Download Chemistry of Art by Daniel Yang

THE CHEMISTRY OF ART
1. From earliest times, people have used colour to decorate themselves and their
surroundings
- Identify the sources of the pigments used in early history as readily available minerals
While organic pigments have been used in history, such as saffron, indigo, madder etc. that are extracted from
plants and animals, in early history inorganic pigments from readily available minerals such as ochre and malachite
have predominately been used.
- Explain why pigments used needed to be insoluble in most substances
While a pigment is an insoluble solid suspended in the binder, a dye is colouring matter dissolved in solution. The
pigment in the binder needs to be insoluble so that as the binder dries, the pigment is left behind.
- Outline the early uses of pigments for:
Cave drawings
The oldest known paintings from over 17000 years ago are Aboriginal rock paintings, and the Lascaux cave paintings
in France from about 15000 years ago. The pigments used were derived from natural sources such as ochre and
charcoal, and used saliva, fats, waxes, water or blood as the binder.
Self-decoration including cosmetics
Egyptians used kohl as eyeliner and mascara, and used different pigments as eyeshadow such as orpiment and
verdigris. Ancient Greeks used white lead as a face paint.
Preparation of the dead for burial
In the Stone Age, humans began burying their dead with red ochre placed on the head and chest of the corpse.
Egyptians painted their dead with pigments in the same way as they were used for the living, sometimes painting the
sarcophagus with the history of their life.
- Explain that colour can be obtained through pigments spread on a surface layer (eg paints) or mixed with the
bulk of material (eg glass colours)
Canvas or panel
Ground: Bottommost layer on support in preparation for painting. For panels, it is usually gesso (gypsum or chalk
with animal glue).
Glass
Finely powdered pigments were added to the glass mixture before melting, or glass was “stained” by painting the
glass with silver nitrate and then firing it in an oven. The number of times the glass was stained then fired gave rise
to a range of colours from pale lemon to deep orange.
- Outline the processes used and the chemistry involved to prepare and attach pigments to surfaces in a named
example of medieval or earlier artwork
St John the Baptist with St John the Evangelist and St James, Nardo di Cione (circa 1365)
- Panel painting using egg tempera
- Panel painting on poplar wood, coated with gesso and underdrawing
- Painting is gilded (coated with gold) underneath
- Pink pigment is crimson lake, probably using cochineal, with layers of white lead. Blue pigment is natural
ultramarine. Green pigment is ultramarine mixed with lead tin yellow. Scarlet pigment is vermillion.
- Describe paints as consisting of:
The pigment
The pigment adds the colour to the paint. It must be chemically inert and should be unaffected by light, heat, acids,
alkalis, moisture etc.
A liquid to carry the pigment
The pigment is suspended in a binder that forms a film when dry, and allows the pigment to adhere to the canvas or
wood. Oil paints use a drying oil such as linseed oil, watercolours use gum arabic, acrylic colours use acrylic resin and
egg yolk is mixed with water to form egg tempera.
- Describe an historical example to illustrate the relationship between the discovery of new mineral deposits and
the increasing range of pigments
In 1770, an orange mineral was found in the Beresof gold mine in Siberia. It was identified by French chemist
Vauquelin as a compound of lead and a new element, chromium. The discovery of chromite (FeO∙Cr2O3), a chromium
ore, led to the development of many chromium pigments such as chrome yellow (PbCrO4), chrome green (chrome
yellow mixed with Prussian blue) and chrome red (PbCrO4∙Pb(OH)2).
- Analyse the relationship between the chemical composition of selected pigments and the position of the metallic
component(s) of each pigment in the Periodic Table
Many pigments are composed of metals that are transition metals.
