2O3.2SiO2.2H2O CaCO3 CaSO4.2H2O Fe2O3 FeO(OH) Fe2O3.H2O Insolubility of Pigments The pigment needed to be insoluble in most substances as when the usually colourless medium dries, the colour remains. An advantage of insoluble pigments is that they are not easily removed when exposed to rain or ground water for cave and rock paintings. In hot climates, cosmetics made from insoluble minerals would not dissolve due to perspiration. Early Uses of Pigments Early humans used pigments in their cave drawings over 17000 years ago using basic colours of red, yellow, black and white. Pigments were ground to powder and combined with a medium (e.g. egg yolk, animals fats), which allows the pigment to be transferred to the cave wall. They were applied using twig tips, fingers or brushes of furs and feathers. These were important in Aboriginal art to depict the Dreamtime, which were not only used on cave walls but also rocks and bark. Pigments were also used for self-decoration including cosmetics. Aboriginal culture had certain dances and rituals (e.g. corroboree, coming of age) where the body was painted using red and yellow ochre. Egyptians used the red pigment, cinnabar, for rouge and lipstick, yellow orpiment for eye shadow, black kohl as eyeliner and mascara and white lead to whiten the skin. In the preparation of the dead for burial, pigments were also used. The Egyptians painted the dead and the containers in which the body was entombed. In Tutankhamen’s tomb, a paint box was found containing powdered gypsum, orpiment, haematite and malachite, suggesting that these were used in the preparation of his body. Red natron was used in body scrubs, which were often buried with the dead along with other cosmetics such as cleansing creams of oil and lime. Processes to Attach Pigments to Surfaces Madonna and Child with Saints Jerome, John the Baptist, Bernardino and Bartholomew is a mid-15th century altar piece with gold leaf and tempera by Sano di Pietro (1405-1481). Paintings have similar basic components including support made of smoothed wood panels and canvas (pine or oak for this painting). As these are still rough and absorbent, ground or priming is used to prepare it for painting. It was usually gesso (gypsum and animal glue) that was applied, and on setting, it was scraped and rubbed smooth. (Coarse gesso grosso, then fine gesso sottile, then fine black powder (e.g soot)) On the ground, the artist may outline their design in an underdrawing. The paint layers consisted of pigments mixed with a binding medium or drying oil. In this case, tempera was used as the medium that binds pigment particles together to achieve better adhesion to the surface. Varnish may be applied to protect the paint layers and to give the colours clarity and depth of saturation. This painting also uses gold gilding on the frame, which is adhered to the surface on top of red paint by oil, varnish or garlic juice (Mordant). The panel is gilded by gold leaf, which is then painted over and the design scratched into the paint to show gold (Sgraffito) Pigment Vermillion Ultramarine (Lazurite) White lead (Cerrusite) Orpiment Charcoal Azurite Malachite Colour Red Blue White Yellow Black Blue Green Chemical Composition HgS 3Na2O.3Al2O3.6SiO2.2Na2S 2PbCO3.Pb(OH)2 As2S3 C 2CuCO3.Cu(OH)2 CuCO3.Cu(OH)2 Obtaining Colours from Pigments Pigments can be spread on a surface layer as in paintings. Until the 19th century, the paint layer of a picture mostly consisted of several layers superimposed one upon the other to achieve complex colour effects, but pigments were rarely mixed with each other, only using white lead to make colours lighter. These pigments can also be used in the production of coloured glass where the finely powdered pigment is directly added to the glass mixture before melting. By mixing the pigment into the bulk of the material, the material is given colour. Cobalt Oxide Blue Manganese Oxide Purple Iron Oxide Green Bright yellow (depends on amount and temperature) For staining, the glass was painted with silver nitrate, and then fired in an oven. Depending on the number of times the glass was stained and fired, a range of yellow tomes from pale lemon to deep orange could be produced. Another method was painting on the glass with a dark pigment of copper or iron oxide mixed with soft powdered glass and wine or urine, which was made permanent by firing in the kiln. Paints Paint consists of a pigment which gives the colour of the paint. The pigment remains insoluble, so that as the medium dries, the colour remains. The binder causes pigment particles to adhere to one another and to the surface. The medium carries the pigment and allows it to be spread over the surface. New Mineral Deposits and the Increasing Range of pigments The first cave paintings used pigments that were easily obtained and extracted from the environment. However, as time progressed new minerals became available for use with artworks due to the discovery of new mineral deposits and improved extraction methods. This is emphasised by comparing the basic pigments used by Aborigines in traditional art and that of the Egyptians and Romans, who had better extraction technologies and therefore, a wider array of pigments available for use in art. An example of this is the pigment, Green Earth. Before it could be