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CHEM4: Chemistry of Art 1. From earliest times, people have used colour to decorate themselves and their surroundings DESCRIBE PAINTS AS CONSISTING OF: THE PIGMENT A LIQUID TO CARRY THE PIGMENT Paints consist of pigment, binder, and medium Colorant pigments add colour to the paint, and extender pigments develop properties of the paint e.g. gloss o Hiding pigments e.g. titanium dioxide makes the paint opaque to protect it from UV light The binder (or resin) is the glue that holds the paint together. It imparts adhesion to bind the pigments together. The medium, or solvent, keeps the pigment and binder flowing together so that the paint achieves better adhesion to the surface o It affects drying time and film thickness. After the usually colourless medium dries, the coloured pigment and binder remain as a solid film o The medium must be viscous/thick 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 Pigments are insoluble so that, after the colourless medium dries, the coloured pigment remains o An insoluble pigment is not easily removed when exposed to rain or ground water (especially in rock/cave paintings) o An insoluble pigment used in cosmetics will not dissolve in perspiration While a pigment is coloured insoluble solid suspended in a medium, a dye is colouring matter dissolved in solution. A lake pigment is produced when a vegetable dye is dissolved in water, then precipitated out with an inert insoluble white powder (the binder) e.g. chalk. o Madder lake (pink-red): alizarin [Cl4H8O4] from root of the madder plant o Indigo lake (blue-violet): vegetable dye indigotin [C16H10N2O2] from the indigo plant o Crimson lake (pink-red): carminic acid [C22H20O13]from female cochineal insect 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) Structure of Early Paintings: Wood panel or canvas Ground o Layers of ground are applied consisting of gesso, which is a mixture of gypsum or chalk with animal glue. The gesso sets to a brittle, creamy layer. o Ground is applied because wood and canvas are too rough and absorbent to paint upon Underdrawing: an outline of the design to be painted Paint layers: artists up until the 19th century did not mix pigments, but rather superimposed layers of paint to achieve complex colour effects Varnish: protects the painting, and gives the colours clarity and depth of saturation Glass Colours: Metal oxide pigments normally used to paint on surfaces were finely powdered and added to the glass mixture before melting. This mixed the pigments within the bulk of the material, making coloured glass o E.g. cobalt oxide for blue, manganese oxide for purple, iron oxide for green/yellow o Impurities in these oxides extend the possible range of colours Flashing: semi-molten clear glass is dipped into molten red glass to form a light red coating o Used to make ‘ruby’ glass from the extremely dense red colour produced by copper (II) oxide o Patterns can be created by scratching away parts of the coloured layer to reveal the clear glass beneath Enamel pigments: Finely ground coloured glass mixed with a medium (e.g. wine/urine) is painted onto a piece of glass. Firing in the kiln causes the ground glass to melt and resolidify, effectively painting the surface beneath Staining: glass is painted with silver nitrate, then fired in an oven o Staining/firing the glass a different number of times produces various colours from light yellow to deep orange OUTLINE THE PROCESSES USED AND THE CHEMISTRY INVOLVED TO PREPARE AND ATTACH PIGMENTS TO SURFACES IN A NAMED EXAMPLE OF MEDIEVAL OR EARLIER ARTWORK Medieval Painting: ‘St John the Baptist with St John the Evangelist and St James’, by Italian artist Nardo di Cione (1365) Poplar wood backing, gesso, underdrawing Medium: egg tempera, egg yolk mixed with water o Pigments used were mixed with egg tempera, and painted on Background was gilded, i.e. coated with gold o Area to be gilded was coated with iron (III) oxide and egg white, polished after it had set, coated with gold leaf and egg white, then polished again Floor was gilded, then covered completely with red lead in egg tempera. The blue birds/flowers were painted on using ultramarine in egg tempera, and finally some paint was scraped away to reveal the gold layer St John the Evangelist (left) o Green robes are ultramarine mixed with lead tin yellow o Scarlet lining of cloak, and book, are vermillion Vermillion is produced by heating mercury and sulfur together in the flask, until the mixture vaporises and reacts to form mercury(II) sulfide, which condenses at the top of the flask. The condensed vermillion was ground to produce the red pigment. Exposure to sun has converted some of the vermillion into the black form of mercury (II) sulfide, darkening the colour St John the Baptist (middle) o Pink robes are crimson lake (from cochineal) with layers of white lead Crimson lake is produced from by boiling dried female cochineal insects in water to extract the carminic acid [C22H20O13], then precipitating out the pigment onto a clear insoluble powder St John the Baptist’s blue robe lining and St James’ (right) robes are natural ultramarine o Natural ultramarine is extracted by wrapping a mixture of ground lapis lazuli stone and melted wax/resins/oils in a cloth, and kneading it under a dilute lye solution of potassium carbonate KCO3. Impurities collect in the wax/resin/oil, and blue pigment particles form a precipitate which can be removed. o Vivid blue colour of natural ultramarine can only be maintained within aqueous solutions and egg tempera Flesh of the saints is brown and red earth pigments painted over an underpaint of green earth 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 Ochres: natural earths of silica and clay which owe their colour (red-yellow) to iron (III) oxide Red ochre: anhydrous iron (III) oxide [Fe2O3] Yellow ochre: hydrated iron (III) oxide, e.g. the mineral goethite [Fe2O3.H2O] o Burning yellow ochre drives off water, converting it to red anhydrous form Brown ochre: nearly pure limonite [FeO(OH)] Black Pigments: Manganese (IV) oxide [MnO2] Charcoal: mainly graphite [C] o From burnt wood or lamp soot White Pigments: Chalk: calcium carbonate [CaCO3] Gypsum: calcium sulfate dihydrate [CaSO4.2H2O] Kaolin: hydrated aluminium silicate [Al2O3.2SiO2.2H2O] OUTLINE THE EARLY USES OF PIGMENTS FOR: CAVE DRAWINGS SELF-DECORATION INCLUDING COSMETICS PREPARATION OF THE DEAD FOR BURIAL Cave Drawings: Rock art and cave drawings from over 17,000 years ago used natural pigments such as red ochre (haematite) for red, yellow ochre (goethite) for yellow, and charcoal for black. They mixed these pigments with saliva, fats, waxes, water or blood The Aboriginal people either used solid pigments directly to paint on cave walls, or used liquid mediums. o Mediums include orchid juice, wax/honey of wild bees, egg yolk (e.g. turtle or seagull), tree/plant gums and resins