- Identify data, gather and process information from secondary sources to identify and analyse the chemical
composition of an identified range of pigments
Aboriginal
Pigment
Chemical name
Colour
Chemical formula
Red ochre
Yellow ochre (goethite)
Charcoal
Chalk
Anhydrous iron (III) oxide
Hydrated iron (III) oxide
Graphite
Calcium carbonate
Red
Yellow
Black
White
Fe2O3
Fe2O3∙H2O
C
CaCO3
Pigment
Chemical name
Colour
Chemical formula
Cinnabar
Orpiment
Galena
White lead
Malachite
Mercury (II) sulfide
Arsenic trisulfide
Lead (II) sulfide
Basic lead carbonate
Basic copper carbonate
Red
Yellow
Black
White
Green
HgS
As2S3
PbS
2PbCO3∙Pb(OH)2
CuCO3∙Cu(OH)2
Egyptian/Greek
Post 18th Century
Pigment
Chemical name
Colour
Chemical formula
Smalt
Prussian blue
Scheele’s green
Chrome yellow
Chrome red
Cobalt blue
Cobalt green
Cobalt yellow
Dark cobalt violet
Cadmium yellow
Cadmium red
Titanium white
Zinc white
Cobalt (II) oxide
Ferric hexacyanoferrate
Acid copper (II) arsenite
Lead (II) chromate
Basic lead (II) chromate
Cobalt-aluminium oxide
Cobalt-zinc oxide
Cobalt-potassium nitrite
Cobalt phosphate
Cadmium sulfide
Calcium sulfide selenide
Titanium dioxide
Zinc oxide
Blue
Blue
Green
Yellow
Red
Blue
Green
Yellow
Violet
Yellow-orange
Red
White
White
CoO
Fe4(Fe(CN)6)3
CuHAsO3
PbCrO4
PbCrO4.Pb(OH)2
CoO.Al2O3
CoO.nZnO
CoK3(NO2)6.H2O
Co3(PO4)2
CdS
CdS.CdSe
TiO2
ZnO
- Process information from secondary sources to identify the chemical composition of identified cosmetics used in
an ancient culture such as early Egyptian or Roman and use available evidence to assess the potential health risk
associated with their use
Cinnabar, HgS – Mercury compound which is poisonous. Used as rouge, lipstick.
Orpiment, As2S3 – Arsenic compound which is toxic. Used as eyeshadow.
Galena – PbS – Absorbed through the skin leading to lead poisoning. Used as eyeliner/mascara/kohl.
White lead - 2PbCO3∙Pb(OH)2 – Absorbed through the skin leading to lead poisoning. Used as facial powder.
2. By the twentieth century, chemists were using a range of technologies to study the
spectra, leading to increased understanding about the origins of colours of different
elements
- Identify Na+, K+, Ca2+, Ba2+, Sr2+, and Cu2+ by their flame colour
Na+: Yellow
K+: Violet
Ca2+: Brick red
Ba2+: Apple green Sr2+: Crimson
Cu2+: Blue-green
- Explain the flame colour in terms of electrons releasing energy as they move to a lower energy level
When electrons of a metal are placed into a flame, they absorb heat energy and jump to a higher energy level.
However, these electrons jump back down for stability and release the energy difference between the two energy
levels as a photon.
- Explain why excited atoms only emit certain frequencies of radiation
Electrons of a particular element can only absorb energy corresponding to specific frequencies. When the electrons
drop back down to the lower level, they only emit certain frequencies corresponding to the energy differences
between the two energy levels. Each atom has a unique set of possible energy levels, thus making energy jumps
unique.
- Distinguish between the terms spectral line, emission spectrum, absorption spectrum and reflectance spectrum
A spectral line is a single line or wavelength, compared with a continuous spectrum.
An emission spectrum is produced by an excited atom, each emitted wavelength appearing as a discrete line on a
black background.
An absorption spectrum is produced when a continuous spectrum is passed through a cool, unexcited gas. The
electrons in the gas absorb specific wavelengths and re-emit them in all directions, effectively leaving the original
spectrum deficient in those wavelengths. The absorption spectrum appears as dark lines on a continuous
background.
The reflectance spectrum is the complement of the absorption spectrum and represents the visible colour.
