Download Chemistry of Art by Ian Wilkinson

HSC Chemistry Module 9.8 Summary
1. From the 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
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The pigments used in artworks (cave drawings) from early history mostly consisted of
minerals from coloured earth and soft rocks, before the technology was available to
manufacture synthetic colours
The predominant colours seen in ancient artwork are red, yellow, black, and white, as these
colours came from readily available minerals from the earth
o Red ochre (Fe2O3), or haematite, can be easily extracted from clay to produce a red
pigment
o Charcoal (C), or graphite, can be extracted from ash to produce a black pigment
o See below for a more detailed list on pigments sourced from readily available
minerals used in ancient artwork
Solve problems and perform a first-hand investigation or process information from
secondary sources to identify minerals that have been used as pigments and
describe their chemical composition with particular reference to pigments
available and used in traditional art by Aboriginal people
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Colour
Red
Aboriginal people used “earth” colours in traditional art (i.e. colours that were readily
available from the environment)
o Such colours included red, yellow, brown, white, and black.
o The various ochres were readily available from the ground, as well as white clay and
manganese (IV) oxide
o Charcoal could be obtained through burning grass or bark
Pigment
Red ochre
Mineral
Haematite
Chemical composition
Fe2O3
Yellow
Brown
White
Black
Yellow ochre (goethite)
Brown ochre (limonite)
White clay
Charcoal
Manganese (IV) oxide
Goethite
Limonite
Kaolin
Graphite
Pyrolusite
Fe2O3.H2O
FeO(OH)
Al2O3.2SiO2.2H2O
C
MnO2
Describe paints as consisting of:
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the pigment
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a liquid to carry the pigment
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Paint consists of a binder, a pigment, and a medium
The pigment provides the colour in paint, and remains on the surface after the medium has
evaporated
o Extender pigments can also be added, which can develop certain characteristics of
the paint, such as gloss
The binder, also caused the resin, is the adhesive that binds the pigments together, which
allows the pigment to remain as a film after the medium has dried
o Binders include egg yolk, urine, blood, saliva, tree gums, and honey
The medium is the liquid solvent that carries the paint so it can be spread over a surface, and
allows the pigment and binder to flowing together
o The medium used affects the film thickness of the paint and drying time
o The liquid used must be viscous enough to prevent the paint from running, but not
so thick that it restricts the artist’s work
Explain why pigments used needed to be insoluble in most substances
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As discussed above, paints consist of insoluble pigments in a colourless liquid medium
Pigments need to be insoluble in the liquid medium so that they remain on the surface of
the artwork after the medium has evaporated
o If a pigment is soluble in the medium (such as dyes), the pigment also evaporates,
thus in such a case no colour would remain
Insoluble pigments are also not easily removed when exposed to rain or ground water,
which is advantageous for cave or rock paintings
In addition, cosmetics made from insoluble minerals do not dissolve in perspiration, which is
particularly useful in hot climates
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)
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As described above, paint consists of a pigment within a binder and liquid medium
Paint can be spread on a surface layer, which can adhere to the surface due to the adhesive
properties of the binder (such as egg yolk) after the liquid medium has evaporated
o See below for a more detailed description of the preparation of panel painting
(under mediaeval artwork)
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Pigments can also be mixed with the bulk of material to produce colour, such as in coloured
glass, rather than just the surface layer
Metal oxide pigments can be finely grounded and added to the glass mixture before melting,
which results in coloured glass
o For example, cobalt oxide can be added to produce blue glass
Another method to produce coloured glass is through staining
o The glass is painted with silver nitrate, then fired in an oven
o Depending on the number of time the glass is stained then fired, a range of yellow
tones could be produced, from pale lemon to deep orange
Outline the early uses of pigments for:
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cave drawings
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self-decoration including cosmetics
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preparation of the dead for burial
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CAVE DRAWINGS
The oldest know paintings are Aboriginal cave paintings from around 17 000 years ago, and
the Lascaux cave paintings in France from around 15 000 years ago
The predominate colours in cave drawings are red, yellow, black, and white (see above for
the relevant pigment names and chemical formulae)
o These pigments used to produce these colours came from readily available minerals,
as technology was too limited to produce a wider range of colours
The pigments were first grinded to a fine powder, then either applied directly to the wall as
a solid mixed with binder in a medium as a paint
o The paints used in cave drawings generally used saliva, honey from wild bees, or
tree gums as the binder, and water as the medium
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SELF-DECORATAION INCLUDING COSMETICS
Ancient Egyptian culture used coloured pigments extensively for self-decoration and
cosmetics, with a wide variety of colours produced from various pigments
o A wider variety of pigments were available to Egyptians than in cave drawings of
earlier cultures because extraction techniques and refining technology had
improved
Black kohl (Sb2S3) was used as a black pigment for eyeliner and eye shadow
Cinnabar (HgS) was used as a bright red pigment for lipstick and rouge
Orpiment (As2S3) was used as a rich lemon-yellow pigment for body paint and eye shadow
Malachite (CuCO3Cu(OH)2) was used as a bright green pigment as body paint, particularly
around the eyes
See below for a more detailed list of cosmetics used in ancient Egyptian culture
PREPARATION OF THE DEAD FOR BURIAL
The cosmetic pigments above were also used by ancient Egyptian culture to prepare dead
bodies for burial and the afterlife
The internal organs of a body were often replaced with a fluid containing a mixture of resins
from coniferous trees, beeswax, and aromatic plant oils
Elemental gold, either as gold leaf or as grinded powder mixed with saliva and water, was
also used to decorate bodies and burial objects
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
CHEMICAL COMPOSITION OF COSMETICS
Pigment/dye
Colour
Composition
Malachite
Bright green
CuCO3.Cu(OH)2
Azurite
Blue
2CuCO3.Cu(OH)2
Gold
Gold
Ag
Cinnabar
Bright Red
HgS
Orpiment
Rich lemon-yellow
As2S3
Egyptian blue
Blue
CaO.Cu.4SiO2
Kohl
Black
Sb2S3
White lead
White
2PbCO3.Pb(OH)2
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Use
Eye paint
Make-up
Body paint
Lipstick
Eye shadow
Body paint
Eyeliner
Face paint
POTENTIAL HEALTH RISKS
Many of the cosmetics used in ancient Egyptian culture comprised various substances that
are toxic if ingested
o Cinnabar contains mercury, which can cause numbness and brain damage if in a
compound
o Orpiment contains arsenic, which can cause stomach and intestinal irritation and
blood vessel damage
o Malachite, azurite, and Egyptian blue contains copper, which can cause anaemia,
liver damage and kidney damage if ingested in high doses
o Kohl contains antimony, which can cause nausea, vomiting, and diarrhoea if ingested
o
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White lead contains lead => if lead builds up in the body, it can cause damage to the
nervous system, mental retardation, and death
Many of these toxic substances also can disfigure the skin after prolonged use
Thus the health risks with their use are significantly hazardous, due to their serious health
problems caused by their prolonged use.