extracted, this shade of green had not been seen in artworks. It only came into use when the iron silicate and clay mixture it was found in, could be extracted. Thus, the extraction of pigments and minerals had a direct link to the use of these various pigments in colour. Greater understanding of how colour results from structure of substances have led to a shift from naturally occurring to synthetic compounds. Pigments and the Periodic Table The metallic components of pigments used as a source of colour are predominantly from the transition metal section or d-block of the Periodic Table. It is the electron configuration of these transition metals that contribute to the colour produced (e.g. chromium, manganese and copper). Metallic Component Copper Iron Cobalt Chromium Colour Range Blue – green Red – yellow Yellow – violet Red – yellow Health Risks of Cosmetics Stibnite, Sb2S3, was used as a kohl, mascara, eye shadow and eyeliner. Exposure to antimony for short periods of time can cause nausea, vomiting and diarrhoea. It is a suspected human carcinogen. Malachite was used as an eye shadow and eyeliner. Azurite was also used as an eye shadow. Copper(II) oxide is black and used as a kohl. Copper in high doses can cause anaemia, liver and kidney damage and intestinal irritation. Galena was used as a kohl and mascara. White lead was used as a facial powder or paint. Lead accumulation can damage the nervous system, result in mental retardation and even death. Cinnabar was red and used as rouge, blush, lipstick and facial paint. Mercury poisoning can result in numbness, tunnel vision and brain damage. Pigments over Time (Not really in Syllabus) A dye is a coloured substance that has an affinity to the substrate to which it is being applied. Lakes are natural organic pigments made from precipitation a dye with an inert binder and are usually fugitive. Refined mineral pigment Natural ultramarine – purified from ground lapis lazuli by mixing with wax and kneading in a dilute lye bath. Lazurite crystals preferentially wash out and are collected. Lazurite is a complex sodium calcium aluminium silicate sulfate. Refined organic pigment Madder lake – derived from madder plant’s root (colouring substance: alizarin). The root is dried, crushed and boiled in a weak acid to dissolve the dye. The pigment is made by dissolving this in hot alum solution, then precipitating with soda or borax. Refined inorganic pigment Burnt umber – formed when raw earth umber is calcinated (causes thermal decomposition), resulting in darker shades of colour from the loss of water Synthetic inorganic pigment Prussian blue – formed by the precipitation reaction of potassium ferrocyanide with iron (III) chloride Emerald green – formed by the precipitation reaction of sodium arsenite and copper (II) acetate Synthetic organic pigments Mauvine – formed by oxidising impure aniline with potassium dichromate Alizarin crimson Part 2: The Colours and Spectra of Elements Flame Colour Cation + Na K+ Ca2+ Flame Colour Yellow Lilac Brick Red Cation Ba2+ Sr2+ Cu2+ Flame Colour Apple Green Crimson Red Blue – Green An atom does not radiate energy when its electrons are in their ground states, but electrons can move to other orbits by absorbing or emitting a photon. When an electron jumps from one energy level to a lower energy level, a photon whose energy is equal to the difference between the two energy levels is emitted. The colour of a flame is created by the strong emission lines of the element. The energy of an atom is not continuous, existing only in fixed amounts, with the photon frequencies absorbed or emitted by an atom being fixed by differences between the possible energy transitions between the energy levels of the atom. Emission, Absorption and Reflectance Spectra When light is dispersed into its separate wavelengths, the emitted light passes through a slit then a prism or diffraction grating, which disperses the light into individual wavelength that is seen as spectral lines. Emission spectra result from an excited atom emitting photons characteristic of the atom when it returns to its ground state, which are seen as bright lines on a dark background. Absorption spectra are produced when cool, gaseous atoms absorb photons of certain wavelengths and become excited, and then re-emit the certain wavelength in all directions, so the original light is deficient in these wavelengths. It appears as black lines against a continuous background. Reflectance spectra results from shining white light onto a surface and examining the spectrum of light that is reflected. The Development of the Bohr Model of the Atom The Rutherford model of the atom did not explain the existence of emission spectra as it implied that an electron would spiral smoothly towards the nucleus, releasing a continuous spectrum. Bohr’s model of the atom attempted to provide an explanation for the observation of emission and absorption spectra of elements that suggested that each atom had a different and complex, internal energy structure. Bohr used Planck’s concept of quantisation of energy 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 and that only orbits of certain radii and particular energies were allowed for the electron. He also explained that electrons could not spiral into the nucleus as they could not lose part