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1. From earliest times, people have used colour to decorate themselves and their
Identify the sources of the pigments used in early history as readily available minerals
Pigments available to early artists were derived from natural resources. They are different to dyes, which are
usually soluble in water, which stain and run on surfaces.
- White: produced in white clay (kaolin), hydrated aluminium silicate (Al2O3.2SiO2.2H2O); chalk,
calcium carbonate; gypsum, calcium dihydrate (CaSO4.2H2O)
- Ochre (natural earth of silica and clay):
o Yellow – various hydrated iron oxides, main one being mineral goethite (Fe2O3.H2O)
o Red – contains anhydrous iron (III) oxide (Fe2O3)
o Brown – nearly pure limonite (FeO(OH))
- Black – naturally occurring manganese (IV) oxide or charcoal, which is mainly graphite, resulted from
the burnt wood, animal bones or the soot from lamps.
Explain why pigments used needed to be insoluble in most substances
Pigments are inorganic (obtained from rock) or organic (obtained from living material) substances of
small particles that is insoluble in the applied medium and is used on account of its colouring,
protective or magnetic properties.
Pigments need to not stain and run on surfaces. When exposed to moisture or water, pigments must
not be easily dissolved so that it remains adhered to surfaces. Also, cosmetics made from insoluble
minerals ensure that it doesn’t dissolve in perspiration. Pigments should also be opaque.
Outline the early uses of pigments for:
cave drawings
Rock paintings have been produced since the Stone Age. Oldest cave paintings confirmed by
radioactive dating are in Chauvet Cave, France. Pigments used include red iron oxide, soot, black
manganese oxide and white porcelain clay. These were mixed with cave water, which had high
calcium content, ensuring good adhesion and long durability.
Aborigines used pigments to produce schematic artworks on rocks located near ceremonial sites or
corroboree grounds. The pigments were obtained from red and yellow ochre, charcoal, chalk and
clay, and mixed with binders that included blood, raw egg, animal fat, sap, honey or wax. Paint
applied using fingers, brush or chewed and frayed end of a twig.
self-decoration including cosmetics
Aboriginal uses of pigments include smearing clay and ochre in geometrical patterns on themselves
for ceremonial purposes. Charcoal was sometimes rubbed on foreheads.
Ancient Egyptians used eye makeup, lipstick and rouge for cosmetic purposes. Kohl was used to
darken eyebrows, for eyelashes and sometimes to darken hair. It has also been thought that black
eye paint (constitutes Laurionite – PbOHCl and phosgenite – Pb2Cl2CO3) were used as bactericide
protection against eye disease in the hot dry Egyptian climate. Laurionite and phosgenite do not
occur naturally so they had to synthesise them.
Ancient Greek women used cerusse (lead carbonate), a white pigment to make their face pale.
Ancient Romans also used lead carbonate and chalk for face paint and rouge made out of lead oxide.
preparation of the dead for burial
Pigment has been used in burial sites all around the world, including Qafzeh Cave, funeral contexts in
Greece and Spain, Lake Mungo in NSW, and Krems in Austria.
Red pigment was obtained from Red Ochre. Red is associated with blood and has been thought to
symbolise life’s meaning and life over death. Red ochre is a suitable material as a pigment as it is soft,
easy to ground in powder and is insoluble.
Red pigment was also obtained from cinnabar. Cinnabar may have had a role as a preservative due
to its toxicity.
Ancient Egyptian tombs contained corpses painted with a representation of the person and that the
walls of the sarcophaguses were painted depicting the history of life. Tutankhamen’s tomb
contained samples of gypsum, orpiment, haematite and malachite.
Outline the processes used and the chemistry involved to prepare and attach pigments to surfaces in a
named example of medieval or earlier outwork
“Madonna and the Child with Saints Jerome, John the Baptist, Bernadino and Bartholomew,” by
Sano di Petro is 600 years old. Originates from Italy.
The surface is made of wooden panels – oak, pine or poplar with canvas (linen) glued with animal
glue. Gesso is coated. Grosso is the coarser and thicker layer that it painted, with sottiel painted on
top as a smooth top surface. Charcoal is used for underdrawing.
Pigments used include cinnabar (red), azurite (blue), gypsum (white), orpiment (yellow) and
malachite (green). Egg tempera used as binders (commonly used in late medieval panel painting)
Decoration produced by mordant gilding. The area to be gilded is coated with mixture of iron (III)
oxide and egg white, which is then polished after it had set hard. The gold leaf applied using glue
such as egg white and burnished with a tool.
Explain that colour can be obtained through pigments spread on a surface layer (e.g. paint) or mixed with
the bulk of material (e.g. glass colours)
Pigments added onto a surface layer
- Paintings contain layers including wood, canvas, ground layers, drawing, paint layers and varnish.
Panels prepared for painting by mixing gypsum or chalk with animal glue. It was applied as a thick
warm liquid and set to a creamy white layer which was rubbed smooth.
- Paint layers included pigments mixed with a bind medium or a drying oil. For example, the blue
colour of natural ultramarine pigment is maintained within aqueous solutions and egg tempera
which binds the pigment.
Mixing with bulk of material
- Glass was made from melted sand (SiO2). This was obtained by adding flux (mixture of sand, soda ash,
lime, potash and lead oxide to lower melting point) and heating sand at 800 degrees Celcius.
- Pigments were added to the molten glass, and the pigment disperses and traps in glass as it cools
and hardens. Details were painted onto glass with one dark pigment of copper or iron oxide mixed
with soft powered glass, then made permanent by firing in kiln. An example of this is stained glass.