Outline the processes used and the chemistry involved to prepare and attach
pigments to surfaces in a named example of medieval or earlier artwork
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Panel painting was a widespread form of painting in mediaeval society, which involved
painting on wooden panels
The processes involved in panel painting include preparation of the wood panel, preparation
of the pigment, and adhesion of pigment to panel
One example of a mediaeval panel painting is St John the Baptist with St John the Evangelist
and St James by Italian artist Nardo di Cione (1365)
PREPARATION OF WOOD PANEL
Wood and canvas are unsuitable to paint onto directly, as they are too rough and absorbent
The wood (poplar in the above painting) is prepared for painting through the application of
layers of grounding or priming
The grounding used was called gesso, which is a mixture of gypsum (CaSO4) or chalk (CaCO3)
and animal glue
Gesso was applied as a thick, warm liquid, but set to a brittle creamy white layer, which was
then scraped and rubbed smooth
PREPARATION OF PIGMENTS
Each colour used on in the artwork comes from different pigments, each of which has to be
extracted separately.
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The pink robes on St John the Baptist (middle) comes from the pigment crimson lake with
layers of white lead
o Crimson lake is produced by boiling dried female cochineal insects in sodium
carbonate (Na2CO3) solution to extract the carminic acid (C22H20O13), then
precipitating out the pigment onto a clear insoluble powder
The scarlet lining of the cloak and the cover of the book are vermilion (HgS)
o Vermilion is produced by heating mercury and sulfur together, until the mixture
vaporises and reacts to form mercury (II) sulfide, which condenses at the top of the
flask.
o The condensed vermilion is then grounded to produce the red pigment
The pigments were then mixed with egg yolk as a binder, then mixed with water as the
medium
ADHESION OF PIGMENT TO PANEL
The artist first would have drawn an outline of the image on the ground, and then would
have painted over the top of the outline
The paint was first applied, and then left so the medium could dry, leaving the pigment and
binder as a film above the gesso ground
Multiple layers of paint were applied, as artists in during the mediaeval era generally did not
mix colours, but instead painted in layers
The gold background was gilded onto the painting, rather than painted on
o The area to be gilded was coated with iron (III) oxide and egg white, which was then
polished and set hard
o The gold leaf was then added using glue such as egg white, then polished again
Describe a historical example to illustrate the relationship between the discovery of
new mineral deposits and the increasing range of pigments
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The discovery of new mineral deposits throughout history, as well as increasing extraction
and refinement technologies, led to an increased range of pigments available to painters
Examples of new discoveries leading to a wider range of pigments available include Naples
Yellow, chromium pigments, and cadmium pigments
NAPLES YELLOW
Chemical composition: Pb3(SbO4)2 [lead (II) antimoniate]
Colour: Yellow
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Discovery: Naples Yellow was originally discovered in mineral deposits on the slopes of
volcanic Mt Vesuvius, but was synthetically produced during the 17th century by the
prolonged roasting of oxides of lead and antimony.
CHROMIUM PIGMENTS
In 1770, an orange-coloured mineral was found in the Beresorf gold mine in Siberia
The mineral was analysed, and determined to be a compound of lead and chromium, which
had been recently discovered by French chemist L. N. Vauquelin
The discovery of vast deposits of the ore chromium, FeO.Cr2O3 in the U.S. in the 1820s led to
the manufacture of a range of chromium compounds, including many pigments, including…
o Chrome yellow (PbCrO4) [lead (II) chromate]
o Chrome red (PbCrO4.Pb(OH)2) [basic lead (II) chromate]
o Chrome green, formed by mixing chrome red with Prussian blue
CADMIUM PIGMENTS
Cadmium was first discovered by Stromeyer in 1817, but cadmium pigments were only
produced until the 1840s due to the scarcity of the metal
Pigments include…
o Cadmium yellow (CdS), prepared by reacting an acid solution of a cadmium salt
(either chloride or sulfide) with hydrogen sulfide gas or an alkali sulfide
o Cadmium red (CdS.CdSe), which was not prepared until 1910
Identify data, gather and process information from secondary sources to identify
and analyse the chemical composition of an identified range of pigments
Pigment
Gypsum
Cerussite
Stibnite
Galena
Graphite
Cinnabar
Malachite
Azurite
Orpiment
Turquoise
Red Ochre/ Haematite
Yellow ochre
Egyptian blue
Prussian blue
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Colour
White, grey
White, grey
Lead-grey, blackish
Lead-grey
Black
Red, brownish red
Bright green
Azure blue to dark blue
Lemon yellow – brownish yellow
Bluish green
Earthy to bright red
Yellow
Blue
Blue
Chemical composition
CaSO4.2H2O
PbCO3
Sb2S3
PbS
C
HgS
Cu2(CO3)(OH)2
Cu3(CO3)2(OH)2
As2S3
CuAl6(PO4)4(OH)8.4H2O
Fe2O3
Fe2O3.H2O
CaO.CuO.4SiO2
Fe7(CN)18.14H2O
An important note is the historic development of pigments, and its relationship to the
chemical composition of pigments
o Early pigments, which came from natural sources, consisted of mostly ochres, which
are oxides of iron
o The development of technology allowed mineral ores to be processed, typically
through roasting, which led to pigments containing Fe, Cu, Mn, and toxic Pb, Hg, and
Cr => all of these elements are transition metals
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The discovery of new elements and processes since the 16th century have allowed
synthetic inorganic pigments to be produced, such as Prussian Blue
Increased understanding in carbon chemistry during the 19th and 20th century have
allowed synthetic organic pigments to be produced
Analyse the relationship between the chemical composition of selected pigments
and the position of the metallic components(s) of each pigment in the Periodic Table
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Below is another table of a range of pigments used throughout history
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As can be seen in the above table, it can be seen that metals are present in each pigment,
and that certain colours are associated with each colour
o Iron produces colours in the red-yellow range
o Copper produces colours in the blue-green range
o Cobalt produces colours in the yellow-violet range
o Chromium produces colours in the red-yellow range
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Most of the metals present occupy the transition metal region of the Periodic Table, due to
the electronic configuration of transition metals (discussed in further detail below)
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
Describe the development of the Bohr model of the atom from the hydrogen spectra
and relate energy levels to electron shells
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If electricity is passed through a discharge tube containing hydrogen gas at very low
pressure, a purple glow can be observed
When this light is passed through a prism, a series of discrete spectral lines can be observed
As can be seen above, the spaces between the lines at the short wavelength (or high energy)
end of the spectrum as smaller than at the low energy end
In 1901, Max Plank developed the revolutionary quantum theory, which showed that light
was emitted and absorbed in discrete units called photons, and that the energy of a photon
was proportional to its frequency
Rutherford’s model of the atom visualised each atom consisting of a dense positivelycharged nucleus surrounded by orbiting electron
o Rutherford based his model under classical physics, which implied that an orbiting
electron would emit a continuous spectrum of electromagnetic radiation, so the
electrons would lose energy and spiral into the nucleus
o His model did not account for the discrete wavelengths of light absorbed in the
emission spectrum of hydrogen
Niels Bohr applied Plank’s quantisation of electromagnetic radiation to Rutherford’s model,
and proposed the Bohr model of the atom
Bohr proposed that the electrons in the hydrogen atom could only occupy certain energy
levels or stationary states around the nucleus. Each state was associated with a specific
circular orbit around the nucleus, and electrons did not emit radiation whilst in a stationary
state.