of a quantum. Every line in the hydrogen spectrum represents a transition from one energy level to another and explained how excited atoms generate emission spectra. However, this model did not match with the experimental results of emission spectra from atoms containing more than one electron. It was also unable to explain why a single emission line for hydrogen was at times composed of 2 or more closely spaced lines, which split even further apart in a magnetic field. In Bohr’s model, the principle quantum number is associated with the radius of the electron’s orbit, which is directly related to the electron’s energy – the lower the principle quantum number, the smaller the radius of the orbit and the lower the energy of the atom. The atom is in its ground state when the electron is at n=1. When an electron jumps from one energy level to a lower energy level, a photon of radiation is emitted, whose energy is equal to the difference in energy between the two levels. For some of these emissions, the energy released is in the visible spectrum range and so can be seen as a visible emission spectrum. Infrared and Ultraviolet Analysis of Pigments Infrared absorption spectroscopy identifies the presence of organic functional groups and allows the molecules to be identified, but is not useful for ionic bonds and requires the dissolution of the sample in a colourless solvent. It is usually used for qualitative analysis, but can be used for quantitative analysis. It is a destructive analytical technique. At room temperature the atoms will vibrate, with the frequency of vibration of atoms in molecules being the same as the range of frequencies of infrared radiation. Vibration absorption occurs when molecules show a dipole moment (i.e. they are IR reactive and polar). The vibrations can be measured by passing infrared light through the sample and recording which wavelengths are absorbed. A particular bond will have approximately the same frequency regardless of the molecule. To identify an unknown substance, the frequency and intensity of each vibration is matched with a standard spectrum of a known substance. The graph is usually plotted as transmission over wavenumber (cm-1). Ultraviolet-Visible Absorption Spectroscopy can be used to determine purity or concentration of a solution of a substance that absorbs UV-Vis light (180-820nm). It can be used as a qualitative tool to identify pigments from their characteristic graph. It is also a quantitative tool as the absorbance of light is directly proportional to the concentration of the pigment in the solution. The concentrations of absorbing materials can be determined by comparing the absorbance with developed calibration curves of the material. It is a destructive analytical technique. This method can be used to identify pigments containing metal ions, but a limitation is that the pigment needs to be in solution. The graph is usually plotted as absorbance over wavelength (nm). Absorption Method Source Detector Measurement Analysis Infrared Ultraviolet Double beam absorption spectrophotometer Heated ceramic such as silicon Tungsten lamp or deuterium discharge carbide rod tube Thermocouple (measures Photomultiplier tube (detects and temperature) amplifies light from faint sources) Frequency is measured in Wavelengths in nanometres -1 wavenumbers (cm ) (1nm = 10-9m) Qualitative and quantitative, but destructive Infrared reflectography is non-destructive and supplies information about the whole painting. In IR reflectance, a lamp with a portion of its output in the cool, near infrared is shone onto the surface of a painting and the reflected light detected. IR radiation penetrates pigments, such as titanium oxide, cadmium sulfide, mercury sulfide and copper containing green pigments, and reflects from the white background. It is particularly good for detecting underdrawings in a painting where the artist has used graphite or charcoal as carbon strongly absorbs IR, producing a black image of the underdrawing and also detects compositional changes. Ultraviolet reflectography involves illuminating the painting with UV from a blacklight. Photons in the UV region are strongly absorbed in the layers of varnish and binding media in the painting. The reflected UV is detected and as different materials absorb UV to varying extents, it can give clues to the composition of pigments. The absorption of ultraviolet photons can cause chemical reactions that can result in the emission of photons in the visible region – fluorescence. The examination of a painting under UV light can reveal retouched areas that have been painted over or restored as aged varnishes and traditional pigments fluoresce differently than new varnishes and modern pigments. UV light can sometimes provide additional information in identifying certain pigments, especially zinc white, which appears yellow in UV light, and cadmium pigments. Relationship between Absorption and Reflectance Spectra Reflectance spectra can only be obtained for opaque coloured objects, so it is used when the substance cannot be dissolved in a colourless solvent. In contrast, to obtain an absorbance spectrum, the substance must be transparent or light is passed through a solution of the substance. Since light can only be absorbed or reflected, the reflectance spectrum is the