Describe paints as consisting of:
the pigment
Pigments must be chemically inert and should be unaffected by light, heat, acids, moisture or
pollution. They impart colour onto surfaces. Pigments are grounded into fine powder and thoroughly
mixed and dispersed into the binding liquid.
Types of pigments include
o Organic pigments – made from animal/vegie sources, are likely to fade eg. Tyrolean purple –
marine slug, Lake pigments such as indigo lake, rose madder lake and carmine lake. Benefit
of a lake pigment is that it gives a softer colour, although it is unstable when exposed to light.
o Inorganic/mineral pigments – natural pigments prepared from minerals eg. ochre,
azurite/malachite. Benefit of inorganic pigment is that it gives a more opaque colour,
although many contain toxic heavy metals.
o Synthetic pigment – 19th/20th century produced eg. Prussian blue, Titanium oxide, cadmium
yellow. Many are expensive to produce.
a liquid to carry the pigment
The liquid binding medium is the vehicle that transports the dispersed pigment and binds it onto a
When the medium dries, unsaturated bonds crosslink after oxidising to form a three dimensional
polymer network. The pigment becomes trapped in the liquid and so it remains adhered onto the
surface. The medium should also allow for expansion and contraction so that the paint will not crack.
o Egg tempera – from 14th century. Egg yolk is misted with pigment – tough and long lasting
o Gum Arabic – principle binder in watercolours (plant resin from acacia plants)
o Linseed oils – best of drying agent, specially prepared for use in oil paint
o Alkynols – oil modified synthetic resins, dry quickly and have good adhesive and nonyellowing properties
o Acrylics – used in quick drying synthetic paints. Durable, flexible and suitable for use on
paper, panels & unprimed canvas
Describe an historical example to illustrate the relationship between the discovery of new mineral
deposits and the increasing range of pigments
Modern pigment industry started in the 18th century. This increased understanding of reactions and
stoichiometry, allowing production conditions such as pH and temperature to be controlled and
altered. This increased the range of colours and sizes of pigments available.
New metals have been discovered, such as cadmium, cobalt, chromium, manganese and
molybdenum. New synthetic colour pigments have been produced, such as cadmium red, cadmium
yellow, manganese blue, cobalt blue, chrome yellow and molybdenum red. Prussian blue and cobalt
blue were developed during this time as an alternative to the expensive Ultramarine blue.
A specific example is the discovery of an orange-red colour mineral called chrocite (PbCrO4) in 1770s
in Siberia, which contained chromium. This mineral is rare. However, discovering large deposits of
chromite (FeCr2O4) in US in 1820 led to manufacture of chromium compounds.
Before the 19th century, yellows were typically compounds of lead, tin and antimony. A richer golden
yellow is orpiment, arsenic sulphide.
Newer yellows include chrome yellow (PbCrO4), strontium chromate (pale yellow), barium chromate
(lemon yellow), zinc chromate (greenish yellow). These have replaced toxic yellow orpiment.
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
Most colour pigments are sulfides or oxides of transition metals.
Naples Yellow
Yellow orpiment
Zinc white
Antimony oxide
Arsenic (III) sulfide
Copper carbonate
Copper carbonate
Iron oxide
Iron oxide
Mercury Sulfide
Zinc oxide
Sb (group V)
As (group V)
Cu (transition)
Cu (transition)
Fe (transition)
Fe (transition)
Hg (transition)
Zn (transition)
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.
Pigments are mineral salts (oxides, phosphates, carbonates.) Naturally occurring and insoluble
Differences is the colour of ochre is due to the presence of different amounts of oxyhydroxides
(mainly goethite) and oxides (hematite)
Water is driven off when siennas and umbers are placed in a fire, changing the colour.
Kaolinite is a soft clay silicate mineral produced by chemical weathering of aluminium silicate
minerals like feldspar
Deep Brown
Red Brown
Name of Pigment
Red Ochre
Yellow Ochre
Brown Ochre
Manganese (IV) Oxide
Bone Black
Chemical Compound
Fe203 (45-70%) Manganese Dioxide (5-20%)
Fe2O3 (50%) Manganese Dioxide (<1%)
CaCO3 (50-98%) and various minerals (clay, hematite, mica, quartz
and pyrite)
Carbon (10-20%) Ca3(PO4)2 (70-80%)
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.
Name of Pigment
Chemical Compound
Yellow Orpiment
Red Orpiment
White lead
Yellow & Red
Ancient Egyptians used red pigment for rouge and lipstick, and green, blue, black and yellow
pigments in eye makeup. Ancient Romans used white lead as face powders and vermillion (HgS) as
red makeup.
Orpiment contains Arsenic. Acute toxicity (immediate) is high for inorganic arsenic compounds, with
symptoms including nausea, diarrhoea, cyanosis, hallucinations and cardiac arrhythmia. Symptoms
of chronic arsenic poisoning include depression, numbness, sleeping disorders and headaches.
Mercury is a cumulative poison that is highly toxic by ingestion and inhalation in high doses. Early
symptoms may include irritability, excitability and insomnia. With long lasting exposure, tremors
develop, causing severe behavioural and personality changes and loss of memory, severely affecting
the central and peripheral nervous system
Malachite may cause gastrointestinal discomfort: may be harmful if inhaled; eye irritant
Galena contains lead which is a cumulative poison that can lead to abdominal pain, muscular
weakness and fatigue and headache. It can cause nervous system disorders.
White lead is a cumulative poison that is toxic by inhalation and ingestion, causing brain damage in
children and irritating skin, eye and lungs.