Self-decorations including Cosmetics: People used pigments (e.g. red ochre and yellow ochre) as body paint. The Egyptians used the red pigment cinnabar for rouge and lipstick, yellow orpiment for eye shadow, kohl for eyeliner and mascara, etc. The Romans used white lead as face powder Preparation of the Dead for Burial: The Egyptians painted the dead, and the containers in which the body was entombed, with pigments in much the same way as they were used for the living. In Tutankhamen’s tomb, a small paint box was found containing powdered gypsum, orpiment, haematite and malachite o In preparing the dead for burial, the Egyptians replaced internal organs with a mixture of resins from coniferous trees, and then treated the body with hygroscopic natron (a sodium salt mixture from dry lake beds) to dry and preserve the body. The Mayans covered graves and skeletal remains with cinnabar, because the colour red was associated with death and rebirth 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 The Egyptians used malachite, azurite, cinnabar, orpiment, and kohl. The Romans used soot, white lead, minium (red lead), vermillion, yellow massicot, and green-blue verdigris Pigment/mineral Malachite Cosmetic Eye/body colouring Health Risks Eye irritant, gastrointestinal discomfort, harmful if inhaled Cinnabar Composition basic copper carbonate [CuCO3.Cu(OH)2] basic copper carbonate [2CuCO3.Cu(OH)2] mercury (II) sulfide [HgS] Rouge, lipstick Orpiment Realgar Arsenic trisulfide [As2S3] Tetraarsenic tetrasulfide [As4S4] Kohl Stibnite: antimony (III) sulfide [Sb2S3] Yellow eye shadow Orange-scarlet eye/body colouring Used as eyeliner/mascara to darken eyebrows and eyelids Mercury is toxic by ingestion/inhalation/skin contact. Cumulative poison that causes CNS damage Arsenic is toxic by ingestion/skin contact. Suspected carcinogen Azurite Manganese (IV) oxide [MnO2] Galena [PbS] Black copper (II) oxide [CuO] Soot from fire Carbon black [C] Black eyeliner Symptoms: nausea, vomiting, diarrhoea. Antimony will cause liver/kidney damage if inhaled/ingested. Skin/eye irritant. Prolonged skin contact may cause skin burns. Symptoms: Sluggishness, weakness in legs, fixed facial expression, emotional disturbances. Harmful by inhalation/ingestion. May cause lung damage, CNS impairment, and reduce fertility in men Symptoms: abdominal pain, constipation, headache. Harmful by ingestion/inhalation. Cumulative poison may cause mental retardation and death. Skin irritant. Gastrointestinal irritation if ingested in large amounts, causing vomiting and diarrhoea. May cause ulceration of respiratory tract Respiratory irritant. Harmful by ingestion/inhalation. White lead Minium/red lead Basic lead carbonate [2PbCO3.Pb(OH)2] Lead tetroxide [PbO2.2PbO] Face powder to lighten face Face/body colouring May contain carcinogenic organic molecules. Symptoms: abdominal pain, constipation, headache. Lead is toxic by ingestion/inhalation. It is both an acute and a chronic poison. Cumulative poison causes mental retardation and death IDENTIFY THE SOURCES OF THE PIGMENTS USED IN EARLY HISTORY AS READILY AVAILABLE MINERALS Pigments used in early history were minerals from coloured earth and soft rocks. These include red ochre, malachite, vermillion, siennas, and Naples yellow Red ochre, anhydrous iron (III) oxide [Fe2O3]: ochre clay was mined from the ground, washed to separate ochre from sand by hand, then dried in the sun Green malachite, basic copper carbonate [CuCO3.Cu(OH)2]: Egyptians prepared malachite by selection, grinding and sieving of ore Cinnabar, mercury(II) sulfide [HgS]: prepared by crushing, then the ground ore is used o Greeks artificially synthesised HgS by heating mercury and sulfur together in the flask, until the mixture vaporised and reacted to form mercury(II) sulfide, which condensed at the top of the flask. The condensed vermillion was ground to produce the red pigment. Egyptian Blue, calcium copper silicate [CaO.CuO.4SiO2]: heating a mixture of sand, copper, and alkali produces a glass, which is ground to make the pigment Green-blue verdigris, basic copper (II) acetate Cu(CH3COO)2.2Cu(OH)2: Greeks produced verdigris by exposing copper to vapours of fermenting grape skins or placing copper over vinegar Sienna: Italian painters began to use the earth pigment sienna, comprising various minerals including Fe2O3.H2O, after good supplies of the substance were found in the city of Sienna Naples yellow, lead (II) antimonate [Pb3(SbO4)2]: Italians originally gathered this from slopes of Mount Vesuvius, but later produced by roasting oxides of lead and antimony DESCRIBE AN HISTORICAL EXAMPLE TO ILLUSTRATE THE RELATIONSHIP BETWEEN THE DISCOVERY OF NEW MINERAL DEPOSITS AND THE INCREASING RANGE OF PIGMENTS Chromium pigments: In 1770, French chemist Vauquelin discovered the element chromium in a Siberian gold mine, as orange lead chromate PbCrO4. Fifty years later in 1820, vast deposits of the ore chromite FeO.Cr2O3 were discovered in the United States, leading to the production of various chromium pigments. o Chrome yellow: lead (II) chromate [PbCrO4] Produced by Vauquelin: When solutions of lead (II) nitrate and sodium chromate are mixed, lead (II) chromate precipitates out o Chrome red: basic lead (II) chromate [PbCrO4.Pb(OH)2] o Chrome green: mixture of chrome yellow and Prussian blue Cadmium pigments: Stromeyer discovered cadmium in 1817, but production of cadmium pigments was delayed until the 1840s because of the metal’s scarcity. o Cadmium yellow: cadmium sulfide [CdS] Prepared by reacting an acid solution of a cadmium salt with hydrogen sulfide gas o Cadmium red: calcium sulfide selenide [CdS.CdSe] IDENTIFY DATA, GATHER AND PROCESS INFORMATION FROM SECONDARY SOURCES TO IDENTIFY AND ANALYSE THE CHEMICAL COMPOSITION OF AN IDENTIFIED RANGE OF PIGMENTS Aboriginal: Pigment/mineral Colour Chemical Composition Chemical Formula Red ochre Red Anhydrous iron (III) oxide Fe2O3 Yellow ochre Yellow Hydrated iron (III) oxides, e.g. the mineral goethite Fe2O3.H2O Brown ochre Brown Nearly pure limonite FeO(OH) Manganese (IV) oxide Black Manganese (IV) oxide MnO2 Charcoal Black Mainly graphite C Kaolinite White Hydrated aluminium silicate Al2O3.2SiO2.2H2O Chalk White Calcium carbonate CaCO3 Gypsum White Calcium sulfate dihydrate CaSO4.2H2O Egyptian/Roman: Pigment/mineral Colour Chemical Composition Chemical Formula Malachite Green Basic copper carbonate CuCO3.Cu(OH)2 Azurite Blue Basic copper carbonate 2CuCO3.Cu(OH)2 Vermillion/Cinnabar Red Mercury (II) sulfide HgS Orpiment Yellow Arsenic trisulfide As2S3 Realgar Orange-scarlet Tetraarsenic tetrasulfide As4S4 Egyptian blue/Blue frit Blue Calcium copper silicate CaO.CuO.4SiO2 White lead White Basic lead carbonate 2PbCO3.Pb(OH)2 Minium/Red lead Red Lead tetroxide PbO2.2PbO Massicot Yellow Lead (II) oxide PbO Verdigris Blue-green Basic copper (II) acetate Cu(CH3COO)2.2Cu(OH)2 Lakes: Madder lake (pink-red): alizarin [Cl4H8O4] from