complement of the absorption spectrum. While obtaining reflectance spectra is non-destructive, obtaining absorption spectra can be destructive. Infrared radiation changes the colour of zinc oxide, in the presence of oxygen, from white to lemon-yellow, although it becomes white again on cooling. However, it permanently breaks down red copper(I) oxide, malachite and verdigris into black copper(II) oxide. Ultraviolet radiation is less harmful than infrared on zinc and copper containing pigments. UV causes some pigments to fluoresce, with zinc oxide fluoresces a yellow colour when it absorbs UV radiation. Malachite fluoresces a dirty mauve colour when exposed to UV light. Emission Lines of Sodium The Zeeman Effect is that in the presence of a magnetic, the spectral lines split into more spectral lines. The stronger the magnetic field, the greater the amount of splitting. Bohr could not explain this effect. The 3p level splits into 2 different energy levels, with total angular momentum j=3/2 and j=1/2, due to the magnetic fields caused by the electrons’ spins and their orbits – spin orbit effect. The line at 589.0nm caused by 3p3/2 has twice the intensity of the line at 589.6nm, which is caused by the 3p1/2. Both these lines of the sodium double represent a transition of electrons from the 3p to 3s energy level. Laser Microspectral Analysis Methodology Material is vaporised when a powerful pulsed laser is focused on a small sample of the pigment to be identified, so it is a destructive analytical technique The vapour is fed through a gap between two electrodes that sparks and excites the atoms and ions, producing an emission spectrum as the excited electrons return to the ground state Importance for Identifying Elements in Compounds The emission spectrum consists of lines corresponding to elements evaporated from the sample surface. It is useful for determining trace elements in solid and liquid samples. It is importance as it has high sensitivity and requires minimal sample preparation. Used to analyse elemental composition of pigments used in the restoration of paintings. A microscopic amount of colour of the painting is vaporised using laser energy. Atomic emission spectroscopy is used to identify and measure the amounts of each element. From this analysis, a synthetic substitute for this colour can be prepared Examples of its Use Analysis of elements in pigments can be used to determine the validity of an artwork Determination of pigments used in artwork before restoration begins Merits and Limitations of Bohr’s Atomic Model Merits Provided an explanation of why electrons did not spiral into the nucleus that Rutherford’s model could not Explained that emission and absorption spectra for different elements had a different and complex energy structure Was successful in predicting the emission spectrum of hydrogen with reasonable accuracy Limitations Not possible to calculate the wavelengths of spectral lines of atoms and ions with more than one electron Does not explain why some spectral lines are more intense (favoured) than others or hyperfine lines (e.g. sodium doublet) Could not account for the splitting of spectral lines whilst in a magnetic field Mixture of classical and quantum physics Part 3: The Periodic Table and Electron Distribution The Pauli Exclusion Principle The Pauli Exclusion Principle states that: An orbital can hold a maximum of two electrons that must have opposite spins It also means that no two electrons in the same atom can have the same set of quantum numbers. Each electron has 4 quantum numbers denoting its shell, sub-shell, orbital and spin. Sub-shells Each shell contains of a number of energy sublevels with slightly different energies which are called subshells. Subshells are a consequence of nucleus-electron attraction and electron–electron repulsion. The order of sub-shell filling can be represented in the following ways: Electrons in their ground-state electron configurations occupy the lowest energy shells, sub-shells and orbitals available to them. 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 levels, which can be by electrical excitation or by the absorbance of electromagnetic radiation. Organisation of the Periodic Table The Periodic Table is organised into blocks that reflect the filling of the outermost subshells with electrons. Atoms with similar electron configurations in their valence shell display similar chemical properties and are represented as a vertical group. The s block relates to groups 1 and 2, the p block corresponds to groups 3 to 8. The d block relates to the transition metals, while the f block corresponds to the lanthanide and actinide metals, with this arrangement reflecting the filling of the respective subshells. Electronegativity Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. Factors affecting electronegativity include: The number of protons in the nucleus The distance from the nucleus The amount of screening by inner electrons Electronegativity increases across periods in the periodic table from left to right, due to an increase in nuclear charge across a period and as atomic radius decreases. This reflects the trend across a period for elements to gain electrons and achieve a noble gas configuration. Electronegativity decreases down groups of the periodic table as when