- Describe the development of the Bohr model of the atom from the hydrogen spectra and relate energy levels to
electron shells
In 1919, Neils Bohr used Planck’s quantum theory to explain the hydrogen spectrum. He proposed that in a hydrogen
atom, the electron moves around the nucleus in a circular orbit. He also proposed that electrons can only exists in
specific quantised energy states; only orbits of particular energies were allowed. Each line of the hydrogen spectrum
is produced when an electron moves from one energy level to another.
However, the limitations of Bohr’s model included:
- Did not explain the relative intensities of the spectral lines
- Could not be sufficiently extended to atoms with more than one electron
- Could not explain that emissions lines were actually sometimes composed of two or more closely spaced lines
- Could not explain the splitting of lines in a magnetic field (Zeeman effect)
- Explain what is meant by n, the principal quantum number
The principal quantum number n corresponds to the energy level or energy shell. n=1 is the energy level closest to
the nucleus followed by n=2, n=3 etc. It was the only quantum number in Bohr’s model, with one electron with a
ground state of n=1.
- Identify that, as electrons return to lower energy levels, they emit quanta of energy which humans may detect as
a specific colour
When an excited electron returns from an excited energy level to a lower energy level, a photon of light is emitted
with energy equivalent to the energy difference between the two energy levels. If the photon emitted has frequency
inside the visible range (wavelength of about 380-750 nm), then it will be visible light.
- Outline the use of infra-red and ultra-violet light in the analysis and identification of pigments and their chemical
composition
Both IR and UV-Vis spectroscopy are destructive techniques and both use double beam spectrophotometers. In
double beam spectrophotometers, the radiation is split into two beams; one passes through the sample and the
other is used as a reference. The intensities of the two beams are compared allowing absorbance (or transmittance)
to be determined.
UV-Vis spectroscopy
UV-Vis spectroscopy is mainly used to identify pigments consisting of metal ions since it is based on the excitation of
electrons which are promoted to higher energy level. The source of UV-Vis spectrophotometers is a tungsten lamp
or deuterium discharge tube, and the detector is a photomultiplier tube. Absorbance is then plotted against
wavelength (nm) and a calibration graph used to identify pigments.
IR spectroscopy
IR spectroscopy is mainly used to identify pigments consisting of organic molecules since it is based on molecular
vibrations and stretching of bonds such as C-H, C-C and C=C. The source of IR spectrophotometers is commonly a
heated ceramic such as silicon carbide and a thermocouple is used as the detector. Transmittance is then plotted
against wavenumber (cm-1) and the peaks can be identified as specific bonds.
IR reflectography
IR reflectography is a non-destructive technique to detect the underdrawing of a painting. While the underdrawing is
usually sketched in carbon-based substances such as graphite or charcoal, the paints usually consist of ionic
pigments. Infra-red radiation is shone onto the painting, and is not absorbed by the upper layers and thus penetrates
to the underdrawing. However, the covalent bonds in the underdrawing absorb UV light and thus do not reflect it
back. The reflected IR light is then detected by a thermographic camera and an IR reflectogram created, and the dark
areas indicate where the underdrawing is.
- Explain the relationship between absorption and reflectance spectra and the effect of infra-red and ultra-violet
light on pigments including zinc oxide and those containing copper
The reflectance spectrum is the complement of the absorption spectrum and represents the visible colour. They are
complementary; a yellow pigment absorbs violet etc.
Reflectance spectra can be analysed instead of absorption spectra by shining white light onto the pigment and then
using the reflectance spectra to compare it to known reflectance spectra to identify the pigments, which is nondestructive compared to absorption spectroscopy which is destructive.
Effect of IR on ZnO and Cu compounds
Infra-red radiation (heat) causes zinc white, ZnO, to turn yellow. However, when the pigment cools, it returns to
white. Infra-red also permanently causes red Cu2O and green CuCO3 to decompose to black copper (II) oxide.
- Gather and process information from secondary sources to analyse the emission spectra of sodium and present
information by drawing energy level diagrams to represent these spectral lines
In sodium, two electrons in the 3p orbital experience spin-orbit splitting, which
causes shifts in an electron's atomic energy levels due to electromagnetic
interaction between the electron's spin and the nucleus's magnetic field. Thus,
the two electrons of the 3p orbital have slightly different energies.