Use of these pigments were harmful as they were used on skin, so they would have been affected by
the symptoms described due to prolonged accumulation of these poisonous chemicals
Identify data, gather and process information from secondary sources to identify and analyse the chemical
composition of an identified range of pigments.
Time Period
Name of Pigment
Type of pigment
Chemical Composition
Natural Inorganic
Chalk White
Natural Inorganic
Natural inorganic
Natural inorganic
Synthetic inorganic
Natural inorganic
Naples Yellow
Carmine Lake
Ultramarine Blue
Synthetic Ultramarine Blue
Cobalt Blue
Chrome Yellow
Titanium White
Natural inorganic
Natural organic
Natural organic
Synthetic inorganic
Natural organic
Synthetic inorganic
Natural inorganic
Synthetic inorganic
Synthetic organic
Synthetic inorganic
Synthetic inorganic
Refined mineral
Fe2O3 (45-70%) MnO2 (15-20%)
Pb(SbO3)2 or Pb(SbO4)2
C22H20O13, C16H10O8
CoO in silicate
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
Identify Na+, K+, Ca2+, Ba2+, Sr2+ and Cu2+ by their flame colour.
Name of Element
Colour in flame
Brick Red
Yellow Green
Scarlet Red
Emerald Green
Colour through blue glass
Purple red
 Perform first-hand investigation to observe the flame colour of Na+, K+, Ca2+, Ba2+, Sr2+ and Cu2+.
Qualitative test determining the identity or possible identity of a metal or metalloid ion
Pour 10ml conc. HCL into 50 mL beaker. HCl is used to clean the wire, as impurities will make results
Caution: HCl is highly corrosive and so it irritates eyes and skin, possibly causing burns if contacted. If
ingested, lining of lungs and oesophagus may be damaged. Wear safety glasses and protective shoes
and clothing.
Dip platinum wire loop in the acid and rinse with distilled water, and heat the loop in the outer edge
of the burner flames. Continue until no colour is observed in flame.
Dip clean wire loop into one of nitrate powders. Place loop in outer eye of burner flame and move
loop up and down. Note colour of flame.
Repeat with a different nitrate powder.
Cobalt blue glass filters out yellow colour (particularly Na+), allowing some flame colours to be
Explain the flame colour in terms of electrons releasing energy as they move to a lower energy level.
An atom does not radiate energy when its electrons are in a fixed orbit
Electrons move from a lower orbit to a higher orbit if it is given energy by an external source. The
orbit that it jumps to depends on the amount of energy the electron absorbs. The energy is absorbed
as a photon of electro-magnetic radiation. The photons energy equals the difference in energy
between the two stationary states of orbit.
The colour of a flame is created by a strong emission in the line spectrum of the element.
The electron will not remain in this orbit, but will decay back to its original orbit. They decay may be
in a single or multiple steps. Each time the electron decays to original orbit, energy is loss and a
photon in emitted. The specific wavelength of colour of the photon depends on amount of energy
Identify that, as electrons return to lower energy levels, they emit quanta of energy which humans may
detect as a specific colour.
The wavelength of EM radiation depends on the energy difference between the energy levels
In hydrogen the visible wavelength are called the Balmer series, referring to the transitions from 3 rd,
4th, 5th or 6th level to the 2nd (n=2) level. The human eye is sensitive to the visible region of the EM
spectrum. Different effects of infra-red and ultraviolet are due to difference in the energy of the
Explain why excited atoms only emit certain frequencies of radiation.
When an electron drops from an orbit farther from the nucleus to one closer to the nucleus, it emits
a photon of specific energy.
Unlike white light from sun, light given off by an energetically excited atom is not a continuous
distribution of wavelengths. When passed through a narrow slit and then through a prism, visible
light emitted by an excited atom is found to consist only of a few wavelengths.
The energy is not continuous but is quantised, that is electrons can be in one fixed orbit or another
fixed orbit, but not between orbits. Each of the orbits relate to a certain energy level of the electron.
The excited atoms emit certain frequencies of radiation (photons) because the differences between
energy levels of the orbits are fixed.
Planck proposed that electromagnetic radiation is transmitted and absorbed in discrete units, or
quanta, called photons. Energy carried by a photon of radiation is proportional to frequency. Greater
the difference in orbit radius, greater is the energy of the emitted photon. Amount of energy
between two energy levels is calculated using equation: E = hf, where h is Planck’s constant
Distinguish between the terms spectral line, emission spectrum, absorption spectrum and reflectance
Spectral Line: a gaseous sample of an element is dissociated into atoms and excited by an electric
discharge. The emitted light passes through a slit and then a prism or diffraction grating, dispersing
light into individual wavelengths, which are represented as spectral lines
Emission spectrum: produced when atoms that have been excited to higher energy levels emit
photons characteristic of the elements as they return to lower energy level. This does not occur in
ground state as atoms do not emit wavelengths as the electrons are at the lowest energy level. The
emission spectrum is unique to each element as they have different atomic structure/configuration,
and thus have different energy changes and wavelengths. In the visible part of the EM spectrum, the
emission spectrum is seen as bright coloured lines that are specific for an individual element (colour
lines on black background)
Absorption spectrum: produced when atoms absorb photons of certain wavelengths and become
excited from lower to higher energy levels. It appears as black lines against a bright background or
continuous spectrum.