the size of the atom increases, the attractive force of the nucleus is shielded by electrons in inner shells. Ionisation Energy The ionisation energy of an atom or ion is the amount of energy needed to remove the outermost electron from a mole of gaseous atoms or ions. Ionisation energy is also affected by the charge on the nucleus, the distance of the electron from the nucleus and the amount of screening electrons. Across a period, the first ionisation energy increases because of the greater attraction to the nucleus from the extra proton, while the extra electron is located a similar distance away from the nucleus. Down a group, there is a gradual decrease in ionisation energy as the outer electrons are further from the nucleus. The successive ionisation energies of atoms have larger values because the electrons are being removed from progressively larger positive charges. The removal of an electron from a shell closer to the nucleus leads to a substantial increase in ionisation energy. Thus, trends in successive ionisation energies of an element can be used to predict the number of electrons in the shells of an element’s atoms. Boron has a first ionisation energy lower than expected as the outermost electron is shielded by the 1s and 2s electrons, so it requires less energy to remove than beryllium. Oxygen has a lower first ionisation energy than nitrogen as the additional electron must be paired up in the same orbital as another electron, which results in greater electrostatic repulsion, making it easier to remove. Hund’s Rule If two or more orbitals of the same energy (subshell) are available, 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 as it produces a configuration of low energy. The mutual repulsion from electrons in the same orbital would result in a higher potential energy than filling of orbitals according to Hund’s rule. Part 4: Transition Metals and Electronic Configurations Transition Metals The transition metals occupy the d block of the periodic table. Transition elements are elements that form at least one ion with a partially filled d subshell. Each d subshell includes 5 orbitals and can accommodate up to 10 electrons, so each transition series consists of 10 elements. Transition metals can exhibit a variety of oxidation states as they can lose electrons from both the 3d and 4s subshells, which have similar energies. The +2 oxidation state that occurs commonly for nearly all transition metals is due to the loss of their two 4s elements. Oxidation states above this result from the additional loss of 3d electrons, which are of similar energies to the 4s electron, when simple charged ions are formed, or when the transition metal is bonded to more electronegative element, such as oxygen. The maximum oxidation state for a transition element corresponds to the total number of 4s and 3d electrons in the atom. Colour Changes and Changing Oxidation States The ability of transition metal compounds to absorb light energy is due to the fact that the d orbitals in transition metal ions can have slightly different energies and are incompletely filled. The small energy differences between the d orbitals are similar to the energies of the photons of visible light, thus the absorption of photons of an appropriate frequency can result in an electron being excited from a lower to a higher energy orbital. The absorption of certain components of white light means that the compound appears coloured and exhibits a colour that is the complement of the wavelengths absorbed. As each ion of a transition metal in a different oxidation state has a different arrangement of filled and unfilled 3d orbitals, they absorb different energies from the range of visible light, giving them 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+) Complex Ions, Transition Metals and Strong Oxidising Agents Transition metal ions in which the metal has a high oxidation state tend to be strong oxidising agents. Complex ions, such as the chromate ion (CrO42-), dichromate ion (Cr2O72-) and permanganate ion (MnO4-), are strong oxidising agents as they contain many oxygen atoms, which only have 6 electrons in the outer shell and are able to readily accept electrons. Transition metal ions in these complexes have high oxidation states and have lost many electrons, giving it a smaller radius that gives greater attraction for electrons and thus, a higher oxidising strength. Part 5: Coloured Complex Ions Hydrated Ions Hydration is when an ionic solid dissolves in water and the ions dissociate and are surrounded by water molecules that are oriented with their positively charged ends directed towards the anions and their negative charged ends towards the cations. In a complex ion, a central metal ion is surrounded by a group of anions or molecules called ligands. Hydrated ions are an example of a complex ion where water is the ligand. Coordinate covalent bonds form between the ligands and the transition metal ion using the electron lone pairs on the ligands (e.g. [Cu(NH3)42+], [Cr(H2O)6]3+) Ligands The number of coordinate covalent bonds between the ligands and the central metal ion is the coordination number. Ligands are all Lewis bases and have at least one unshared pair of electrons that can be used to form a coordinate covalent bond to a metal ion. Ligands which bond