root of the madder plant Indigo lake (blue-violet): vegetable dye indigotin [C16H10N2O2] from the indigo plant Crimson/carmine lake (pink-red): carminic acid [C22H20O13]from female cochineal insect Italian: Natural ultramarine (blue): complex sodium aluminosilicate with sulfur [Na8-10Al6Si6O24S2-4] Naples yellow: lead (II) antimonate [Pb3(SbO4)2] Lead tin yellow: lead-tin oxide [Pb2SnO4] Italian ochres: siennas, umbers and green earths Post-18th century: Pigment/mineral Colour Chemical Composition Chemical Formula *Smalt Blue Cobalt (II) oxide CoO *Verditer Blue Basic copper carbonate 2CuCO3.Cu(OH)2 Prussian blue Blue Ferric hexacyanoferrate Fe4(Fe(CN)6)3 Scheele’s green Green Acid copper (II) arsenite CuHAsO3 Emerald green Green Copper (II) acetoarsenate Cu(CH3COO)2.3Cu(AsO2)2 *Chrome yellow Yellow Lead (II) chromate PbCrO4 *Chrome red Red Basic lead (II) chromate PbCrO4.Pb(OH)2 Cobalt blue Blue Cobalt-aluminium oxide CoO.Al2O3 Cobalt green Green Cobalt-zinc oxide CoO.nZnO Cobalt yellow Yellow Cobalt-potassium nitrite CoK3(NO2)6.H2O *Dark cobalt violet Violet Cobalt phosphate Co3(PO4)2 Light cobalt violet Violet Cobalt arsenate Co3(AsO4)2 *Cadmium yellow Yellow-orange Cadmium sulfide CdS *Cadmium red Red Calcium sulfide selenide CdS.CdSe Monastral blue Blue Copper phthalocyanine blue C32H16N8Cu *Titanium white White Titanium dioxide TiO2 *Zinc white White Zinc oxide ZnO 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 Metal components of pigments used as a source of colour are mostly from the transition metal or d-block of the Periodic Table. Electron configurations of transition metals contribute to the colours produced. Each metal, in its various compound pigments, is usually associated with a colour range. Iron, copper, arsenic, mercury, lead 2. By the twentieth century, chemists were using a range of technologies to study the spectra, leading to increased understanding about the origins of colours of different elements IDENTIFY NA+, K+, CA2+, BA2+, SR2+, AND CU2+ BY THEIR FLAME COLOUR Element Cation Flame Colour Sodium Na+ Potassium K+ Yellow Violet Calcium Ca2+ Brick red Barium Ba2+ Apple green Strontium Sr2+ Crimson Copper Cu2+ Blue-green PERFORM FIRST-HAND INVESTIGATIONS TO OBSERVE THE FLAME COLOUR OF NA+, K+, CA2+, BA2+, SR2+, AND CU2+ DESCRIBE THE DEVELOPMENT OF THE BOHR MODEL OF THE ATOM FROM THE HYDROGEN SPECTRA AND RELATE ENERGY LEVELS TO ELECTRON SHELLS It was known that when atoms of a particular element were excited (e.g. by electric discharge), the atoms emitted light. When this light was passed through a prism, it was found that it consisted of a few discrete wavelengths. Each element had a unique emission spectrum. Rutherford’s model of the atom visualised each atom as consisting of a dense positively charged nucleus surrounded by orbiting electrons, similar to the arrangements of planets in the solar system o According to classical physics, any charged particle moving on a curved path emits EM radiation, so the electrons in Rutherford’s model would lose energy and spiral into the nucleus, releasing a continuous spectrum. Therefore, Rutherford’s model could not explain atomic line spectra. In 1901, Max Planck had formulated a relationship between frequency of radiation and energy of a discrete unit of that radiation (a photon): 𝐸 = ℎ𝑓 Niels Bohr developed a model that explained the observed line spectra of hydrogen o Bohr proposed that electrons could only exist at certain allowable energy levels. Each energy level corresponded to a circular path of different radius. EXPLAIN WHAT IS MEANT BY N, THE PRINCIPAL QUANTUM NUMBER In Bohr’s model, each energy level/electron shell is assigned a principal quantum number. The energy level closest to the nucleus is given n=1, and called the ground state o Energy levels with greater energy, and greater orbital radius, are assigned larger principal quantum numbers, i.e. n = 2, n=3, etc. An electron may jump from one of the allowed energy levels to another by absorbing/emitting energy equivalent to the energy difference between the electron shells. An electron may jump to a higher energy level by absorbing the required energy from heat, incoming EM radiation, or electric discharge. When the electron falls back from a higher energy level to a lower level, the precise energy difference is released as radiation, with frequency of the radiation given by 𝐸 = ℎ𝑓. The hydrogen spectrum appears as lines, because there are discrete energy levels in the atom, and hence a limited number of possible energy transitions. SOLVE PROBLEMS AND USE AVAILABLE EVIDENCE TO DISCUSS THE MERITS AND LIMITATIONS OF THE BOHR MODEL OF THE ATOM Merits: Explained the emission of line spectra by excited atoms Explained the existence of different line spectra for different elements Predicted the observed frequencies in the hydrogen emission spectrum with reasonable precision Determined ionisation energy for the hydrogen atom, and its atomic radius Limitations: Did not explain the different intensities of these lines Did not explain why only some radii were allowed Did not explain why the electrons orbiting in stationary states did not lose energy via EM radiation Bohr’s model failed to predict the spectrum of any element other than hydrogen Did not explain the fact that single emission lines for hydrogen were actually sometimes composed of two or more closely spaced lines (no knowledge of subshells or orbitals) Could not explain the splitting of spectral lines under the influence of a magnetic field (Zeeman effect) IDENTIFY THAT, AS ELECTRONS RETURN TO LOWER ENERGY LEVELS, THEY EMIT QUANTA OF ENERGY WHICH HUMANS MAY DETECT AS A SPECIFIC COLOUR An electron may absorb energy to jump to higher energy levels. When the electron returns back to lower energy levels, it emits photons with frequency corresponding to the energy drop. For some of these emissions, the energy released in the visible light part of the EM spectrum, allowing humans to detect it as a specific colour o e.g. excited hydrogen atoms emit violet light comprised of the wavelengths shown in the diagram EXPLAIN THE FLAME COLOUR IN TERMS OF ELECTRONS RELEASING ENERGY AS THEY MOVE TO A LOWER ENERGY LEVEL When a granule of an ionic compound or a drop of its solution is placed into a flame, heat energy from the flame is absorbed by electrons, which jump up to higher energy levels. When the electrons drop back to lower energy levels, they emit photons corresponding to the energy difference between the two orbits. Each element has a unique set of possible energy transitions, and hence a unique emission spectrum. Some elements produce a very intense spectral line (or several closely spaced ones) in the visible range that serves as a marker for the elements’ presence. Therefore, the colour of the flame is created by the specific emission spectrum of the element. EXPLAIN WHY EXCITED ATOMS ONLY EMIT CERTAIN FREQUENCIES OF RADIATION When an excited electron