Reflectance spectrum: produced as white light, when shone, is scattered without absorption from
the surface. Light as a function of wavelength
Absorbed wavelength in Observed
that has been reflected or scattered from a nm (colour)
complementary colour
solid, liquid or gas can be studied to determine 400 (violet)
Greenish yellow
the makeup of a mixture of pigments in paint. 450 (blue)
At different wavelengths, different pigments
570 (yellow-green)
have different reflection and absorption
600 (orange)
coefficients. The substance exhibits the colour 650 (red)
complementary to that absorbed.
Describe the development of the Bohr model of the atom from the hydrogen spectra and relative energy
levels to electron shells.
Electricity passing through a discharge tube containing H gas split H molecules into individual atoms
Bohr applied Planck’s concept of the quantisation of energy and Rutherford’s planetary model of the
atom (neutrons and protons occupy dense central region and electrons orbit the nucleus)
Describes the electrons in atoms as having discrete amounts of energy and that move from a lower
energy level to a higher energy level when a photon is absorbed with sufficient energy
When excited electrons return to lower energy level, a photon of radiation, or a specific frequency of
electromagnetic radiation is emitted.
The main features of the model:
o Atoms with electrons in fixed orbits with definite and discrete energy levels that may be
identified with a principal quantum number, n. Shell is also used to describe an energy level.
o Electrons in orbit do not radiate energy in ground state (lowest energy level)
o When excited electrons return to their ground state, the energy of the photon is exactly
equal to the difference in energy between the two levels in the atom.
o Amount of energy at a particular level depends on radius. The further the energy level from
the nucleus, the greater the energy
Lyman series: Transitions from excited energy states to the ground state; Balmer series: transitions
down to n=2. Occur in visible wavelength range.
Explain what is meant by n, the principal quantum number.
In the Bohr model, the principal quantum number is associated with the radius of the discrete
electrons orbit, which specifies the electrons energy.
Quantum numbers can have whole number values 1,2,3 …to infinity.
Each orbit corresponds to a different value of n, and as the radius of the orbit gets larger, the value
of n increases.
When the electron is closet to the nucleus (n=1) the atom is in its lowest energy level or ground state.
(i.e. greater probability of the electron being closer to the nucleus than n=2 level)
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.
The 3p level is split into two different energy levels by the magnetic energy of the electron spin in
the presence of the internal magnetic field caused by the orbital motion. This effect is called the spin
orbital effect.
Additional external magnetic field cause further splitting, called the Zaeman effect.
Visible line spectrum of the energetically excited sodium atom is dominated by closely spaced pair of
yellow lines called the doublet
Lines appear at 589.0 and 589.6 nm. The line at 589.0 has twice the intensity of the line at 589.6nm.
Represent the transition of electrons within the n=3 energy level from 3p energy level to 3s energy
Solve problems and use available evidence to discuss the merits and limitations of the Bohr model of the
Important ideas that are incorporated into our current model
- Electrons exist only in certain discrete energy levels, described by quantum numbers
- Energy is involved in moving an electron from one level to another.
- Excited atoms generate line emission spectra
- Predicts the observed frequencies in the line emission spectrum for hydrogen
- Able to determine ionisation energy for hydrogen atom and its atomic radius
- Does not predict the spectrum of any other element (does not work for atoms other than hydrogen)
- Weakness of Rutherford’s model is that any charged particle moving on a curved path emits EM
radiation, so the electrons lose energy and spiral into the nucleus. Thus electrons do not move in
circular orbits. Electrons are present in orbitals – areas around the nucleus where there is high
probably electrons exist.
- Cannot account for different intensity of spectral lines
- Cannot explain the hyperfine structure of spectral lines (eg. sodium doublet) as only principle
quantum numbers were used (no knowledge of subshells or orbitals)
- Cannot explain the splitting of spectral lines under the influence of a magnetic field (Zaeman effect)
Outline the use of infra-red and ultra-violet light in the analysis and identification of pigments and their
chemical composition
Infra-red (heat) radiation causes molecules to undergo stretching or bending vibrations, becoming
excited. Each group of atoms has its own characteristic set of vibrational frequencies, which can be
measured by passing infra-red light through them and recording at which wavelengths absorptions
Molecules absorb IR radiation at particular wavelengths depending upon what chemical bonds are
Comparing the IR spectrum produced by the sample with result from a known compound enables a
contaminant or the presence of a certain pigment to be detected.
Ultra-violet radiation cause electronic transitions to take place in the atoms, with electrons excited
and promoted to higher energy levels which are detected as absorption or emission spectra.
Explain the relationship between absorption and reflectance spectra and the effect of infrared and ultraviolet light on pigments including zinc oxide and those containing copper
Absorption Spectra
- Absorption is a transition from a lower level of a higher level with transfer of energy from the
radiation field to the atom or molecule.
- If a beam of white light is passed through a coloured substance, photons of certain wavelengths are
absorbed but those of other wavelengths pass through. When the transmittance (infra-red) or
absorbance (UV-visible) of a sample is measured and plotted as a function of wave no. cm-1 (infrared) or wavelength (UV-visible), an absorption spectrum is obtained
- The absorption spectra of elements show up as thin dark lines, whereas for molecules and complex
ions, it consist of broad bands rather than sharp lines
- Absorption spectroscopy is destructive but qualitative. Sample needs to be in solution. It is used
mainly to identify pigments containing metal ions.
- Double beam absorption spectrophotometers are commonly used to record IR and UV-visible
spectra. The detector measures the radiation passing through the sample and solvent, and a
comparison of the intensity of the two beams allows the absorption of radiation to be determined.