The transition between the 3p energy level and the 3s energy level causes two
lines; the D1 line of 589.592 nm and the D2 line of 588.995 nm, known as the
sodium doublet. The two energy levels are split even further apart when an
external magnetic field is applied, known as the Zeeman effect.
- Gather, process and present information about a current analytical technology
Laser microspectral analysis
A high-energy pulse of laser light vaporises a small sample, which is then fed through a gap between two electrodes
that sparks and excites the atoms and ions of the sample, causing them to release emission spectra. A
spectrophotometer is then used to identify the pigments using the emission spectra. While it is a destructive
technique, only a minute amount of the sample is vaporised. It can identify many elements at a time, is highly
sensitive, is minimally destructive, requires minimal sample preparation, only requires optical contact with the
painting, can test many objects in a short amount of time and can be used in-situ as it can be transported. It is used
to determine the authenticity of the painting by analysing the pigments used, and is also used for restoration as a
synthetic substitute pigment can be prepared by knowing the chemical composition of the original pigment used by
the artist.
3. The distribution of electrons within elements can be related to their position in the
Periodic Table
- Define the Pauli exclusion principle to identify the position of electrons around an atom
The Pauli exclusion principle states that no two electrons in the same atom may occupy the same quantum state
simultaneously – i.e. no two electrons in an atom can have identical sets of four quantum numbers. This means that
if two electrons share an orbital, they must have opposite spin.
Quantum numbers: principal quantum number, n (electron shell)
azimuthal quantum number, ℓ (subshell)
magnetic quantum number, mℓ (orbital)
spin projection quantum number, ms (spin)
- Identify that each orbital can contain only two electrons
Since there are only two spin projection quantum numbers, ½ and -½, each orbital can have a maximum of two
electrons due to the Pauli exclusion principle and they must have opposite spins.
- Define the term sub-shell
Each energy shell consists of a number of energy sublevels called sub-shells, with slightly different energies. The first
four subshells are s (one orbital, two electrons), p (three orbitals, six electrons), d (five orbitals, ten electrons) and f
(seven orbitals, fourteen electrons).
- Outline the order of filling of sub-shells
Electrons are filled into the sub-shells from lower to higher energy.
Hund’s rule of maximum multiplicity states that electrons are added to the
orbitals of a subshell with parallel spin until all the orbitals are half-full,
then with opposite spin to a half-filled orbital.
- Identify that electrons in their ground-state electron configurations occupy the lowest energy shells, sub-shells
and orbitals available to them and explain why they are able to jump to higher energy levels when excited
In their ground state, electrons occupy the lowest energy shells, sub-shells and orbitals available. In an excited state,
electrons have more energy and can jump to a higher energy level by absorbing the energy.
- Explain the relationship between the elements with outermost electrons assigned to s, p, d and f blocks and the
organisation of the Periodic Table
Elements in the same group have similar outer shell electron configurations.
- Explain the relationship between the number of electrons in the outer shell of an element and its
electronegativity
Electronegativity is a measure of the tendency of an atom to attract electrons towards itself.
From left to right across a period, electronegativity increases as electrons are added to the same outer shell but the
nuclear charge increases.
From top to bottom down a group, electronegativity decreases as electrons are added to an increasingly further
outer shell. The increased distance of the valence electrons from the nucleus reduces the attractive force to the
nucleus, and electron shielding from the inner shells shield the valence electrons from the nuclear attraction.
- Describe how trends in successive ionisation energies are used to predict the number of electrons in the
outermost shell and the sub-shells occupied by these electrons
Ionisation energy is the amount of energy required to remove an electron from a mole of atoms in gaseous state.
Large increases in successive ionisation energies indicate that an electron is now being removed from a full shell; it
requires much less energy to remove a single electron from a shell than to remove an electron from a full electron
shell. In general, if the nth ionisation energy is much larger than the previous, then we can predict that the atom has
n-1 electrons in its outermost shell. Furthermore, by knowing the electron configuration of an atom, we can
determine its subshell structure using the order of filling.