Bohr also stated that an electron can move from one energy state to another by emitting or
absorbing a photon with an energy equal to the energy difference between the two states
o This was Bohr’s explanation for the observed emission spectrum of hydrogen
In summary, the postulates of the Bohr model of the atom include:
o Electrons orbit around the nucleus in circular orbits under the influence of Coulomb
attraction to nucleus
o Orbital angular momentum is quantised, hence only certain orbits are possible
o Electrons in stable orbits do not radiate
o Electrons can change orbits by radiating (larger to smaller), absorbing radiation
(smaller to larger), or by collisions (either larger to smaller or smaller to larger)
Solve problems and use available evidence to discuss the merits and limitations of
the Bohr model of the atom
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See above for an explanation of the Bohr model of the atom
MERITS
The Bohr model introduced the newly developed quantum theory to the model of the atom,
which established many important features of the atom, such as quantised energy
differences between orbitals
o This explained the observation that excited atoms release generate line emission
spectra
His model successfully accounted for the observed emission spectrum , atomic radius, and
ionisation energy of hydrogen with reasonable precision
LIMITATIONS
Attempts to extend the Bohr model of hydrogen to other elements and ions disagreed with
the experimental line spectrum frequencies observed
Bohr could not explain why only a restricted number of energy levels existed, or why the
accelerating electrons did not lose energy
Bohr also couldn’t explain the further analysis of the emission spectrum of hydrogen, such as
the splitting of spectral lines when a magnetic field was applied
His model couldn’t explain the relative intensities of the observed line spectra
His model describes electrons in defined orbits around the nucleus, but this violates
Heisenberg’s uncertainty principal as an electron’s there is a limit to the certainty of an
electron’s position and momentum around the nucleus
ASSESSMENT
Whilst the Bohr model could not explain observations beyond the line spectra, his
application of quantum theory to the model of the atom allowed the further for further
development, thus it was an important development in the study of atoms
Explain what is meant by n, the principal quantum number
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As discussed above, the Bohr model of the atom introduced the quantisation of angular
momentum, which implied that electrons orbited around the nucleus in certain allowable
radii that was associated with a specific energy level
Each orbital radii or energy level can be represented by the principal quantum number, or n,
which indicates the energy level of an electron
The quantum number can only take an integer value (i.e. 1, 2, 3, 4…)
Energy levels with greater energy are assigned larger principal quantum number, so n=1
denotes the orbit closest to the radius, n=2 denotes the next energy level up, and so on
Identify that, as electrons return to lower energy levels, they emit quanta of energy
which humans may detect as a specific colour
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When an electron returns from an excited energy state to its ground state, it emits a quanta
of energy called a photon
The energy of the photon equals the energy difference, thus discrete wavelengths of light
are emitted by excited atoms or ions
If part of the emission spectrum of an atom or ion lies within the visible spectrum of
electromagnetic radiation, the human eye can detect these photons of light as specific
colours
The combination of wavelengths in the visible spectrum that are emitted can produce
specific colours, which can be observed in a flame test (see below)
o For example, the emission spectrum of the sodium ion in the visible spectrum
consists of two very close emission lines of wavelength that corresponds to yellow,
thus we see a yellow colour in a sodium ion flame test
Identify Na+, K+, Ca2+, Ba2+, Sr2+, and Cu2+ by their flame colour
Substance
Na+
K+
Ca2+
Ba2+
Sr2+
Cu2+
Flame colour
Yellow
Violet
Orange-red
Apple green
Red
Green-blue
Explain the flame colour in terms of electrons releasing energy as they move to a
lower energy level
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Recall that light does not consist of a continuous spectrum of light, but discrete ‘packets’ of
light of quantised or discrete wavelengths called photons
When a substance is sprayed into a flame, the cations absorb heat from the flame
If the quantity of heat absorbed equals the energy difference between two energy levels
within the ion, an electron can become excited and exist in a higher energy level
When the electron returns to a lower energy level, it emits a photon of light with energy
equal to the energy difference between the two energy levels, thus discrete wavelengths of
light are emitted
Electrons can either return directly to their ground state, or by a series of ‘jumps’ to lower
energy levels, thus a spectrum of discrete wavelengths are emitted => this is the emission
spectrum of an atom or ion
Each element and ion has a unique set of possible electron transitions, hence they emit a
unique emission spectrum
Thus a flame test can serve as an analytical tool for identifying the presence of certain
elements or ions
Explain why excited atoms only emit certain frequencies of radiation
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When an electron in an excited state drops from a higher energy level to a lower energy
level, it emits a photon of a discrete frequency that equals the energy difference between
the two energy levels according to E=hf
o Thus the greater the difference in orbit radius, the greater the energy of the photon
emitted
As the energy levels of electrons in the atom are quantised (i.e. there are a discrete number
of allowable atomic radii for electrons to occupy), there are only certain frequencies of
radiation that excited atoms can emit that correspond to the various energy difference
between energy levels
Distinguish between the terms spectral line, emission spectrum, absorption
spectrum and reflectance spectrum
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SPECTRAL LINE
If an atom or ion is excited, it will emit photons of certain frequencies as the excited
electrons drop to lower energy levels
If the emitted light is passed through a narrow slit, then a glass prism or diffraction grating,
the light disperses into its individual wavelengths
The individual wavelengths of light emitted by the exited atom or ion can then be seen as
spectral lines, which are separated by blank areas
EMISSION SPECTRUM
An emission spectrum is produced when electrons in atoms or ions have been excited to a
higher energy state, and emit photons of EMR characteristic of the element as the electron
returns to a lower energy state
An emission spectrum consists of emission lines separated by blank spaces, and can be
observed through a spectroscope (see above), which disperses light into its individual
wavelengths
Below is the emission spectrum of hydrogen and iron
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ABSORPTION SPECTRUM
If white light is passed through a cool gas, the atoms or ions within the gas absorb quantised
wavelengths of light that correspond to the energy difference between the energy states of
the atoms or ions, then re-emit the photons in all directions
This causes the light passing through the gas to have reduced intensity in certain
wavelengths, so if observed with a spectroscope they appear as black lines against a
continuous spectrum that correspond to the emission spectrum of elements present
REFLECTANCE SPECTRUM
Atoms and molecules absorb and reflect energy at wavelengths related to their atomic
structure, as the radiation excites surface molecules, leading to absorption => the spectrum
obtained by reflected light is called the reflectance spectrum
If the radiation reflected from a surface is analysed and compared to a non-absorbing or
white sample, the reflectance spectrum can be obtained, which has distinctive features that
can be used to identify minerals
The reflectance spectrum is the complement of the absorption spectrum, and represents the
visible colour of the object
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
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Below is the emission spectrum for sodium
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As can be seen, the spectrum consists of two closely-spaced emission lines in the yellow
region of the visible spectrum called a doublet
o The emission lines lie at 589nm and 589.6nm
The two lines represent the transition of electrons within the n=3 level, from the 3p orbital
to the 3s orbital
Each electron within an atom is spinning along its own axis, and electrons have quantised
angular momentum
According to Pauli’s exclusion principle, two electrons in the same orbital must have
different spins
Within the 3p orbital, one electron has angular momentum j=3/2, whilst the other has j=1/2