- In infra-red spectrophotometers, the source is a heated ceramic such as silicion carbide rod, and a
thermocouple (measures temperature) is used as the detector. Commonly used for qualitative
analysis, particularly useful for identifying the presence of various organic functional groups.
- For example, the IR absorbance spectrum shows the frequency and intensity of absorption peaks
that are characteristic of the molecular structure of the material and are used to study changes in
composition, purity and degradation.
- In ultra-violet-visible spectrophotometers, the source is a tungsten lamp or a deuterium discharge
tube, and the detector is a photomultiplier tube. It is a quantitative tool as absorbance is directly
proportional to concentration. Ideal for use to detect oil, tempura and acrylic polymers that form a
large network polymer over time.
Reflectance Spectra
- Reflectance measures radiation that is reflected from a surface, which can be compared with the
radiation reflected from a non-absorbing or white sample
- The reflectance spectrum is the complement of absorption spectrum, which can be compared to
known spectra for identification of pigments
- Reflectance spectroscopy is non destructive, and can be taken to the painting. IR can be used to
detect underdrawings.
- In infra-red reflectography, a lamp with a portion of its output in the cool near infra-red is shone
onto the surface of a painting. The beam of IR energy is absorbed at wavelengths where it is
normally absorbed, so the reflected light detected is the IR energy at other wavelengths. This is
compared with reflection from a material (a control) that does not absorb IR (transmit IR light) such
as ionic compounds like sodium chloride
In ultraviolet reflectography, white light is shone on the painting and then the spectrum of light that
is reflected is examined. The reflected radiation from a pigment surface is compared with reflection
from a material (a control) that does not absorb UV such as SiO2
Effect of IR and UV
- Zinc oxide fluoresces yellow in ultraviolet. It changes from white to yellow in infra-red. (returns to
normal colour on cooling)
- Copper: green malachite fluoresces dirty mauve in ultraviolet. Red copper (I) oxides, malachite and
vergedris changes to black copper (II) oxide in infrared. (permanent due to heating)
Gather, process and present information about a current analytical technology to:
Describe the methodology involved
Laser Microspectral Analaysis involves high energy laser light being focused on a surface to vaporise
a sample. The vapour is fed through a gap between two electrodes that sparks and excites the atoms
and ions, producing an emission spectrum that consists of line corresponding to the elements
evaporated from the sample surface, as the excited electrons return to ground states
Obtained spectra consist of lines corresponding to the elements evaporated from the sample surface,
so can be used to identify specific pigments
Raman spectroscopy uses intense monochromatic visible light provided by lasers. Laser light is shone
on a material, scattering light, and a tiny fraction of it is shifted in frequency as atoms in the
substance being analyses vibrate. It reveals the characteristic vibration frequencies of the atoms and
hence the chemical composition of the substance.
Assess the importance of the technology is assisting identification of elements in samples and in
Laser induced plasma spectroscopy (LIPS) is based on the above principal, and is capable of
performing trace element measurements in all solid. Process is non-destructive, requires minimal
sample penetration, qualitative, can identify more than one element at a time, and is highly sensitive.
Quantitative measurements can be obtained by constructing calibration curves, so that a number of
samples with known elemental composition can be referenced.
Raman spectroscopy is non-contact, non-destructive, characterises chemical composition of
pigments, minerals, binders and of painting grounds and different laser excitations enable various
possibilities for identification of organic and inorganic compounds.
Provide examples of the technology’s use
Laser analysis analyses elemental composition of pigments in restoring paintings and can identify
short-lived chemical reactions. It can identify the validity of an art work.
Raman spectroscopy is used in forensic analysis, analysis of art paintings, and can give information
about the age, production processes, restorations and falsifications of art.
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
Identify that each orbital can contain only two electrons
Cannot be precise about the exact location of an electron with a given energy. The region around the
nucleus through which an electron with a given energy may move is called an orbital
Pauli Exclusion Principle states that no more than two electrons can occupy each atomic orbital and
those two electrons must have opposite spin. (no two electrons in an atom may have identical sets
of four quantum numbers) The four quantum numbers:
o principal-quantum (n) indicating orbital energy (size)
o subsidiary (l) indicating orbital shape
o magnetic (mL) indicating orbital orientation
o Spin (ms) indicating direction of orbital spin
Define the term sub-shell
Shells are distinct energy levels of atoms in which electrons are located. These a denoted by their
principal quantum number n=1, n=2 etc.
Each principal energy level or shell consists of a number of energy sublevels, or sub-shells, with
slightly different energies. The number of subshells in any shell is the same as the principal quantum
number of the shell. Subshells designate the shape of the orbital.
Sub-shells are closely spaced energy levels that consist of a set of atomic orbitals. They are a
consequence of two types of interactions – nucleus – electron attractions and electron –electron
o Lowest energy subshell: s-subshell – spherically symmetric, contains one s-orbital
o Second lowest energy subshell: p-subshell – dumb-bell shaped, contains three p-orbitals
o Third lowest energy subshell: d-subshell – contains five d-orbitals
o Fourth lowest energy subshell: f-subshell – contains seven f-orbitals
Outline the order of filling of sub-shells
Place the electrons into orbitals starting with the lowest energy orbital first
Place a max of two electrons in each orbital, but where more than one orbital with the same energy
is available, place on electron in each orbital before pairing electrons up
1s  2s  2p  3s  4s  3d  4p  5s 4d  5p  6s  4f  5d  6p  7s  5f  6d
Identify that electrons in their ground-state 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
When an electron is in the orbital closest to the nucleus it is in its ground state or lowest energy level.
