Download Chemistry of Art by Jonathan Chan

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