The interaction of the magnetic field produced by the spinning electrons and the internal
magnetic field cause the 3p orbital to split slightly, though the energy difference between
the two orbitals is very small
The transition from the 3p3/2 orbital to the 3s orbital causes the 589.0nm line, whilst the
transition from the 3p1/2 orbital to the 3s orbital causes the 589.6nm line
Thus the opposite spins of electrons within the 3p orbital cause the hyperfine splitting of
spectral lines
Outline the use of infra-red and ultra-violet light in the analysis and identification
of pigments and their chemical composition
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DETECTION EQUIPMENT FOR ABSORPTION ANALYSIS
The equipment used in the analysis of pigments using infra-red and ultra-violet light is called
a double-beam spectrometer
Below is a flow-chart of the operation of a double-beam spectrometer
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The basic features of the double-beam spectrometer are the same for both infra-red and
ultra-violet analysis, but the source of electromagnetic radiation and the detector are
different
o In infra-red analysis, the source is commonly a heated ceramic such as a silicon
carbide rod, and the detector is a thermocouple (used for measuring temperature)
o In ultra-violet analysis, the source is a tungsten lamp or a deuterium lamp, and the
detector is a photomultiplier (a device that can detect and amplify light from very
faint sources)
In the double-beam spectrometer, the radiation is split into two beams => one is passed
through the sample, and the other through a reference
o The reference is usually the solvent used to dissolve the sample
The detector measures the radiation passing through the sample and the solvent, and a
comparison of the intensity of the two beams allows the absorption of radiation to be
determined
The absorption spectrum is a graph of absorbance against frequency or wavelength, which
can be used to detect a contaminant in a sample, or detect the presence of a certain
pigment
Differences between infra-red and ultra-violet analysis include…
o For infra-red analysis, sample is held in a glass case made out of KBr (which does not
absorb IR) and compared to a sample that does not absorb IR such as NaCl, whilst in
the sample is in solution for UV analysis
o Infra-red analysis is based upon the molecular vibrations/stretching due to the
absorption of IR radiation, whilst UV analysis is based upon the excitation of
electrons to higher energy states
o Infra-red analysis is mainly used to identify organic molecules, though the molecules
must be polar, whilst UV analysis is mainly used to identify pigments containing
metal ions
INFRA-RED REFLECTANCE SPECTROSCOPY
Infra-red radiation covers wavelengths between 700nm and 1mm within the
electromagnetic spectrum
o Near infra-red radiation lies closest to 700nm, and is commonly called ‘cool’ infrared
o Far infra-red radiation lies closest to 1mm, and is commonly called ‘hot’ infra-red =>
far infra-red radiation is can be detected as heat
Infra-red reflectography is a non-destructive technique to detect an artist’s preliminary
drawings
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A lamp directs near infra-red radiation towards the surface of a painting, and the reflected
light is detected
Infra-red radiation penetrates the pigments on the surface layer of the painting, and reflects
from the white ground of the painting
This technique is useful for detecting the underdrawing where the artist has used graphite
pencil, charcoal, or black ink, as carbon strongly absorbs IR, so incident light does not reflect
back
Infra-red radiation is also absorbed by most copper-containing green pigments, so this
technique can also be used to detect copper-based pigments
ULTRA-VIOLET REFLECTANCE SPECTROSCOPY
UV reflectance spectroscopy is a non-destructive technique that compares the reflected
radiation from a pigment on the surface layer of a painting to the a material that does not
absorb UV radiation, such as SiO2
UV radiation causes the specific fluorescence in materials depending on the sample’s
chemical composition and age, so ultra-violet reflectance spectroscopy can reveal changes in
elemental composition on the surface, as well as identify certain pigments on the surface
o NOTE: Fluorescence is where a substance absorbs and re-emits radiation at different
frequencies
o NOTE: For analysing paintings, absorption spectroscopy is a destructive technique,
whilst reflectance spectroscopy is a non-destructive technique
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
RELATIONSHIP BETWEEN ABSORPTION AND REFLECTANCE SPECTRA
 Absorption spectra are obtained by passing light through a solution of the sample, and
detecting the percentage of radiation absorbed, also called absorbance. The wavelength
of radiation is then plotted against absorbance.
o Absorption is directly proportional to the concentration, so an absorption
spectrum can be used to determine the concentration of a substance in a given
sample
 Reflectance spectra are obtained by reflecting light off the surface of a sample, and
detecting the percentage of radiation reflected (rather than absorbed), which is called
reflectance. The wavelength of radiation is then plotted against reflectance.
o The reflectance spectra shows the visible colour of the pigment, as the colour
we see a pigment is produced by the reflected light, NOT the absorbed light (i.e.
colour that is seen is not absorbed)
 The intensity of absorbance/reflectance at a given wavelength indicates the
concentration of a given pigment, whilst the shape of the curve indicates the purity of
the pigment
 Reflectance spectra is the complement of absorption spectra, so peaks in the reflectance
spectrum of a given substance correspond to the trough in its absorption spectrum, and
vice versa
o
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This is because radiation that is absorbed is not reflected, so the greater the
absorbance, the smaller the reflectance
EFFECT ON ZINC OXIDE
Zinc oxide, also known as Chinese white, is very opaque white pigment, which has high
reflectance for all wavelengths of visible light
Ultra-violet light causes zinc oxide to fluoresce from white to pale yellow
Far infra-red radiation changes the colour of zinc oxide from white to yellow in the presence
of oxygen, although it returns to its natural white colour once cooled
EFFECT ON PIGMENTS CONTAINING COPPER
Ultra-violet light causes malachite to fluoresce from a green to a dirty mauve colour
Far Infra-red radiation permanently breaks down red copper (I) oxide, green malachite, and
verdigris to black copper (II) oxide, as all these substances decompose when heated
o Care must be taken when using far infra-red radiation to analyse artworks, as it may
cause these copper pigments to decompose
Gather, process and present information about a current analytical technology to:
•
describe the methodology involved
•
assess the importance of the technology in assisting identification of
elements in samples and in compounds, and
•
provide examples of the technology’s use
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Laser microspectral analysis is a current technique used to identify and analyse elements
and molecules in pigments
METHODOLOGY INVOLVED
A powerful pulsed laser is focussed on the surface, which vaporises a tiny amount of the
surface material
The vaporised sample is fed through a gap between two electrodes that sparks and excites
the vapour, which produces an emission spectrum as the excited particles return to their
ground state
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The radiation released is fed through a spectrometer, which can identify the released
emission spectrum
By comparing the spectrum obtained to the known spectrum of other elements and
molecules, the chemical composition of the pigments on the surface can be determined
EXAMPLES OF USE
Laser microspectral analysis is used to analyse the elemental composition of pigments used
in restoring paintings, as it is a highly sensitive technique that requires minimal sample
preparation and causes minimal damage to the painting
Analysis of elements present in pigments can also be used to identify the validity and
authenticity of an artwork
ASSESSMENT
Laser microspectral analysis is a very important analytical technique because…
o It is a highly sensitive technique, so it can identify trace elements in any kind of solid
and liquid samples
o It requires minimal sample preparation
o It can rapidly identify more than one element at a time, as opposed to AAS
o Whilst it is a destructive technique, only a negligible amount of sample is destroyed,
thus it is safe to use on paintings
Perform first-hand investigations to observe the flame colour of Na+, K+, Ca2+, Ba2+,
Sr2+, and Cu2+
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METHOD
A Bunsen burner flame was ignited and placed on a heat mat, and the aqueous nitrate salts
of sodium, potassium, calcium, barium, strontium, and copper were sprayed into the flame.