Electrons are only able to move to a higher energy level when they absorb a photon whose energy
equals the difference in energy between the two stationary states.
The atom then would be in an excited-state electronic configuration
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 listed in order of increasing atomic
number display similar outer shell or energy level
electron configuration at regular intervals
Electrons with similar electron configurations in
their outermost shells display similar chemical
Explain the relationship between the number of electrons in the outer shell of an element and its
Electronegativity is the measure of the ability of an atom in a molecule to attract electrons to itself.
Expressed as a relative scale. Fluorine, the most electronegative element, is assigned a value of 4.0
Increases across periods from left to right as nuclear charge increases. Extra electrons being added
are not fully shielded from the increased number of protons in the nucleus. Reflects the increasing
trend across a period for elements to gain electrons and achieve a noble gas electron configuration.
Noble gases have no electronegativty as they have no tendency to form molecules with other atoms.
Decreases down the group since the size of the atom increases, which diminishes the attractive force
of the shielded nucleus.
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
Absorption of energy by an atom leads to a change in its electron configuration as valence or outer
shell electron is promoted to a higher energy orbital
Ionisation energy is the energy needed to remove the negatively charged electron from the
electrostatic attraction of the positively charge nucleus. It is the energy needed to remove the
outermost electron from a mole of gaseous atoms or ions.
First ionisation energy
- Ionisation energy gradually increases across a period as each successive element has one more
proton, which is located at nearly the same distance from the nucleus. Valence electrons are more
tightly held by the nucleus due to the increase nuclear charge, simultaneously decreasing the atomic
size. However, there are some irregularities
o Despite nuclear charge on B>Be, ionisation energy for B is lower than Be, due to 2p subshell
having a slightly higher energy than 2s orbital. P electron is well shield by the 2s electron. It
requires less energy to remove a single p electron than to remove a paired s electron from
the same energy level.
o Ionisation energy of nitrogen is more than that of oxygen. 3 unpaired electrons in the 3
outermost p orbitals. Half-filled 2p orbitals is more stable than oxygen.
o Oxygen has slightly lower ionisation energy as it takes slightly less energy to remove the first
paired electron than to remove an unpaired p electron as there is greater electrostatic
repulsion between two paired p electrons making it easier to remove.
Ionisation energy decreases down a group as the outer electrons are further from the nucleus and
the number shielding the outer electrons from the nuclear charge also increases. The outer electrons
are attracted by a similar positive charge but at a progressively greater distance.
Successive Ionisation Energies
- Are larger values because electrons are being removed from progressively larger values
- Successive electrons are removed from their orbitals in the reverse order to which they’re filled.
- The greater the charge, the greater the energy required to remove electron
- Removal of an electron from an energy level closer to nucleus results in increase in ionisation energy
- After the valence electrons have been removed and a noble gas configuration is reached, there is a
steep rise in ionisation energy for the next ionisation.
Process information from secondary sources to analyse information about the relationship between
ionisation energies and the orbitals of electrons
Factors affecting ionisation energy
- Ionisation energy decreases with the increase in atomic size as the outermost electrons are held less
tightly by the nucleus
- Increases with increase in nuclear charge as the attractive force between the nucleus and the
electron makes it more difficult to remove an electron
- Outermost electrons which are shielded from the nucleus by the inner electrons cause the outer
electrons to not feel the complete charge of the nucleus, resulting in a decrease in ionisation energy
- More penetration power of electron causes it to be closer to the nucleus, resulting in an increase in
ionisation energy
- The more stable the electronic configuration, the higher is the ionisation energy. i.e. elements with
high ionisation energies are helium, neon, argon. Elements with low ionisation energies: lithium,
sodium, potassium as they only have one electron in their outer shell
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
Hund’s rule states that if two or more orbitals with the same energy (in the same subshell) are
available, one electron goes into each orbital until all are half-full, keeping the spins parallel until
forced to pair by lack of additional empty orbitals. (no. of electrons with same spin is maximised)
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
Transition metals occupy the ‘middle’ d block of the periodic table.
They have properties that are intermediate between the s-block elements and the p-block elements.
Define the term transition element
Metal in which the available electron energy levels are occupied so that the d-orbitals are containing
less than its maximum number of 10 electrons per number
Must have at least one ion with a partially filled d shell. Scandium and zinc are d-block elements, but
are not transition metals
They occur in the fourth, fifth and sixth rows of the periodic table and result from the filling of the dsubshell of the third, fourth and fifth shells. Each d-subshell includes five orbitals and can
accommodate up to 10 electrons.
Explain why transition metals may have more than one oxidation state
2,3,4,5 2,3,4,6 2,3,4,6,7
Maximum number of oxidation states occurs at manganese (+7) and decreases on either side.
Maximum oxidation states equal total number of 4s and 3d electrons in the atom
Transition metals can lose electrons from both the 3d and 4s subshells, which have similar energies.
The +2 oxidation state is due to loss of two 4s electrons. Oxidation states above +2 results from
additional loss of 3d electrons when simple charged ions are formed, or when transition metal is
bonded to more electronegative elements such as oxygen.