drops from a higher energy level to a lower energy level, it emits a photon with frequency corresponding to the energy difference between the 2 electron shells by 𝐸 = ℎ𝑓 o A greater drop in radius corresponds to a greater energy difference, and will cause the emission of a photon of greater frequency (i.e. closer to the violet end of the spectrum) Therefore, the frequencies of radiation emitted will depend on the possible energy transitions within the atom’s energy levels. The possible energy transitions are unique to each element because each element has a unique set of possible energy levels. Hence, excited atoms only emit certain frequencies of radiation specific to that element DISTINGUISH BETWEEN THE TERMS SPECTRAL LINE, EMISSION SPECTRUM, ABSORPTION SPECTRUM AND REFLECTANCE SPECTRUM Spectral Line: Light released by an atom which has been excited (by EM radiation, electric discharge, or heat), is not a continuous distribution of wavelengths, but rather consists of a few wavelengths unique to the element When the emitted light is passed through a narrow slit, and then through a prism/diffraction grating to disperse the light into separate wavelengths, we can see the wavelengths present as discrete spectral lines, which are separated by blank areas Emission Spectrum: It was found that each element produces a unique line spectrum when excited. From this information, it became possible to identify the elements present in minerals. Each element’s unique line spectrum is called its emission spectrum Absorption Spectrum: White light (continuous spectrum) is passed through a substance in vapour phase. Atoms of the gas only absorb photons corresponding to quanta of energy involved in possible transitions of electrons from ground state to higher energy. Hence, the wavelengths absorbed by the gas are those in its emission spectrum. As excited electrons drop back to their ground state, they re-emit the absorbed photons in all directions. By passing a continuous spectrum of light through the cool gas, the absorbed wavelengths have been reduced in intensity. When this light is directed through a spectroscope or prism, the absorption spectrum produced will have dark lines at the absorbed wavelengths, across a continuous spectrum of colour The absorption spectrum of an element is the complement of its emission spectrum Reflectance Spectrum: Atoms and molecules reflect energy at wavelengths related to their atomic structures. A mineral can be identified by shining white light on it, and examining the spectrum of light that is reflected. The reflectance spectrum is the complement of the molecular absorption spectrum, and represents the visible colour of an object *??: When a photon strikes a molecule, it normally imparts the energy as heat and the photon is absorbed. Molecules have specific resonant frequencies at which they can rotate or vibrate, each corresponding to discrete energy levels (vibrational modes). When an incoming photon has energy equal to the resonance energy of the molecule, the photon is absorbed and the molecule is put into an excited state, i.e. its pattern of vibration changes. The molecule than relaxes back to its ground state by emitting a photon with energy/frequency equal to that of the original photon. The photon appears to have been reflected. 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 When scientists studied the line spectra of many atoms, they noticed that lines originally thought to be singlets were actually closely spaced doublets or triplets. In some respects, electrons behave as if they are spinning on an axis (clockwise or anti-clockwise). By Pauli’s Exclusion Principle, each orbital holds two electrons that must have opposite spins. Within the 3p orbital, one electron has angular momentum j=3/2 and the other electron has j=1/2. Spin-orbit splitting: The magnetic energy due to electron spins interacts with the internal magnetic field caused by the electrons’ orbital motion. This causes spin-orbit splitting, i.e. the energy levels move slightly as a result of electron spin, splitting the 3p level into two states. In effect, atomic energy levels depend partially on angular momentums. The emission spectrum of sodium is dominated a closely spaced pair of yellow lines (a doublet), together called the D-lines. The next strongest visible line has an intensity about 0.7% that of the strongest line. The line at 589.0 nm has twice the intensity of the line at 589.6 nm. They are together called the D-lines. Both of the D-lines represent the transition of electrons within the n=3 energy shell, from 3p energy level to 3s energy level. The transition from 3p3/2 to 3s causes the 589.0nm line. The transition from 3p1/2 to 3s causes the 589.6nm line. Therefore, the opposite spins in each orbital result in the hyperfine splitting of spectral lines. Zeeman effect: When an external magnetic field is applied, these energy levels and spectral lines are further split apart OUTLINE THE USE OF INFRA-RED AND ULTRA-VIOLET LIGHT IN THE ANALYSIS AND IDENTIFICATION OF PIGMENTS AND THEIR CHEMICAL COMPOSITION Double-beam spectrophotometers are used to record infra-red and ultra-violet absorption spectra o A diffraction grating splits the continuous spectrum of the light source into a band of wavelengths. Then, a narrow slit only lets a narrow range of wavelengths through. o A beam splitter splits the light beam into two beams. One beam is passed through the sample. The other beam is passed through the solvent alone. o A detector measures radiation passing through sample and solvent. A comparison of the intensity of the 2 beams allows absorption to be determined UV-visible Spectroscopy Based on excitation of electrons, which are promoted to higher energy levels Source is a tungsten lamp (for visible part) and a deuterium discharge tube (for UV part), and the detector is a photomultiplier tube Pigment needs to be in solution Mainly used to identify pigments containing metal ions Infra-red Spectroscopy: Based on molecular vibrations/stretching due to absorption of IR radiation Source is a heated ceramic such as silicon carbide, and the detector is a thermocouple which measures temperature. Glass cell which contains the sample is made of KBr which does not absorb IR. Molecules must be polar Mainly used to identify organic molecules, so its good for older paintings Infra-red Reflectography: Infra-red reflectography is a non-destructive technique to detect the artist’s preliminary drawings A lamp emitting near infra-red is shone onto the surface of a painting, and the reflected light is detected Infra-red penetrates pigments such as titanium dioxide, cadmium red or vermillion, reflects from the white ground, and is detected by a thermocouple o The underdrawing is usually marked out with carbon-based substances such as graphite, charcoal or black ink. These substances absorb infra-red radiation and do not reflect it back, producing a black image on the computer readout. o Infra-red reflectography can also be used to identify many copper-containing green pigments as these also absorb infra-red Ultra-violet Reflectance Spectroscopy: Compares reflected radiation from a pigment surface with reflection from a material that does not absorb UV radiation such as SiO2 o Note: Ultraviolet light can cause some substances to fluoresce (absorb and re-emit light at different frequencies), so pigments with these chemicals can be identified using UV light. 