- Process information from secondary sources to analyse information about the relationship between ionisation
energies and the orbitals of electrons
Down a group, the first ionisation energy decreases as the electronegativity decreases and less energy is required to
remove the first electron.
Across a period, the first ionisation energy increases as the electronegativity increases and more energy is required
to remove the first electron.
In the first period, however, there are two instances in which the general trend is not observed:
Be (1s2 2s2) to B (1s2 2s2 2p1) – the 2p1 electron being removed from B is further away from the nucleus than the 2s
sub-shell, and is also party shielded by the 2s electrons, so less energy is required to remove it.
N (1s2 2s2 2p3) to O (1s2 2s2 2p4) – the three electrons in the 2p sub-shell of N occupy separate orbitals in accordance
with Hund’s rule. However, the extra electron in oxygen must be paired with another electron in an orbital, resulting
in greater electrostatic repulsion, and hence less energy is required to remove it.
4. The chemical properties of the transition metals can be explained by their more
complicated electronic configurations
- Identify the block occupied by the transition metals in the Periodic Table
The transition metals are a group of metals that occupy the d-block of the periodic table.
- Define the term transition element
A transition metal is an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an
incomplete d sub-shell.
- Explain why transition metals may have more than one oxidation state
When electrons are back-filled into the 3d sub-shell, they shield the 4s electrons from the nucleus leaving the 4s
electrons more exposed. Transition metals exhibit a variety of oxidation states since they can lose electrons from
both the 3d and 4s sub-shells, which have similar energies. The 2+ oxidation state which occurs for nearly all of the
transition metals is due to the loss of the two 4s electrons. Oxidation states above 2+ are a result of the loss of 3d
electrons as well.
- Account for colour changes in transition metal ions in terms of changing oxidation states
The different oxidation states of transition metals arise from a variety of non-degenerate d-orbitals. Each transition
metal ion in a different oxidation state has a different arrangement of its 3d orbital, which leads to a different set of
possible electron jumps, and hence a different set of frequencies of light that can be absorbed. As such, the
reflectance spectrum of each oxidation state is different and it thus appears a different colour.
- Explain, using the complex ions of a transition metal as an example, why species containing transition metals in a
high oxidation state will be strong oxidising agents
In general, transition metal ions with a high oxidation state tend to be strong oxidising agents. Species containing
transition metals in a high oxidation state readily accept electrons (and are thus oxidants). Complex ions such as
dichromate Cr2O72- (oxidation state +6) and permanganate MnO4- (oxidation state +7) are strong oxidants; such
species will have a high tendency to accept electrons and thus be oxidising agents.
- Process and present information from secondary sources by writing electron configurations of the first transition
series in terms of sub-shells
The electron configurations of the first
transition metal series generally follows
Hund’s rule and fills normally. However,
there are two exceptions to the expected
filling pattern:
The reason for the unexpected electron configuration of Cr and Cu is the fact that a half-filled or completely filled 3d
orbital is more stable than if the electron were to go into the 4s orbital.
- Perform a first-hand investigation to observe the colour changes of a named transition element as it changes in
oxidation state
Acidified vanadate (VO3-), prepared by the dissolution of 3g NH4VO3 and 100 mL 1M NaOH in a conical flask followed
by acidification by adding 75 mL 2M H2SO4, was swirled in a conical flask with zinc.
All oxidation half-equations were:
Zn(s) ⇌ Zn2+(aq) + 2eYellow (+5) to blue (+4):
VO3-(aq) + 4H+(aq) + e- ⇌ VO2+(aq) + 2H2O(l)
Blue (+4) to green (+3):
VO2+(aq) + 2H+(aq) + e- ⇌ V3+(aq) + H2O(l)
Green (3+) to violet (+2):
V3+(aq) + e- ⇌ V2+(aq)
- Choose equipment, perform a first-hand investigation to demonstrate and gather first-hand information about
the oxidising strength of KMnO4
Equal volumes of 0.01M KMnO4 and 1M H2SO4 were mixed in a beaker to prepare acidified permanganate. The
permanganate was then added to KCl, KBr, KI, Mg, Zn, Sn, Fe and Cu in test tubes. All solutions turned colourless
except with KCl, which remained purple (similar oxidation potential).