Transition metals with high oxidation states tend to be strong oxidising agents. Generally, maximum
oxidation states are found when metals are combined with the most electronegative elements (O, F)
Scandium and Zinc show one oxidation state (not transition metals)
Copper is a transition metal since Cu2+ has an incomplete d orbital
The formation of ions depends on the amount of energy provided to ionise the metal. (ionisation
energy) A compound is most stable when the ion is more highly charged and releases more energy
Account for colour changes in transition metal ions in terms of changing oxidation states
The d-orbitals of a free transition metal atom or ion are
However, when transition metal form complex ions, the d
orbitals of the metal interact with the electron cloud of ligands,
so the d-orbitals become non-degenerate. They are split into sets
of orbitals separated by energies that correspond to the
wavelengths of EM radiation in the visible spectrum. d orbitals
therefore have slightly different energies and are incompletely
Lemon yellow
Colour arises from the absorption of photons of light by electrons in the transition metal ions. Small
energy differences between d orbitals are similar to the energies of photons of visible light. The
absorption of photons of appropriate frequency results in electron being excited from a lower to a
higher energy level
As the oxidation state of the metal increases, so does the amount of splitting of the d orbitals, and
therefore require more energy to reach higher energy levels. Changes of oxidation state change the
colour of the light absorbed and also the observed colour of the compound (which is complement to
the colour absorbed)
o Red is removed from Chromium – 2+ ion [Cr(H2O)6]2+ to give cyan colour
o Yellow is removed from Chromium – 3+ ion [Cr(H2O)6]3+ to give darker blue 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
Oxidation states of +4 to +7 exist only when the element is covalently bonded or forms complex ions.
Complex ions (CrO42-, Cr2O72-, MnO4-) are strong oxidising agents (electron acceptors) because many
contain oxygen atoms (only have six electrons in outer shell) which is electronegative
Strong oxidising agents in acidic solutions are widely used as oxidants in chemistry such as the
dichromate ion (6+) and the permanganate ion (7+). Acidified potassium permanganate readily
reacts with reducing agents such as halide ion, Fe (II) and bromide ion.
The oxides of transition metals in higher oxidation states are acidic (want to gain electrons) and
therefore strong oxidising agents. The oxides of lower oxidation states are more basic and make
good reducing agents
Process and present information from secondary sources by writing electron configurations of the first
transition series in terms of sub-shells
21. Sc
22. Ti
23. V
24. Cr
25. Mn
26. Fe
27. Co
28. Ni
29. Cu
30. Zn
 Perform a first-hand investigation to observe the colour changes of a name transition element as it
changes in oxidation state
Oxidation states of vanadium: reduction of V5+ to V2+
- In a 250ml conical flask, dissolved 3g of ammonium vanadate (NH4VO3) in 100ml of
1M NaOH. Acidify solution by adding 75mL of H2SO4.
- Pour off 20mL of yellow solution into a large test tube
- Add 6 granules of zinc into the conical flask and add rubber stopper.
- Swirl gently, and save 20mL of solution at each successive colour change
- All
- +5 to +4
+4 to +3
+3 to +2
Zn  Zn2+ + 2eVO3- + 4H+ + e-  VO2+ + 2H2O
Zn + 2VO3- + 8H+  2VO + 4H2O + Zn2+
VO2+ + 2H+ + e-  V3+ + H2O
Zn + 2VO2+ + 4H+  2V3+ + 2H2O + Zn2+
V3+ + e-  V2+
Zn + 2V3+  Zn2+ + 2V2+
Solve problems and process information from secondary sources to write half-equations and account for
the changes in oxidation state.
A change in the oxidation state is a loss or gain in electrons. Writing reduction half equations:
MnO4- (+VII)  Mn2+ (+II)
Balance the atoms
MnO4  Mn + 4H2O
Balance oxygen with water
MnO4- + 8H+  Mn2+ + 4H2O
Balance hydrogen with H+
MnO4- + 8H+ + 5e+  Mn2+ + 4H2O
Balance with charge of electrons
 Choose equipment, perform a first-hand investigation to demonstrate and gather first-hand information
about the oxidising strength of KMnO4
MnO4- (+VII) are strong oxidants as they are bonded to the very electronegative O atoms, which, in
redox reactions bond to H atoms to form H2O. The electrons taken by the O atoms as they leave are
replaced by a species being oxidised.
Add equal quantities of 0.01 M potassium permanganate and 1M H2SO4 to make acidified KMnO4
Use acidified potassium permanganate to test its oxidising ability on the other reagents. Test
solutions directly with 2-3 drops of potassium permanganate in a test tube. Solutions to test:
o Halides: 0.5 M solutions of potassium iodide (purple to orange), potassium bromide (purple
to pale yellow), potassium chloride (remains purple)
o few crystals of iron ammonium sulfate dissolved in a few mL of water (remains colourless)
o small pieces of metal: Zn powder, Mg, Cu turnings, Sn, Fe (purple – brown – colourless)
Wear eye and skin protection. Potassium permanganate stains skin and clothing. Sulfuric acid is
strongly corrosive to skin and eyes. Potassium iodide, bromide and chloride and iron (II) ammonium
sulfate are toxic if ingested
Reduction equation: MnO4- + 8H+ + 5e-  Mn2+ + 4H2O
5. The formation of complex ions by transition metal ions increases the variety of
colour compounds that can be produced
Explain what is meant by a hydrated ion in solution
A hydrated ion is an ion in which a specific number of water molecules is associated with each
formula unit, for example, colourless Cu2+ forms blue Cu(OH2)42+ in aqueous solution
It is formed an ionic solid dissolves in water. The ions dissociate and are surrounded by water
For many cations the ion is surrounded by a fixed number of tightly bound water molecules
Describe hydrated ions as example of a coordination complex or a complex ion and identify examples
Hydrated metal ions are examples of complex ions
A complex ion contains a central metal ion surrounded by a group of anions or molecules called
ligands. They are commonly formed between transition metal ions and anions or polar molecules
with at least one lone pair of electrons
Compounds containing complex ions are called coordination compounds
Describe molecules or ions attached to a metal ion in a complex ion as ligands
Ligands are molecules or anions attached to the central metal cation (transition metals), forming
Examples of Ligands include chloride, cyanide, polar molecules: water and ammonia
Coordinate covalent bonds form between the ligands and the transition metal ion using the electron
lone pairs on the ligands
The number of coordinate covalent bonds between the ligands and the central metal ion is known as
the coordination number of the metal ion. Most common coordinate number in transition metals is
A Lewis acid is an ion or molecule that can accept a pair of electrons. A lewis base is an ion or
molecule that can donate a pair of electrons
Ligands are therefore Lewis Bases, as they have at least one atom with a lone pair of electrons and
can donate that pair of electron.