The resulting flame colour was observed and recorded
SAFETY:
o Substances used and fumes produced may be toxic if ingested => use low
concentrations of nitrate salts
o General safety precautions of using a Bunsen burner => wear safety goggles, spray
the flame from a safe distance
o Flame colours produced may be bright and cause eye damage => use low
concentrations, do not look at the flame for long periods of time
RESULTS
Salt
Cation
Flame colour
+
NaNO3
Na
Yellow
KNO3
K+
Violet
2+
CaNO3
Ca
Orange-red
BaNO3
Ba2+
Apple green
2+
SrNO3
Sr
Red
CuCl2
Cu2+
Blue-green
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ACCURACY/RELIABILITY/VALIDITY
The substances were repeatedly sprayed into the flame, which improved reliability
The results were checked against reliable sources such as textbooks and university websites,
which confirmed the validity of the results obtained
The spray bottles were cleaned before use, but they may have been impurities in the water
used to produce the aqueous solution, so this may have reduced the accuracy of the
resultant flame colours
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
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Each electron has four sets of quantum numbers that defines its energy state
The four quantum numbers are the principal quantum number (n), the azimuthal quantum
number (l), the magnetic quantum number (mL), and the electron-spin quantum number (mS)
o The principal quantum number indicates the relative size of the orbital, and takes
integer values from 1 to infinity => the higher the principal quantum number, the
greater the energy of the electron
o The azimuthal quantum number is related to the shape of the orbital within each
energy level, and takes integer values from 0 to n-1 for each value of n (i.e. n=1 can
only contain l=0, n=2 contains l=0 and 1 etc.)
o The magnetic quantum number describes the orientation of the orbital in the 3D
space about the nucleus, and takes integer values from –l to l for each value of l
o The electron-spin quantum number indicates the direction of an electron’s spin, and
can take one of two possible values: -1/2 or ½ => whilst the other three quantum
numbers describe an electron’s orbital, the electron-spin quantum number
describes a property of the electron itself
Pauli’s exclusion principle states that no two electrons in the same atom may have the same
set of four quantum numbers
o Another way of stating Pauli’s exclusion principle is that an orbital can hold a
maximum of two electrons that must have opposite spins (i.e. different energy
states)
o This is used to explain the occurrence of doublets due to the interaction of the
magnetic field produced from a spinning electron and the internal magnetic field,
such as in sodium (see above)
Identify that each orbital can contain only two electrons
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An orbital is defined by the first three quantum numbers as described above (i.e. each
orbital has a unique size, shape, and orientation within an atom)
An electron within an orbital can have two possible spins, which are defined by the electronspin quantum number
Thus a consequence of Pauli’s exclusion principle is that each orbital can contain a maximum
of two electrons, which must have opposite spins
Define the term sub-shell
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The principal quantum number defines the energy level of an electron (i.e. the greater the
principal quantum number, the further away an electron is from the nucleus)
Within each energy level there are sub-shells, each with slightly different energies
Sub-shells result from electron-nucleus attractions and electron-electron repulsions
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Sub-shells are defined by the second quantum number defined above (the azimuthal
quantum number), so each sub-shell has an unique shape within each energy level
Sub-shells are also assigned letters, which are listed below in order of increasing energy
Sub-shell
Number of orbitals
s
p
d
f
1
3
5
7
Maximum number of
electrons
2
6
10
14
Shape
Spherical
Dumb-bell
Complex
Complex
Process information from secondary sources to use Hund’s rule to predict the
electron configuration of an element according to its position in the Periodic Table
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ORBITAL NOTATION
The two methods of orbital notation are electron configurations or orbital diagrams
Electron configurations
Writing the electron configuration of an element involves writing the sub-shells present in
an element, and the number of electrons within each subshell
The electron configuration of sodium is show below
The letter (s, p, d, or f) indicates the orbital
The prefix is the principal quantum number for that orbital
The superscript suffix indicates the number of electrons in each orbital
The order of the subshell indicates increasing energy (i.e. 1s has the lowest energy, 2s the
second-lowest energy etc.)
Condensed electron configurations involves listing the number the electrons in each energy
shell (i.e. the Na would have a condensed electron configuration of 2, 8, 1)
Another condensed form of electron configurations is to write a noble gas in square bracket
to represent its electron configuration, then write the relevant electron configuration of the
element. For example, the electron configuration of iron can be written as [Ar] 3d6 4s2
Orbital diagrams
Orbital diagrams give an indication of the various orbitals within an element, and the spin of
each electron
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Each orbital is represented by a box, and each electron by an arrow => an upwards arrow
indicates an electron spinning in one direction, a downward arrows indicates an electron
spinning in the other direction
HUND’S RULE
Hund’s rule states that orbitals of equal energy each acquire one electron before any orbital
acquires two electrons, and all electrons in singly occupied orbitals have the same spin
Below are the orbital diagrams for various elements
Notice that the carbon atom has two half-filled 2p orbitals, and that the first 2p orbital only
gains an electron pair once all the other 2p orbitals have been filled
This rule becomes more complicated when considering d orbitals, which is discussed below
Outline the order of filling of sub-shells
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To determine the order of filling sub-shells, apply the following principles
o Electrons are placed into orbitals starting with the lowest-energy orbital first
o A maximum of two electrons are placed into each orbital, but where more than one
orbital with the same energy is available, electrons are placed into each orbital
before pairing electrons up
The diagram below shows the energy level diagram for an atom with multiple electrons
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As can be seen, the 4s orbital has a lower energy than the 3d orbital, so the 4s orbital is filled
before the 3d orbital
For example, the electron configurations of potassium are given below
The 4s orbital is filled before the 3d orbital, as the 4s orbital has a lower energy than the 3d
orbital
The diagram below provides a useful method for working out the order of filling orbitals=>
follow each arrow from bottom to top for the order of filing the s, p, d, and f orbitals
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
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According to Heisenberg’s uncertainty principle, the position of an electron cannot be
determined with certainty simultaneously with its momentum
Hence when considering the position of electrons, we must consider the probability of
finding an electron in a particular state
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An electron has a greater probability of being in the lowest state of energy possible, thus
electrons in their ground-state electron configurations occupy the lowest energy shells, subshells, and orbitals available to them
When an electron becomes excited (e.g. it absorbs a photon of EMR), it gains the energy
required to occupy orbitals of higher energy states, thus an excited electron can jump to
higher energy levels
o Remember that the energy absorbed by an electron when transitioning to a higher
energy level is equal to the difference between the two energy levels