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 Absorption and Reflectance Spectra: Absorption spectra are measured by passing light through a solution of the substance. Absorbance is directly proportional to concentration. Therefore, the concentration of a chemical pigment in a paint can be determined. A reflectance spectrum is used when the substance cannot be dissolved in a colourless solvent. To obtain a reflectance spectrum, white light is shone onto the painting’s surface, and the spectrum which is reflected instead of absorbed, is measured. o For analysing paintings, absorption spectroscopy is destructive while reflectance spectroscopy is nondestructive. Absorbance is the % of light absorbed, while reflectance is the % of light reflected. The reflectance spectrum is the complement of the absorption spectrum. For both of these methods, wavelength is plotted against intensity of absorbance/reflectance, because colours are a result of wavelength. The intensity of absorption/reflectance at certain wavelengths tells us the concentration of pigments, and the shape of the curve relates to the purity of the pigment. o Absorbance and reflectance for a blue pigment are shown below Effect on Zinc Oxide and Copper Pigments: Infra-red (heat) changes the colour of zinc white [ZnO] from white to yellow, but the pigment returns to its normal white colour on cooling Infra-red radiation permanently breaks down red copper (I) oxide, green malachite, and verdigris into black copper (II) oxide Zinc white absorbs ultraviolet radiation and fluoresces pale yellow Malachite absorbs ultraviolet radiation and fluoresces a dirty mauve colour 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 PROVIDE EXAMPLES OF THE TECHNOLOGY’S USE Laser Microspectral Analysis: Methodology: A pulse of laser light vaporises a small sample from the surface. The vapour is fed through a spark gap between 2 electrodes, which excites the vapour particles. When electrons in the excited atoms/ions return to their ground state, they release photons producing an emission spectrum characteristic of the elements present. The radiation released is focused to a spectrophotometer, which tells us the emission spectrum, i.e. intensity of emission as a function of wavelength Comparing the spectrum produced with thousands of known pigment spectra tells us the pigments that are present on the surface of the painting Identification of Elements: This is a destructive technique, but only a minute/negligible amount of the sample is destroyed Can identify more than one element at a time, unlike AAS This technology is important in identifying trace elements in solid and liquid samples because it is highly sensitive, minimally destructive, and requires minimal sample preparation Uses: Laser microspectral analysis is used to analyse the pigments present by their component elements, which tells us if an art work is authentic In restoring paintings, laser microspectral analysis tells us the elemental composition of pigments originally used by the artist. From this information, a synthetic substitute for this pigment can be prepared for restoration work. 3. The distribution of electrons within elements can be related to their position in the Periodic Table DEFINE THE PAULI EXCLUSION PRINCIPLE TO IDENTIFY THE POSITION OF ELECTRONS AROUND AN ATOM The Pauli Exclusion Principle states that no two identical fermions can occupy the same quantum state simultaneously. As applied to electrons, it states that no two electrons in the same atom can have the same set of four quantum numbers: n (energy shell), l (sub-shell), ml (orbital), and ms (spin). o The Pauli Exclusion Principle underlies the partitioning of energy levels in atoms into discrete energy shells, sub-shells and orbitals, thus explaining why matter occupies space o A consequence of this is that an orbital can hold a maximum of 2 electrons that must have opposite spins IDENTIFY THAT EACH ORBITAL CAN CONTAIN ONLY TWO ELECTRONS Heisenberg’s Uncertainty Principle tells us that we cannot know the precise position and momentum of an electron at the same time. Hence, we describe the positioning of electrons as 3-dimensional probability clouds which we call orbitals. A higher dot density at a point represents greater probability of finding the electron at that position. The outer ‘boundary’ of the orbital shape is decided such that an electron has a high probability of being found at any point inside the region, and a very low probability of being found outside it. The probability cut-off to determine the orbital’s boundary is arbitrarily chosen. Because there are only 2 possible values for electron spin (-½ and ½), and by the Pauli Exclusion Principle all electrons in an orbital must have different spin, an orbital can contain at maximum two electrons Therefore, an orbital can hold zero, one or two electrons. If two electrons are present, they must have opposing spins. DEFINE THE TERM SUB-SHELL The azimuthal quantum number, l, refers to the sub-shell. Sub-shells are energy sublevels within energy shells with slightly different energy levels. There are 4 types of sub-shells: s,p,d,f (in increasing order of energy). The shapes of orbitals comprising each sub-shell are such that together they fill most of the space available, since electrons naturally repel each other. o The ‘s’ sub-shell is comprised of 1 orbital (s). The s-orbital is spherically symmetrical around the nucleus. It has 2 electrons when full. o The ‘p’ sub-shell is comprised of 3 orbitals (pX, pY, pZ). Each of these orbitals is dumb-bell shaped, with the nucleus between the two halves of the dumb-bells. It has 6 electrons when full. o The ‘d’ sub-shell is comprised of 5 orbitals (dXY, dXZ, dYZ, dX^2-Y^2, dZ^2). It has 10 electrons when full. o The ‘f’ sub-shell is comprised of 7 complex orbitals. It has 14 electrons when full. The number of sub-shells in any energy shell is n, the same as the principle quantum number. The number of orbitals in any energy shell is given by n2, and the maximum number of electrons in any energy shell is given by 2n2 o The n=1 energy shell has 1 sub-shell (s), comprised of 1 orbital, or 2 electrons o The n=2 energy shell has 2 sub-shells (s, p), comprised of 1 + 3 = 4 orbitals, or 8 electrons o The n=3 energy shell has 3 sub-shells (s, p, d), comprised of 1 + 3 + 5 = 9 orbitals, or 18 electrons o The n=4 energy shell has 4 sub-shells (s, p, d, f), comprised