using the electron pair of a single donor atom (e.g. Cl-, NH3 and F-) are monodentate ligands, while ligands that bond through electron pairs on more than one donor atom are said to be polydentate. Polydentate ligands are also known as chelating agents because of their multipoint attachment to a metal ion. The donor atoms of chelated ligands bond simultaneously to the central metal ion. (e.g. bidentate: oxalate ion –OOC-COOand ethylenediamine; quadridentate: haemoglobin; hexadentate: EDTA) Importance of Models Models are important for explaining: how metal ligands form, why certain geometrical shapes are preferred, why these compounds are brightly coloured and often magnetic and why some ligands can chelate. Naming Complex Ions Ligand Coded by H2O Aqua NH3 Ammine OH Hydroxo Cl Chloro F Fluoro CNCyano No of ligands 1 2 3 4 5 6 Coded by Mono Di Tri Tetra Penta Hexta Metal Cobalt Aluminium Chromium Vanadium Copper Iron Negative ion name Cobaltate Aluminate Chromate Vanadate Cuprate Ferrate Valence Bond Theory This explains the bonding in coordination complexes, which depends on the orbitals available for coordinate covalent bond formation, the tendency of the ions or groups to share a pair of electrons, the coordination number of the complex and the geometry assumed by the ligands. The basis of Valence Bond Theory is that atomic orbitals are reorganised for form hybrid orbitals and attempts to explain the shape and magnetic properties of these complexes. e.g. Mn2+ + 4Cl- [MnCl4]2A pair of electrons on each chlorine ion can form a coordinate covalent bond with the 4s and 4p orbitals on the manganese ions, forming a sp3 complex, which is tetrahedral in shape. e.g. Co3+ + 6NH3 [Co(NH3)6]3+ The pairs of electrons from the ammonia molecules are bonded in the two d, one s and three p orbitals of the metal ion, which is a d2sp3 complex, which is octahedral in shape. Since only the inner d orbitals are used in the complex, it is an inner orbital complex. The field strength of the ammonia ligand causes forced pairing, allowing for the use of the inner d orbitals. e.g. Fe3+ + 6H2O [Fe(H2O)6]3+ The pairs of electrons from the water molecules are bonded in the one s, three p and two d orbitals of the metal 3 2 ion, which is a sp d complex, which is octahedral in shape. Since a set of outer d orbitals are involved in the hybridisation, it is an inner orbital complex. Complexes that have no unpaired electrons are diamagnetic and repelled by a magnetic field. Complexes that have unpaired electrons are paramagnetic and attracted by a magnetic field. Crystal Field Theory This theory was devised to interpret the colours of the ions and their magnetic properties. It assumes that the metal ion is a positive point charge and the ligand is a negative point charge, and the bonds between them are completely ionic. The d orbitals of an atom in the absence of an electrical field are of the same energy and said to be fivefold degenerate. However, the field produced by attached ligands is concentrated in particular directions, so the degenerate levels are separated into two or more sets of levels. The magnitude of the separation and the number of orbitals in the new levels depend on the type of arrangement of the ligands and will be different according to their geometry. e.g. Fe3+ + 6H2O [Fe(H2O)6]3+ In this complex ion, the six water ligands are in an octahedral arrangement and point directly towards the 3dx2-y2 and 3dz2 orbitals, increasing repulsion and raising the energy of these two orbitals more than the other three d orbitals, leading to the separation of the d orbitals. As the water molecules have a weak crystal field around the iron(III) ion, it has a smaller separation between the lower and higher d energy levels. As a result, the energy used to promote two electrons to higher levels is less than would be needed to overcome the repulsion of another electron in two lower energy levels, if pairing was to occur. This type of complex is called a high spin complex, having many unpaired electrons, which accounts for their paramagnetic properties. e.g. Co3+ + 6NH3 [Co(NH3)6]3+ As a result of the strong crystal field strength of ammonia, a large separation between the energies of the lower and higher d orbitals result. 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 higher orbitals, so the electrons pair in the lower orbitals. This type of complex is called a low spin complex, having few unpaired electrons, accounting for their diamagnetic properties. Whether these complexes are low spin or high spin depend on the strength of the field established by the ligands around the central ion, the energies of the d orbitals of the central ion and the overlap between the orbitals of the central ion and the ligands. Ligands arranged in a tetrahedral pattern (e.g. [MnCl4]2- ) will have a triply degenerate upper d level as ligands point directly towards the 3dxy, 3dxz and 3dyz orbitals, giving a doubly degenerate lower level. Ligand CNNO2En NH3 SCNC2O42H2O OHFClBrI- The crystal field theory can be used to explain the colours of transition metals. Electrons in the lower d levels may absorb photons and move to higher d levels, with the transmitted light from the complement of the absorbed photons. Relative Strength Strong field ligands Moderate field ligands Weak field ligands