Thus the electron configuration of an excited atom or ion is different to when it is in its
ground state, as it can occupy higher energy levels
NOTE: When drawing the electron configuration of excited elements or ions, the electron
with the greatest energy is the one that jumps to a higher energy state, but REMEMBER to
write/draw any unfilled orbitals (e.g. for an excited Na atom, write/draw the 3s orbital, but
leave it unfilled)
EXAMPLE
Ca atom: 1s2 2s2 2p6 3s2 3p6 4s2
Ca+ ion: 1s2 2s2 2p6 3s2 3p6 4s1
Ca atom (excited): 1s2 2s2 2p6 3s2 3p6 4s1 4p1
NOTE: Once the electrons are actually in their orbitals, the energy order can change, namely
the order of 4s and 3d reverses. So the electron configuration of iron is 1s2 2s2 2p6 3s2 3p6
3d6 4s2, and the 4s orbital behaves as the valence orbital as it has the highest energy. So
when writing the electron configuration of elements with 3d and 4s electrons, write 3d
before 4s
Explain the relationship between the elements with outermost electrons assigned to
s, p, d and f blocks and the organisation of the Periodic Table
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Recall that the Periodic Table was organised to group elements with similar chemical
properties together, and to reflect the periodic trends of many chemical properties
Elements with similar electron configurations in the outermost shells display similar
chemical properties. Further, when elements are listed in increasing atomic number, similar
outer shell or energy level electron configurations are observed to recur in periodic intervals
Thus elements with similar outer shell configurations occur in the same vertical group, and
elements with the same subshell as their outer shell exist in blocks on the periodic table
o For example, all group 1 elements have an s1 valence shell electron configuration,
whilst group 13 elements all have a s2p1 valence shell electron configuration
Consideration of outer shell electron configurations enables the Periodic Table to be divided
into four major blocks that reflect the filling of valence sub-shells, as shown below
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The above diagram can also be used to remember the order of filling sub-shells
Explain the relationship between the number of electrons in the outer shell of an
element and its electronegativity
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Electronegativity is a measure of the ability of an atom in a molecule to attract electrons to
itself in a chemical bond
Electronegativity is measured on a relative scale, with the most electronegative element,
fluorine, assigned a value of 4.0, and the rest of the elements compared to this
Electronegativity values for the main groups of elements are shown below:
As can be seen, electronegativity increases left to right across a period, and decreases down
a group
As the number of electrons in the valence shell increases, so too does its electronegativity
for the following reasons:
o There is an increase in nuclear charge across periods, so the force of attraction
between the valence electrons and the nucleus is greater
o
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Atomic radius decreases across a period due to increased nuclear charge, which also
increases the force of attraction between the valence electrons and the nucleus
Electronegativity decreases down a group because of increased atomic radius, and the
increased shielding of nuclear charge by lower energy electrons, thus decreasing the
attractive force of the nucleus on valence electrons
The noble gases have no recorded electronegativity values due to their tendency not to form
molecules with other atoms
Process information from secondary sources to analyse information about the
relationship between ionisation energies and the orbitals of electrons
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Ionisation energy is the energy required to remove one mole of the most loosely-held
electrons from one mole of gaseous atoms or ions
o Units are typically kilojoules per mole (kJmol-1)
o The energy required to remove the most loosely-held electron from a neutral atom
is called the first ionisation energy
The first ionisation energies of the first twenty elements is shown below:
Generally, first ionisation energies increase across a period for the following reasons:
o The nuclear charge increases as atomic number increases, so the attractive force
between the most loosely-held electron and the nucleus increases, thus more
energy is required to remove overcome the stronger bond and remove this electron
o Atomic radius decreases across a period due to the greater nuclear charge, thus the
most loosely-held electron is more tightly held by the nucleus (attractive force
between an electron and the nucleus is inversely proportional to distance)
In addition, first ionisation energies decrease down a group for the following reasons:
o The atomic radius increases down a group, so the attractive force between the most
loosely-held electron and the nucleus decreases
o Nuclear charge increases down a group, so the nuclear force on the valence
electrons increases
o The inner electrons effectively shield the nuclear charge from valence electrons, so
the outer electrons experience a decreased nuclear charge, which is called the
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effective nuclear charge. Elements further down a group have a greater number of
inner electrons shielding the nuclear charge on the valence electrons, so the
attractive force on most loosely-held electron decreases.
Whilst the explanations above describe the general trends in first ionisation energy across a
period, there are a few significant anomalies that occur due to the varying penetrating
power of sub-shells and the electron configurations of various elements
Penetration power
Electrons in the s orbital have a greater probability of being found closer to the nucleus than
p electrons.
o The order of penetrating power is s > p > d > f
If an electron is closer to the nucleus, it will be more firmly held than electrons in less
penetrating orbits. Thus electrons in less penetrating orbits have greater energy, and require
less energy to be removed from the atom or ion.
In addition, electrons in a less penetrating orbit are shielded from the attractive nuclear
charge by electrons in more penetrating orbits
Thus it takes less energy to remove an electron from a less penetrating orbit
For example, consider the first ionisation energies of beryllium and boron
Beryllium has an electron configuration of 1s2 2s2, whilst boron has an electron configuration
of 1s2 2s2 2p1
The 2p orbit has greater energy than the 2s orbit for the above reasons, thus the ionisation
energy for boron is slightly less than that of beryllium, despite the increased nuclear charge
A similar effect occurs for aluminium
Electron configurations
Recall that according to Pauli’s exclusion principle, each orbit can only contain a maximum of
two electrons, and according to Hund’s rule every orbital in a sub-shell is singly occupied
with one electron before any one orbital is filled with an electron pair
When two electrons occupy the same orbital, electron repulsions within the orbital raise the
energy of its electrons
Thus less energy is required to remove an electron from an orbital with an electron pair than
an orbital containing a single electron
For example, consider the first ionisation energies of nitrogen and oxygen
Nitrogen has an electron configuration of 1s2 2s2 2p3, whilst oxygen has an electron
configuration of 1s2 2p2 2p4 => the three 2p electrons occupy separate orbitals, but the
additional electron in oxygen is paired up in the same orbital as one of these electrons
Thus less energy is required to remove the most loosely-held electron in oxygen than in
nitrogen despite the increased nuclear charge, which is why oxygen has a lower first
ionisation energy than nitrogen
A similar effect occurs for sulfur
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
elements
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The energies required to remove subsequent electrons from a gaseous atom or ion are
called successive ionisation energies
o The energy required to remove the first most loosely-held electron is called first
ionisation energy, the energy required to remove the second electron is called
second ionisation energy etc.