of 1 + 3 + 5 + 7 = 16 orbitals, or 32 electrons OUTLINE THE ORDER OF FILLING OF SUB-SHELLS Aufbau Principle: Orbitals in subshells with the lowest energy are filled first o Energy shells with lower principle quantum numbers generally have lower energies, and s < p < d < f within any energy shell. o However, the 4s sub-shell has less energy than the 3d subshell. This means that after 3p is filled, electrons go into 4s before 3d, even though 4s is in a higher energy shell. Hund’s Rule of Maximum Multiplicity: If 2 or more orbitals in the same sub-shell are available, one electron goes into each orbital keeping spins parallel until all orbitals in the sub-shell are half-full. The next electron is forced to add to a half-filled orbital, and is given opposite spin to the electron already in it. o Filling orbitals this way reduces screening of electron-nucleus attractions since electrons are further apart, and thereby increases stability. http://employees.oneonta.edu/viningwj/sims/atomic_electron_configurations_s1.html IDENTIFY THAT ELECTRONS IN THEIR GROUND-STATE ELECTRON CONFIGURATIONS OCCUPY THE LOWEST ENERGY SHELLS, SUB-SHELLS AND ORBITALS AVAILABLE TO THEM AND EXPLAIN WHY THEY ARE ABLE TO JUMP TO HIGHER ENERGY LEVELS WHEN EXCITED In their ground state, electrons occupy the lowest energy shells, sub-shells and orbitals available to them, as per the Aufbau Principle Electrons can jump to higher energy levels (different sub-shells or energy shells) by absorbing the energy difference (as electrical, heat or EMR) when they are excited 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 The electron configuration of atoms may be denoted with a standard notation, or using orbital diagrams o Standard notation for oxygen in its ground state: 1s22s22p4 or [He] 2s22p4 o Orbital diagram for oxygen in its ground state: see image o Standard notation for oxygen in an excited state: 1s22s22p33s1 o Orbital diagram for oxygen in an excited state: see image o Standard notation for sodium in its ground state: 1s22s22p63s1 or [Ne] 3s1 In standard notation, coefficients represent energy shell, letters represent sub-shell, and indices represent the number of electrons in each sub-shell In orbital diagrams, the up and down arrows represent electrons with opposite spins EXPLAIN THE RELATIONSHIP BETWEEN THE ELEMENTS WITH OUTERMOST ELECTRONS ASSIGNED TO S, P, D AND F BLOCKS AND THE ORGANISATION OF THE PERIODIC TABLE The sub-shell of the outermost electrons in an element affects its properties. More significantly, elements with the same number of valence electrons in these common outer sub-shells display similar chemical properties. The Periodic Table is organised into 4 blocks such that elements with the outermost electrons assigned to the same sub-shell type (s, p, d or f) are grouped together. Elements with the same number of valence electrons are positioned in the same vertical group. The s-block includes Groups 1 and 2. Group 1 elements have the valence configuration s1. Group 2 elements have a completely filled ‘s’ sub-shell. The p-block includes Groups 3-8. Group 3 elements have the valence configuration s2p1. Group 5 elements (e.g. nitrogen) have 3 half-filled p-orbitals. Group 8 noble gases have the valence configuration s2p6 o For the s and p blocks, period number corresponds to principle quantum number The d-block includes the transition metals, as the outermost electrons are being added to the ‘d’ sub-shell o Because the first ‘d’ sub-shell (3d) gets filled after 4s, the d-block begins in period 4 The f-block includes the lanthanide and actinide metals, as the outermost electrons are being added to the ‘f’ sub-shell 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. It is expressed on a relative scale with fluorine assigned the highest electronegativity of 4.0 Noble gases are not assigned electronegativities because they tend not to form molecules with other atoms since their valence shell is filled Going from left to right across a period, electronegativity increases as more electrons are added. o This is because increasing the number of protons increases the electrostatic attraction between the positive nucleus and electrons. Hence, non-metals in the p-block tend to attract electrons and be reduced to a noble gas electron configuration, while metals in the s-block tend to give up electrons to achieve a noble gas configuration (they are said to be electropositive) Going down a group, electronegativity decreases as more energy shells are added. o This is because the increased distance between valence electrons and the positive nucleus reduces the attractive force, and the inner electrons provide a shielding effect against the nuclear attraction. PROCESS INFORMATION FROM SECONDARY SOURCES TO ANALYSE INFORMATION ABOUT THE RELATIONSHIP BETWEEN IONISATION ENERGIES AND THE ORBITALS OF ELECTRONS The first ionisation energy, I1, of an atom or ion is the amount of energy needed to remove the outermost electron (in the highest occupied sub-shell) from one mole of the gaseous atoms or ions o Na(g) Na+(g) + eI1 = ∆H = 502 kJ mol-1 As we go from left to right across a period of the Periodic Table, electrons being added go into the same shell and are about the same distance from the nucleus, but the positive nuclear charge increases. The attraction between the outermost electron and the positive nucleus increases, so first ionisation energy increases across periods of the Periodic Table. o This explains why the graph of first ionisation energy against atomic number rises to a peak at noble gases (a lot of energy is needed to remove an electron from the stable octet of a noble gas), then drops sharply at the following Group 1 metal (the single electron in the outermost ‘s’ sub-shell is easily removed) o Halogens with high ionisation energies tend to gain electrons, while alkali metals with low ionisation energies easily lose electrons to achieve noble gas configuration As atomic size increases down a Periodic table group, valence electrons are further from the positive nucleus, and increasingly shielded by inner electrons, so the force of attraction decreases. Hence, the first ionisation energy decreases down groups of the Periodic Table. o This explains why subsequent peaks on the graph have smaller and smaller first ionisation energies To explain the local minima/maxima within each period, we must look at how orbitals are filled within each shell o Group 3 elements have a lower ionisation energy than expected. This is because the last electron is in the ‘p’ sub-shell, which has a higher energy than the 2s sub-shell in the case of boron, or 3d sub-shell in the case of gallium. Also, the last electron is well-shielded from the increased nuclear charge by the full 2s or 3d sub-shell E.g. boron: 1s22s22p1 1s22s2 + electron o Group 6 elements have a lower ionisation energy than expected. This is because a Group 5 element has all three of its outer