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

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 pigmento 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

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) oxideo Charcoal could be obtained through burning grass or bark

Colour Pigment Mineral Chemical compositionRed Red ochre Haematite Fe2O3

Yellow Yellow ochre (goethite) Goethite Fe2O3.H2OBrown Brown ochre (limonite) Limonite FeO(OH)White White clay Kaolin Al2O3.2SiO2.2H2OBlack Charcoal Graphite C

Manganese (IV) oxide Pyrolusite MnO2

Describe paints as consisting of:

• the pigment

• a liquid to carry the pigment

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

evaporatedo 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 driedo 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 timeo 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

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 evaporatedo 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)

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 evaporatedo See below for a more detailed description of the preparation of panel painting

(under mediaeval artwork)

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 oveno 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:

• cave drawings

• self-decoration including cosmetics

• preparation of the dead for burial

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 painto The paints used in cave drawings generally used saliva, honey from wild bees, or

tree gums as the binder, and water as the medium

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 pigmentso 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 UseMalachite Bright green CuCO3.Cu(OH)2 Eye paintAzurite Blue 2CuCO3.Cu(OH)2 Make-upGold Gold Ag Body paintCinnabar Bright Red HgS LipstickOrpiment Rich lemon-yellow As2S3 Eye shadowEgyptian blue Blue CaO.Cu.4SiO2 Body paintKohl Black Sb2S3 EyelinerWhite lead White 2PbCO3.Pb(OH)2 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 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

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 smoothPREPARATION OF PIGMENTS

Each colour used on in the artwork comes from different pigments, each of which has to be extracted separately.

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

mediumADHESION 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 ono The area to be gilded was coated with iron (III) oxide and egg white, which was then

polished and set hardo 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

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 pigmentsNAPLES YELLOW

Chemical composition: Pb3(SbO4)2 [lead (II) antimoniate] Colour: Yellow

Discovery: Naples Yellow was originally discovered in mineral deposits on the slopes of volcanic Mt Vesuvius, but was synthetically produced during the 17 th 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 Colour Chemical compositionGypsum White, grey CaSO4.2H2OCerussite White, grey PbCO3

Stibnite Lead-grey, blackish Sb2S3

Galena Lead-grey PbSGraphite Black CCinnabar Red, brownish red HgSMalachite Bright green Cu2(CO3)(OH)2

Azurite Azure blue to dark blue Cu3(CO3)2(OH)2

Orpiment Lemon yellow – brownish yellow As2S3

Turquoise Bluish green CuAl6(PO4)4(OH)8.4H2ORed Ochre/ Haematite Earthy to bright red Fe2O3

Yellow ochre Yellow Fe2O3.H2OEgyptian blue Blue CaO.CuO.4SiO2

Prussian blue Blue 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

o The discovery of new elements and processes since the 16th century have allowed synthetic inorganic pigments to be produced, such as Prussian Blue

o 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

Below is another table of a range of pigments used throughout history

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 rangeo Copper produces colours in the blue-green rangeo Cobalt produces colours in the yellow-violet rangeo Chromium produces colours in the red-yellow range

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

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 positively-charged 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 possibleo Electrons in stable orbits do not radiateo 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

See above for an explanation of the Bohr model of the atomMERITS

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 precisionLIMITATIONS

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 nucleusASSESSMENT

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

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

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 Flame colourNa+ YellowK+ VioletCa2+ Orange-redBa2+ Apple greenSr2+ RedCu2+ Green-blue

Explain the flame colour in terms of electrons releasing energy as they move to a lower energy level

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

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

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

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

Below is the emission spectrum for sodium

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

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

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 spectrumo Near infra-red radiation lies closest to 700nm, and is commonly called ‘cool’ infra-

redo 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

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 pigmentsULTRA-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 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 heatedo 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

Laser microspectral analysis is a current technique used to identify and analyse elements and molecules in pigmentsMETHODOLOGY 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

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 determinedEXAMPLES 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 artworkASSESSMENT

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 sampleso It requires minimal sample preparationo It can rapidly identify more than one element at a time, as opposed to AASo 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+

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 saltso General safety precautions of using a Bunsen burner => wear safety goggles, spray

the flame from a safe distanceo Flame colours produced may be bright and cause eye damage => use low

concentrations, do not look at the flame for long periods of timeRESULTS

Salt Cation Flame colourNaNO3 Na+ YellowKNO3 K+ VioletCaNO3 Ca2+ Orange-redBaNO3 Ba2+ Apple greenSrNO3 Sr2+ RedCuCl2 Cu2+ Blue-green

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

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

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 electron-spin 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

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

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 Maximum number of electrons

Shape

s 1 2 Sphericalp 3 6 Dumb-belld 5 10 Complexf 7 14 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

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

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

To determine the order of filling sub-shells, apply the following principleso Electrons are placed into orbitals starting with the lowest-energy orbital firsto 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

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

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

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, sub-shells, 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

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

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

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 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

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 decreaseso Nuclear charge increases down a group, so the nuclear force on the valence

electrons increaseso The inner electrons effectively shield the nuclear charge from valence electrons, so

the outer electrons experience a decreased nuclear charge, which is called the

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

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

Transition metals occupy the d-block of the Periodic Table, which corresponds to the filling of the d orbital

Define the term transition element

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

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

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

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

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 half-equations and account for the changes in oxidation state

NOTES ON DETERMINING OXIDATION/REDUCTION HALF-EQUATIONS Balance the atoms on each side of the equation

o Oxygen is balance with H2Oo 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+

To balance the four oxygen atoms on the reactants side, add four water molecules to the products side

To balance the eight hydrogen atoms on the products side, add eight hydrogen ions to the reactants side

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

Finally, write the relevant states

OTHER EXAMPLES OF CHANGING OXIDATION STATE

Change in oxidation number Half-equationCr(+6) to Cr(+3) Cr2O72-(aq)+ 14H+(aq) + 6e-   2Cr3+(aq) + 7H2O(l)

Fe(+3) to Fe(+2) Fe3+(aq) + e-   Fe2+(aq)

Cu(+2) to Cu(+1) Cu2+(aq) + e-   Cu1+(aq)

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 d-orbitals and oxidation states)

Account for colour changes in transition metal ions in terms of changing oxidation states

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

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

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 (Cr2O7

2-) has an oxidation state of +6, and manganese in permanganate (MnO4

-) has an oxidation state of +7o 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

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 waterRESULTS

Test tube Vanadium ion Oxidation state of vanadium

Colour

1 VO3- +5 Yellow

2 VO2+ +4 Blue3 V3+ +3 Green

4 V2+ +2 Violet

The electron configurations for the various oxidation states of vanadium are shown below:

The reduction half-equations for each step of the reaction are as follows:

The oxidation half-equation for each step of the reaction was the same:

The above redox reactions account for the changes in oxidation state for vanadium, as electrons in the d orbital are gainedACCURACY/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

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

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 reductionSAFETY:

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 OBSERVATIONKI/KBr/KCl Solution turned from purple to colourless with slight pink/brown

colourIron ammonium sulfate, Zn, Mg, Cu, Sn, Fe

Solution turned from purple to a red/orange solution

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

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 iono The above formula can also be written as Cu(OH2)4

2+, 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

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

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

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 ligandso 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

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

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

Chromium displays a wide range of colours in its complexes

Formula Colour[Cr(H2O)6]2+ Blue[Cr(H2O)6]3+ Violet[Cr(H2O)5Cl]2+ Green[Cr(OH)4]- 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

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 orbitalso The energy changes involvedo The rate at which chemical interactions occuro What initiates electron, ionic, and molecular interactionso 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