For a single atom, the ionisation energy increases for successive ionisation energies for its
ions, as the number of electrons decreases whilst the nuclear charge remains the same
Another factor that contributes to a rise in successive ionisation energy is the removal of an
electron from an energy level closer to the nucleus
This leads to a substantial increase in ionisation energy, as the most loosely-held electron is
much closer to the nucleus than the previously-removed electron
Consider the successive ionisation energies listed in the table below
Note the proportionally large jumps in ionisation energy between the first and second for
Na, the second and third for Mg, the third and fourth for Al, and the fourth and fifth for Si
This indicates that the electron removed has come from a lower energy level than the
previously-removed electron
Thus we can conclude that that Na has one valence electron, Mg has two valence electrons,
Al has 3 valence electrons, and Si has four valence electrons
These trends in successive ionisation energies can be used to predict the number of
electrons in the valence shell, which can then be used to identify the sub-shells occupied by
these elements
Combined with the trends described in the dot point above, the electron configurations of
elements can be predicted from the relevant ionisation energies of elements
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
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Transition metals occupy the d-block of the Periodic Table, which corresponds to the filling
of the d orbital
Define the term transition element
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A transition element is one which forms one or more stable ions which have incompletely
filled d orbitals
On the basis of this definition, scandium and zinc are not transition elements, even though
they are members of the d block
o Sc3+ is the only stable ion of scandium, which has no d electrons => Sc3+ has an
electron configuration of 1s2 2s2 2p6 3s2 3p6 (it loses its two 4s electrons and its
single d electron during ionisation)
o Zn2+ is the only stable ion of scandium, which has a full d orbital => Zn2+ has an
electron configuration of 1s2 2s2 2p6 3s2 3p6 3d10 (it loses its two 4s electrons during
ionisation)
The examples above show that the terms d-block element and transition element are not
interchangeable
Process and present information from secondary sources by writing electron
configurations of the first transition series in terms of subshells
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The table below lists the electron configurations of the first transition series

Notice that the electron configurations of chromium and copper do not follow the general
pattern, as the 4s orbital is not completely filled in both cases
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An explanation for the anomaly is that a more stable electron configuration is achieved
when all the 3d orbitals are either half-filled with one electron each (chromium) or
completely filled with two electrons each (copper)
o This is not the case with tungsten (W) however, which has the same number of
outer electrons as chromium, but has a different outer structure of 5d4 6s2 => the
proper explanation for the electron configurations of chromium and copper are
beyond the scope of the HSC course
Explain why transition metals may have more than one oxidation state
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Recall from the first module that oxidation state is a degree of oxidation of an atom in a
chemical compound => see module 1 for rules for calculating oxidation states
One of the characteristic features of transition metals is the occurrence of multiple oxidation
states
Multiple oxidation states occur because the 3d and 4s sub-shells in transition metals have
similar energies, thus transition metals can lose electrons from both these sub-shells in the
formation of chemical bonds and compounds
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The 4s electrons are ALWAYS lost first to produce an oxidation state of +2, as the 4s orbital is
higher in energy than the 3d orbital as soon as electrons occupy the 3d orbitals
Oxidation states above +2 result from the additional loss of 3d electrons
The table below shows the possible oxidation states for the first transition series, and their
relative occurrences
Solve problems and process information from secondary sources to write halfequations and account for the changes in oxidation state
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NOTES ON DETERMINING OXIDATION/REDUCTION HALF-EQUATIONS
Balance the atoms on each side of the equation
o Oxygen is balance with H2O
o Hydrogen atoms are balanced with H+
Balance the charges with electrons (i.e. add electrons to either side)
Make sure to write the relevant states
EXAMPLE
Consider the reduction of MnO4- to Mn2+
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To balance the four oxygen atoms on the reactants side, add four water molecules to the
products side
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To balance the eight hydrogen atoms on the products side, add eight hydrogen ions to the
reactants side
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The charge on the reactants side is +7, whilst the charge on the products side is +2, so add
five electrons to the reactants side to balance the charges
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Finally, write the relevant states
OTHER EXAMPLES OF CHANGING OXIDATION STATE
Change in oxidation number Half-equation
Cr2O72-(aq)+ 14H+(aq) + 6eCr(+6) to Cr(+3)
Fe3+(aq) + eFe2+(aq)
Fe(+3) to Fe(+2)
Cu2+(aq) + eCu1+(aq)
Cu(+2) to Cu(+1)
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2Cr3+(aq) + 7H2O(l)
In the case of reduction for the above equations, use the single arrow inside of the double
arrow
The oxidation states in the reduction equations above have decreased due to a gain of
electrons in the d sub-shell (see the dot-point above for a more detailed explanation of dorbitals and oxidation states)
Account for colour changes in transition metal ions in terms of changing oxidation
states
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Another characteristic of transition metals is that their compounds are often coloured
The d orbitals in transition metal ions have slightly different energies, and are incompletely
filled
The small energy differences between the d orbitals are similar to the energies of visible
light, thus the electrons in transition metal ions can jump to slightly higher-energy d orbitals
by absorbing photons of visible light
The excited electrons usually returns to its ground state via a different set of energy
transitions that do not emit photons of visible light, so the overall effect is that some
frequencies of visible light are absorbed
The absorption of certain components of white light causes the compound to appear
coloured, with the colour being the complement of the wavelengths absorbed
o For example, if red wavelengths are absorbed, the compound will appear blue-green
Each transition metal ion in a different oxidation state has a different arrangement of filled
and unfilled 3d orbitals, so the energy differences between orbitals will be different for
changing oxidation states
In addition, different oxidation states for transition metal ions result from different numbers
of electrons, which determines the d-orbitals that are filled or empty, and which d-electrons
can gain energy and occupy them
Thus the changing oxidation states of a transition metal ion can result in the production of
different coloured compounds, as the photons of visible light absorbed changes
NOTE: Cu+ is not coloured, as it has a full d sub-shell. Cu2+ is coloured however, as it has 9
electrons in d orbitals
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
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The amount of energy required to change the oxidation state of a transition metal, since the
4s and 3d orbitals are of similar energy levels
Thus transition metals are easily oxidised and reduced, and are useful as oxidants or
oxidising agents
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o An oxidising agent is one that causes another substance to be oxidised
The strength of an oxidising agent depends on the ease with which the compound will
accept electrons (and hence be reduced)
Transition metals with high oxidation states have lost a large number of electrons, thus have
a smaller radius as the electrons are removed
This smaller radius gives a greater attraction for electrons, thus species containing transition
metals in a high oxidation state are strong oxidising agents
For example, the reduction half-equations for Cr2O72- and MnO4- are given below
Chromium in chromate (CrO42-) dichromate (Cr2O72-) has an oxidation state of +6, and
manganese in permanganate (MnO4-) has an oxidation state of +7
o Permanganate works best as an oxidising agent in acidified solution, so add 1M of
sulfuric acid when using permanganate as an oxidising agent
The high EΘ indicates that these compounds are strong oxidising agents, which is because of
the high oxidation states of chromium and manganese in each compound respectively
NOTE: For substances to be reduced, they require reducing agents such as Fl- ions. For
substances to be oxidised, they require oxidising agents such as fluorine gas => refer to the
table of standard potentials if asked to identified a required oxidising/reducing agent for
transition metal oxidations/reductions
Perform a first-hand investigation to observe the colour changes of a named
transition element as it changes in oxidation state
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METHOD
In a 250mL conical flask, 3g of ammonium vanadate was dissolved in 100mL of 1M NaOH,
then acidified by adding 75mL of 1M sulfuric acid. Approximately 20mL of the initially yellow
solution was poured into a large test tube. 6-8 granules of zinc were then added, and the
flask was stoppered with a rubber bung. The flask was the gently swirled, and each time a
colour change occurred, 20mL of the solution was poured into a test tube, until no more
colour change occurred.
SAFETY:
Ammonium vanadate can cause irritation to the skin and eyes, and can cause damage to
upper respiratory tract => clean up spills immediately, wear safety glasses, transfer
ammonium vanadate with a dropper, use small volumes of ammonium vanadate to
minimise exposure
Acids are corrosive, bases are caustic => use low concentrations, clean up spills on skin
immediately with running water
RESULTS
Test tube
Vanadium ion
Oxidation state of
Colour
vanadium
1
VO3+5
Yellow
2+
2
VO
+4
Blue
3
V3+
+3
Green
4
V2+
+2
Violet
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The electron configurations for the various oxidation states of vanadium are shown below:
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The reduction half-equations for each step of the reaction are as follows:
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The oxidation half-equation for each step of the reaction was the same:
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The above redox reactions account for the changes in oxidation state for vanadium, as
electrons in the d orbital are gained
ACCURACY/RELIABILITY/VALIDITY
As this was a qualitative experiment, accuracy is not important
The results of others in the class were the same, and corroborated with reliable sources such
as scientific journals, thus the results were reliable
All equipment was cleaned beforehand, so there was minimal chance for contamination
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The experiment was conducted under normal conditions (i.e. not at extreme temperatures
or pressures), so the change in colour can be attributed to the changing oxidation state of
vanadium
Choose equipment, perform a firsthand investigation to demonstrate and gather
first-hand information about the oxidising strength of KMnO4
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METHOD
10mL of 0.01M potassium permanganate was mixed with 10mL of 1M sulfuric acid. Five
drops of the mixture was added to test tubes containing potassium iodide, potassium
bromide, potassium chloride, iron ammonium sulfate solution, zinc powder, magnesium
strip, copper turnings, tin, and iron. Any colour changes were observed.