p-orbitals half-filled, whereas a Group 6 element has the last electron paired up in the same p-orbital as another electron. Because of the electrostatic repulsion between electrons in the same orbital, less energy is needed to remove the extra electron. E.g. oxygen: 1s22s22pX12pY12pZ2 1s22s22pX12pY12pZ1 + electron 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 The second ionisation energy is the energy required to remove the second electron from one mole of the gaseous atom or ion. o Na+(g) Na2+(g) + eI2 = ∆H = 4569 kJ mol-1 In general, the nth ionisation energy is the energy required to remove the nth electron after the first n-1 electrons have been removed. Electrons are removed from their orbitals in the reverse order to that in which they were filled. Successive ionisation energies are always greater as negative electrons are being removed from an increasingly positive ion. Moreover, removing an electron in a shell that is closer to the nucleus requires a significantly greater ionisation energy. Therefore, if the second ionisation energy is much greater than the first ionisation energy, this indicates that the atom has 1 electron in the outermost shell, as it requires much less energy to remove a lone electron from a shell that to remove an electron from the full shell of a noble gas configuration In general, if the nth ionisation energy is abnormally high after a steady progressive increase in ionisation energies, this indicates that the atom has (n-1) electrons in the outermost shell. For instance, aluminium’s ionisation energies are: I1 = 584, I2 = 1823, I3 = 2751, I4 = 11 584, I5 = 14 837 o Because there is a steep rise in ionisation energy after removing 3 electrons, we can conclude that aluminium has 3 valence electrons. If we know that aluminium has 3 energy shells, we can also conclude that the sub-shells occupied by these outer electrons are s2 and p1, since this is the order that sub-shells are filled in. o Aluminium’s electron configuration is 1s22s22p63s23p1 4. The chemical properties of the transition metals can be explained by their more complicated electronic configurations IDENTIFY THE BLOCK OCCUPIED BY THE TRANSITION METALS IN THE PERIODIC TABLE The transition metals occupy the ‘middle’ d-block of the periodic table DEFINE THE TERM TRANSITION ELEMENT Transition elements are those who atom has a partially filled ‘d’ sub-shell, or which can form at least one ion with a partially filled ‘d’ sub-shell o Group 12 elements such as zinc are not transition metals, because their ‘d’ sub-shells are completely filled, and ions such as Zn2+ generally retain the complete ‘d’ sub-shell o Scandium and Yttrium in Group 3 are generally not considered transition metals even though they have 1 electron in the ‘d’ sub-shell because all their compounds contain the ions Sc3+ or Y3+ PROCESS AND PRESENT INFORMATION FROM SECONDARY SOURCES BY WRITING ELECTRON CONFIGURATIONS OF THE FIRST TRANSITION SERIES IN TERMS OF SUB-SHELLS Elements of the first transition series Electron Configuration scandium [1s22s22p63s23p 6] 3d14s2 titanium [1s22s22p63s23p 6] 3d24s2 vanadium [1s22s22p63s23p 6] 3d34s2 chromium [1s22s22p63s23p 6] 3d54s1 manganese [1s22s22p63s23p 6] 3d54s2 iron [1s22s22p63s23p 6] 3d64s2 cobalt [1s22s22p63s23p 6] 3d74s2 nickel [1s22s22p63s23p 6] 3d84s2 copper [1s22s22p63s23p 6] 3d104s1 zinc [1s22s22p63s23p 6] 3d104s2 The 3d sub-shell fills as per Hund’s rule, with one electron adding to each of the 5 d-orbitals, before a second electron is paired up with any of them. There are 2 exceptions to this: chromium and copper. These exceptions arise because the 3d and 4s sub-shells have a very small energy difference. o Chromium is [Ar] 3d54s1 instead of the expected 3d44s2 because a more stable electronic configuration with lower energy is achieved when all the 3d orbitals are exactly half-filled o Copper is [Ar] 3d104s1 instead of the expected 3d94s2 because a more stable configuration is achieved when the 3d sub-shell is completely filled EXPLAIN WHY TRANSITION METALS MAY HAVE MORE THAN ONE OXIDATION STATE s-block and p-block elements usually have oxidation states equal to their Group number, but transition metals exhibit a variety of oxidation states This is because the ‘s’ and ‘d’ sub-shells are very close in energy, so valence electrons can be lost from the ‘s’ and/or ‘d’ sub-shells in transition metals The common +2 oxidation state is formed by the loss of two ‘s’ electrons. Oxidation states above +2 involve the additional loss of ‘d’ electrons. The maximum oxidation state of the first transition series is the total number of 4s and 3d electrons, i.e. all the electrons in the outermost energy shell. However, beyond manganese [Ar] 3d54s2, maximum oxidation states decrease to +3 and +2. This is because transition metals in high oxidation states tend to be easily reduced. 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 Because 4s and 3d sub-shells are of similar energy, it does not require much energy to oxidise or reduce a transition metal. Strong oxidising agents have high positive Eθ values. Transition metal ions in which the metal has a high oxidation state tend to be strong oxidising agents. For example, in complex ions such as dichromate Cr2O72- and permanganate MnO4-, the transition metal ions have high oxidation states (+6 in dichromate, +7 in permanganate), and have smaller atomic radii as many electrons have been removed. This smaller radius gives a greater attraction for electrons, so that such species will be strong oxidising agents (electron acceptors). Further, because these transition elements are combined with highly electronegative oxygen and fluorine, they can release oxygen to oxidise substances by accepting electrons. Cr2O72- and MnO4- are widely used as oxidants in analytical chemistry SOLVE PROBLEMS AND PROCESS INFORMATION FROM SECONDARY SOURCES TO WRITE HALF-EQUATIONS AND ACCOUNT FOR THE CHANGES IN OXIDATION STATE Change in Oxidation Number Half equation chromium: +VI to +III Cr2O72-+ 14H+ + 6e- 2Cr3+ + 7H2O manganese: +VII to +II MnO4- + 8H+ + 5e- Mn2+ + 4H2O iron: +III to +II Fe3+ + e- Fe2+ copper: +II to +I Cu2+ + e- Cu1+ CHOOSE EQUIPMENT, PERFORM A FIRST-HAND INVESTIGATION TO DEMONSTRATE AND GATHER FIRST-HAND INFORMATION ABOUT THE OXIDISING STRENGTH OF KMNO4 ACCOUNT FOR COLOUR CHANGES IN TRANSITION METAL IONS IN TERMS OF CHANGING OXIDATION STATES The small energy differences between the d-orbitals in transition metals are similar to the energies of photons of visible light. Thus, d-orbital electrons in transition metals can absorb visible light photons and move to slightly higher d-orbitals. Because each ion of a transition metal in different oxidation state has a different arrangement of filled/unfilled 3d orbitals, each ion has a different set of possible electron transitions, and hence a different set of wavelengths of light that can be absorbed. The absorption of some components of white light