Transition metals act as Lewis acids, accepting electron from ligands.
Explain that ligands have at least one atom with a lone pair electrons
Ligands have at least one atom with a lone pair of unbonded electrons
Ligands which bond using the electron pair of a single donor atom such as NH3, H2O, Cl-, CN- are
monodentate ligands
Ligands that bond through electron pairs on more than one donor atom are polydentate, such as the
oxalate ion -OOC-COO- and the triphosphate ion [P3O10]5-
Identify examples of chelated ligands
Polydentate ligands are also referred to as chelating agents due to their multi-point attachment to a
metal ion. Chelated ligands have more than one donor atom. Ligands with two or more donor atoms
tend to form rings in the complex ion.
Chelating ligands bind much more tightly to their metal cations than do ligands with one donor atom.
Stabilisation of a metal complex by a ligand with more than on donor atom is known at the chelate
Chelating agents hides or removes the metal ion
An example is EDTA. It is used to treat lead poisoning due to its chelating ability. With six donor
atoms, EDTA is a hexadentate ligand that wraps around a metal ion (eg. Hg) and remove the metal
ion by forming a complex ion.
Discuss the importance of models in developing an understanding of the nature of ligands and chelated
ligands, using specific examples
Crystal Field Theory
- Shows ionic bonding and repulsion due to like charges.
- Accounts for the colour and magnetic properties of transition metal complexes in terms of the
splitting of the energies of the metal ion, d-orbitals by the electrostatic interaction of ligands
- Assumptions are that the interactions between the metal ion and ligands are purely electrostatic
(ionic). Also, electron-electron repulsions between electrons on the metal/ligan is present,
influencing energy of the d-orbitals
- To minimise electron repulsion, the d-orbitals slightly change orientation. This results in unequal
energy states and different wavelengths absorbed/reflected
- Ligands cause splitting of d-orbitals resulting in different pattern of energy levels. There is also
different orientation of d-orbitals since ligands are arranged around the central ion differently, which
changes the symmetry of the electric field around ion.
Valence Bond Theory
- Shows transition and ligand bonded via coordination bond – covalent
- The formation of complex ions depends on
o the orbitals available for coordinate covalent bond formation
o tendency for ions or groups to share a pair of electrons
o the number of molecules of ions that can be placed around the central ion
o the geometry assumed by the ligands and the metal ion
- Empty orbital hybridise to form a mixed set of orbitals that point towards corners and accept pairs of
non bonding electrons from a ligand, forming a complex in which a valence shell is filled
- Explains chemical properties (structure and number of bondings of coordination complexes)
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
Coordination number indicates the number of coordinated bonds (the number of ligands linked to
central metal).
Geometric shape of a complex ion is Ligand
determined by number of coordinated H2O
4-square NH3
5trigonal OHHydroxo
pyramidal, 6-octahedral
Negative ion metal modified to ending F
in –ate
See below for Lewis structures of complex ions
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
Strong field ligands (eg. CN ) produce large crystal field
splitting and low spin complexes (absorb higher energy)
- Weak field ligands (eg. Cl-) produce small crystal field
splitting and high spin complexes (absorb lower energy)
- Complex appears as the complementary colour to that
absorbed. Colour wheel: ROYGBV
Coordination complex ion
Colour (colour absorbed)
Dark blue (green)
Hexaaquachromium (III) ion
Blue green (purple)
Pentaaquasulfate chromium (III) ion
Green (purple)
Tetraaquachlorochromium (III) ion
Blue-grey (purple)
Hexahydroxychromate (III) ion
Blue-grey (red)
Triaquatrihydroxychromium (III) ion
Hexaamminechromium (III ion
pink-purple (yellow)
Relative strength
Strong (large splitting)
Strong (large splitting)
Moderate (medium splitting)
Moderate (medium splitting)
Weak (small splitting)
Weak (small splitting)
Lewis Structure
Explanation of splitting:
- Ammonia has a strong ligand field so it results in a larger splitting of the d orbitals, causing a greater
gap between the orbitals. A high energy wavelength of yellow is absorbed, so the ion appears the
complement of yellow, appearing violet.
- Hydroxide has a weak ligand field so it results in a smaller splitting of the d orbitals, causing a smaller
gap between the orbitals. A low energy wavelength of red is absorbed, so the ion appears the
complement of red, appearing green.
- Water has a moderate ligand field.
Factors affecting the colour of a transition metal complex
- Geometry: splitting is greater if the ion is octahedral than tetrahedral
- Oxidation state: as the oxidation state of the metal increases so does the amount of splitting of the
- Ligand: different ligands have different effects on the energies of d-orbitals of the central ions. Some
ligands have strong electrical field which cause a large energy gap when the d-orbitals split into two
groups others have much weaker field producing smaller gaps.