Potassium permanganate needs to be dilute, otherwise the colour will be too dark to
observe any colour changes. Similarly, the other tested substances need to be in excess in
the reactions for a complete colour change to occur
To compare oxidising strengths, repeat the experiment with potassium dichromate
(K2Cr2O7), and compare the ease of reduction
SAFETY:
Potassium permanganate is a skin irritant as it is a strong oxidant, and also stains skin and
clothing => use dilute concentrations, use a pipette dropper to transfer solutions
General risks of using acids (corrosive) and other chemicals (may be toxic)
RESULTS
SUBSTANCE
OBSERVATION
KI/KBr/KCl
Solution turned from purple to colourless with slight pink/brown
colour
Iron ammonium sulfate,
Solution turned from purple to a red/orange solution
Zn, Mg, Cu, Sn, Fe
For the KI/KBr/KCl reactions, the permanganate ions are reduced to manganese(II) ions,
which are colourless in dilute solutions, but pink to brown in colour in concentrated solution
For the metal reactions, the permanganate ions are reduced to manganese (IV) dioxide,
which is a red/orange colour in dilute solutions, but brown when more concentrated
ACCURACY/RELIABILITY/VALIDITY
Dilute concentrations of potassium permanganate were used so the colour change would be
pronounced. Similarly, the other reagents were more concentrated and greater volumes
used so that they would be in excess, thus the colour change was complete
The results were compared to others in the class, who found similar results, thus the
experiment was reliable
The beaker containing acidified potassium permanganate was used as a control, with no
reagents tested in it (no colour change was observed).
The validity of the experiment could have been improved by comparing the oxidising
strength of potassium permanganate to potassium dichromate, whilst controlling all other
variables.
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
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Recall that water is a polar molecule, with the negative dipole forming on the oxygen atom,
and positive dipoles forming on the hydrogen atoms
When substance is ionised or dissolved in water to form ions, the water molecules surround
the ions ion-dipole forces => this is called a hydrated ion in solution
For many cations, the ion is surrounded by a fixed number of tightly-bound water molecules
(i.e. a specific number of water molecules is associated with each formula unit)
For most metal cations, especially transition metal cations, the water molecules bond via
coordinate covalent bonds, where a lone pair of electrons from the oxygen atom is donated
to fill the vacant orbitals in the cation
For example, colourless Cu2+ forms blue Cu(H2O)42+ in solution, so four water molecules
surround each Cu ion
o The above formula can also be written as Cu(OH2)42+, as the atom or atoms of a
ligand that coordinate to a metal atom are listed first (see below for a more detailed
description of ligands)
NOTE: The charge on the ion does not change despite being hydrated, as water is a neutral
molecule
Describe hydrated ions as examples of a coordination complex or a complex ion and
identify examples
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A complex ion, also known as a coordination complex, is a class of ions that consist of a
central metal cation surrounded by molecules or anions called ligands
o Hydrated ions are an example of complex ions, with the water molecules acting as
ligands
Compounds that contain complex ions are called coordination compounds
The following are examples of complex ions:
o Cobalt (II) chloride hexahydrate (CoCl2.6H2O)
o Magnesium sulfate heptahydrate (MgSO4.7H2O)
o Calcium sulfate dehydrate (CaSO4.2H2O)
Describe molecules or ions attached to a metal ion in a complex ion as ligands
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The molecules or anions attached to a metal ion in a complex ion are called ligands
The number of ligands bonded to the central metal atom is the coordination number
Examples of ligands include anions, such as chloride (Cl-) and cyanide (CN-) (the carbon atom
in cyanide bonds), and polar molecules, such as water (H2O) and ammonium (NH3)
Explain that ligands have at least one atom with a lone pair of electrons
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Ligands have at least one atom with a lone pair of electrons that can be used to form a
coordinate covalent bond to a metal ion
Ligands with one atom with a lone pair of electrons are called monodentate ligands
o Examples include Cl-, NH3, and F-
Ligands with multiple atoms with lone pairs of electrons are called polydentate ligands or
chelated ligands (see dot-point below for more information and examples)
Identify examples of chelated ligands
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Chelated ligands are ligands that bond through electron pairs on more than one donor atom
to the central metal cation in a complex ion.
Chelated ligands are able to bond with the central atom in multiple locations, and tend to
form rings in complex ions
Examples include the oxalate ion (-OOC-COO-, or C2O42-) and the triphosphate ion [(P3O10)5-]
Use available evidence and process information from secondary sources to draw or
model Lewis structures and analyse this information to indicate the bonding in
selected complex ions involving the first transition series
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LEWIS STRUCTURES
Lewis structures provide a useful model for visualising the structure of complex ions, which
allows for a greater understanding of their chemistry
To draw Lewis structures for complex ions, follow the following steps
o Write the chemical symbol for the central transition metal in the centre, and draw
ligands surrounding it
o For the ligands, draw the valence shells using dots ONLY => ensure the right number
of valence electrons for each element (e.g. hydrogen should only have two for a full
valence shell, whilst chlorine should have eight for a full shell)
o Draw a single, unbroken line between the lone electron pair and the central
transition metal ion for each coordinate covalent bond
o If the ion has a charge, draw square brackets around the structure, and write the
charge at the top right outside the brackets
Below are some examples of Lewis structures of complex ions involving the first transition
series. Note that the arrows should be unbroken straight lines, and that all valence electrons
should be represented by dots, not by crosses
Memorise some of the complex ions listed above and below so they can be easily
reproduced in an exam
NOTE: If asked to identify free electron pairs in a given ligand, make sure to double check all
bonding locations, not just those marked with a charge. One common bonding site is on the
lone pair of electrons in nitrogen if nitrogen only has three other bonds.
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
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Chromium displays a wide range of colours in its complexes
Formula
[Cr(H2O)6]2+
[Cr(H2O)6]3+
[Cr(H2O)5Cl]2+
[Cr(OH)4]
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Colour
Blue
Violet
Green
Deep green
The different colours result from the effect of both the oxidation state and the surrounding
ligand groups on the energies of the d orbitals in the chromium ion
As a result of the small variations in the energy separation of the d orbitals, the frequency of
the photons of visible light absorbed in electron transitions are affected, thus different
colours are produced
Complexes that have no d electrons or have a full d sub-shell cannot absorb visible light,
because none of the energy differences in such complexes equal the energy of photons of
visible light
o These complexes are colourless, such as [Zn(H2O)6]2+
Discuss the importance of models in developing an understanding of the nature of
ligands and chelated ligands, using specific examples
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In order to visualise how ligands and chelated ligands bond in complex ions, models such as
Lewis structures, molecular modelling kits, and computer images are used
Such models are useful for explaining the structure, bonding, and hence the chemical
interactions that occur in complex ions
Models are also useful for indicating the geometry of chelated ligands, and visualising
difficult concepts
Such models are limited, however, as they do not show the following important features:
o Electrons and their transitions within orbitals
o The energy changes involved
o The rate at which chemical interactions occur
o What initiates electron, ionic, and molecular interactions
o The mobility or flexibility of the bonds involved
Thus whilst models are useful for enhancing our understanding of ligands and chelated
ligands, they should only be used as a guide for explaining more complex chemical ideas
Below is a sample response to an HSC question on the contribution of Lewis models in the
development of our understanding of the structure of complex ions formed by transition
metals, using examples