means that complementary colours are reflected/transmitted, e.g. a transition metal ion that absorbs red wavelengths will appear blue-green. This is why transition elements which are changing oxidation states in redox reactions have colour changes. E.g. Cu (II) compounds are usually blue because Cu2+ has a 3d9 sub-shell configuration, and colour results from electron transitions from a full d-orbital to the half-filled d-orbital. Cu (I) compounds are not especially coloured because Cu+ has a full 3d10 sub-shell configuration Small amounts of transition metal impurities cause the colour of various gemstones, e.g. red Cr3+ in rubies PERFORM A FIRST-HAND INVESTIGATION TO OBSERVE THE COLOUR CHANGES OF A NAMED TRANSITION ELEMENT AS IT CHANGES IN OXIDATION STATE 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 Ions dissolved in water become ‘hydrated’, i.e. they become surrounded by water molecules. The water molecules orient their negatively charged sides (O atoms) towards the ion in the case of a cation, or their positively charged ions towards it in the case of an anion A hydrated ion is an ion in which a specific number of water molecules is associated with each formula unit of the complex ion DESCRIBE HYDRATED IONS AS EXAMPLES OF A COORDINATION COMPLEX OR A COMPLEX ION AND IDENTIFY EXAMPLES Hydrated ions are examples of complex ions. In a complex ion, the ligand anions/molecules donate electrons to the metal cation, forming coordinate covalent bonds. In hydrated ions, the ligands are water molecules To maintain electrical neutrality, these complex ions may pair up with simple ions. Compounds containing at least one complex ion are called coordination compounds Examples of hydrated coordination complexes / complex ions: o Hexaaquamagnesium (II): [Mg(H2O)6]2+ (image 1) 2+ o Tetraaquacopper (II): [Cu(H2O)4] (image 2) + o Diaquasilver (I): [Ag(H2O)2] Examples of hydrated coordination compounds: o Magnesium sulfate heptahydrate: MgSO4.7H2O (image 3) o Calcium sulfate dihydrate: CaSO4.2H2O Examples of non-hydrated coordination complexes: o Tetraaminecopper (II): [Cu(NH3)4]2+ DESCRIBE MOLECULES OR IONS ATTACHED TO A METAL ION IN A COMPLEX ION AS LIGANDS The molecules or anions attached to a central metal cation in a complex are called ligands. Coordination number is the number of ligands attached to the central cation, e.g. Tetraaminecopper (II) has a coordination number of 4 EXPLAIN THAT LIGANDS HAVE AT LEAST ONE ATOM WITH A LONE PAIR OF ELECTRONS Ligands attach to the metal cation with coordinate covalent bonds, by donating electrons which are accepted by the cation. Therefore, the metal acts as a Lewis acid while the ligand acts as a Lewis base, and the formation of a coordinate covalent bond between them is a Lewis acid-Lewis base reaction o Lewis acids are substances that accept an electron pair o Lewis bases are substances that donate an electron pair In order to act as Lewis bases and donate an electron pair, ligands must have at least one atom with a lone pair of electrons IDENTIFY EXAMPLES OF CHELATED LIGANDS Monodentate ligands have exactly one atom with an unshared pair of electrons, e.g. H2O, NH3, Cl Bidentate ligands have exactly two atoms with unshared pairs of electrons, e.g. oxalate C2O4 In general, polydentate ligands can form coordinate covalent bonds by donating lone electron pairs from two or more atoms. Polydentate ligands are also known as chelating agents (chele means ‘claw’ in Greek), and generally form more stable complexes than monodentate ligands. Chelation is defined as the formation of bonds between two or more separate binding sites (atoms) within the same ligand, and a single central atom For example, ethylenediaminetetraacetic acid (EDTA) is a hexadentate ligand, so it forms stable complexes o Hard water is treated by adding EDTA, which removes Ca2+ ions from solution by forming stable harmless complexes o EDTA is used to treat lead poisoning, as it sequesters lead (II) ions as stable [Pb EDTA]2-, which is excreted via the kidney o Iron chelate, an iron (II)-EDTA complex, is used to treat iron-deficient soils. This is because the complex remains insoluble in alkaline soils, while other iron salts will precipitate out iron hydroxide. 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 See images in above DotPoints DISCUSS THE IMPORTANCE OF MODELS IN DEVELOPING AN UNDERSTANDING OF THE NATURE OF LIGANDS AND CHELATED LIGANDS, USING SPECIFIC EXAMPLES We can observe the properties of chelated ligands by performing experiments, but must rely on models for explain these properties in terms of chemical structure. o We use 2-dimensional drawings such as Lewis dot diagrams, 3-dimensional computer images, and molecular model kits, to indicate the geometry of complex ions Models are important in explaining: o How ligands join with cations to form coordination complexes o How electron pairs from ligands fill up cation orbitals, and thereby form covalent bonds o How the number/type of metal ion’s orbitals, and number/positioning of ligand’s orbitals, determine the ratio of ligands to cations o Why electrostatic repulsions between electrons in ligands result in certain geometrical shapes o The mechanisms by which some ligands can chelate Models do not show: o Electrons and their transitions in orbitals o The energy changes involved and hence the stability of the complex ion o What initiates cation-ligand interactions o The mobility/flexibility of the bonds formed Example: Haemoglobin o The haemoglobin molecule consists of four haem groups. Each haem group is a planar tetradentate ligand that uses lone pairs of electrons on each of its four N atoms to bond to the iron (II) ion. A fifth nitrogen atom from the protein bonds to the Fe2+ ion from below, and an oxygen molecule can bond to the iron Fe2+ from above the chelate ring 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 The energies of d-orbitals are affected by the metal ion, its oxidation state, and surrounding ligand groups These factors affect the energy separations between d-orbitals, and hence the frequencies of photons that can be absorbed to promote electrons to higher d-orbitals. Therefore, the visible colour of an ion complex depends on the metal ion, its oxidation state, and the ligand Effect of different oxidation states and ligands Iron Complex Hexaaquairon (II) Hexaaquairon (III) Hexacyanoiron (III) Formula [Fe(H2O)6]2+ [Fe(H2O)6]3+ [Fe(CN)6]3- Colour Green Yellow Red Effect of different ligands Nickel Complex Hexaaquanickel (II) Hexaamminenickel (II) Triethylenediaminenickel (II) Ligand H2O H2O CN- Formula [Ni(H2O)6]2+ [Ni(NH3)6]2+ [Ni(en)3]2+ Ligand H2O NH3 en Colour Green Blue Violet Chromium Ion Complexes Chromium Complex Hexaaquachromium (II) Hexaaquachromium (III) Pentaaquachromium (III) chloride Chromium (III) tetrahydroxide Formula [Cr(H2O)6]2+ [Cr(H2O)6]3+ [Cr(H2O)5Cl]2+ [Cr(OH)4]- Ligand H2O H2O H2O OH Colour Blue Violet Green Deep green