MnO4-(aq) + 8H+(aq) + 5e(purple)
MnO4- (7+): purple
MnO42- (6+): green
MnO2 (4+): brown
Mn2+ (2+): colourless
Mn2+(aq) + 4H2O(l)
(colourless)
5. The formation of complex ions by transition metal ions increases the variety of coloured
compounds that can be produced
- Explain what is meant by a hydrated ion in solution
When an ionic solid dissolves in water, the ions dissociate and are surrounded by water molecules, which orient with
their negatively charged end (O) towards cations and their positively charged end (H) towards anions. This is a
hydrated ion.
- Describe hydrated ions as examples of a coordination complex or a complex ion and identify examples
Hydrated ions are examples of complex ions, in which a central metal ion is surrounded by a ligand. The unpaired
electrons on the ligand are donated to the central ion, forming a coordinate covalent bond. Examples of ligands
include H2O, Cl-, CN- and NH3. The number of coordinate covalent bonds that the central ion can have with a ligand
is called its coordination number.
- Describe molecules or ions attached to a metal ion in a complex ion as ligands
A ligand is an ion or molecule that binds to a central metal ion to form a coordination complex or complex ion.
- Explain that ligands have at least one atom with a lone pair of electrons
For a ligand to form a complex ion with a metal ion, it must donate a lone pair of electrons to the metal ion. As such,
most ligands act as a Lewis base, and the metal ion acts as a Lewis acid.
Monodentate ligands bond with only one electron pair, while ligands that bond through electron pairs on more than
one donor atom are called polydentate ligands or chelating agents.
- Identify examples of chelated ligands
- Discuss the importance of models in developing an understanding of the nature of ligands and chelated ligands,
using specific examples
In valence bond theory, coordinate covalent bonds are formed when two electrons from a ligand are placed into
vacant hybridised orbitals of a transition metal ion. VBT is useful for predicting and explaining the shape of
coordination complexes and their magnetic properties.
VBT provides no insight into the origins of colour of transition metal complexes and the need for an explanation of
colour resulted in the development of an alternative model of coordination complexes, called crystal field theory.
Crystal field theory assumes that the interaction between a transition metal ion and a ligand is electrostatic in
nature, arising from the attraction between the positively charged metal cation and the negatively charged ligand. As
the ligand approaches the metal ion, the electrons in the d-orbitals and the electrons in the ligand repel each other,
causing the orbitals to become non-degenerate. This results in d-orbital splitting, affected by:
- The nature of the metal ion.
- The metal's oxidation state. A higher oxidation state leads to a larger splitting.
- The arrangement of the ligands around the metal ion.
- The nature of the ligands surrounding the metal ion. The stronger the ligand, the greater splitting of the d-orbital,
and the higher the difference between the high and low energy d-orbitals.
The most common type of complex is octahedral; six ligands form an octahedron around the metal ion. Due to the
splitting of the d-orbital energy levels, the coordination complex now has a difference in energy levels that can
absorb energy to cause an electron to jump to the higher level. Because only certain wavelengths of light are
absorbed – those matching exactly the energy difference – the compounds appear the complementary colour.
- Process information from secondary sources to give an example of the range of colours that can be obtained
from one metal such as Cr in different ion complexes
Chromium complex ion
Hexaaquachromium (III)
Hexaaquachromium (II)
Pentaaquasulfatochromium (III)
Tetraaquadicholorochromium (III)
Hexahydroxochromate (III)
Hexaamminechromium (III)
Chemical formula
3+
[Cr(H2O)6]
[Cr(H2O)6]2+
[Cr(H2O)5SO4]+
[Cr(H2O)4Cl2]+
[Cr(OH)6]3[Cr(NH3)6]3+
Colour
Violet
Blue
Dark green
Dark green
Green
Mauve