chemistry hsc full notes best notes

Chemistry HSC

Note: This is not the final revision of my notes (I’m constantly revising them as I do papers), and there may be a few areas of error or unclear explanations. However, I’ve gone through it a number of times, and it should be mostly very accurate and comprehensive. If you find anything wrong, it would be nice if you could tell me on [email protected] so I can either discuss it or change it. Good luck for the HSC guys

Organic Chemistry – the study of compounds containing carbonTHIS IS BACKGROUND INFOOrganic chemistry is separate because we can look at all of the included chemical groups in a unifying way, through the bonding properties of carbon. Study of major groups:

o Oxygen-containing compounds e.g. alcoholso Hydrocarbons e.g. petroleumo Carbohydrates e.g. sugarso Nitrogen-containing compounds e.g. amino acids proteins

Etc.

HydrocarbonsWhen all bonds are single, they are called alkanes. This is a family of compounds, represented by a general formula CnH2n+2, aka a homologous series. They have similar properties and reactions. There are ‘straight’ chain alkanes.e.g.

The 109o, zig-zag bonding shape is due to the tetrahedral nature of single bonds. Carbon atoms always form 4 bonds. If they don’t you’re doing something wrong.

Branched chain (one is attached to at least 3)

Methane, CH4, ethane, C2H6, Propane, C3H8, and Butane, C4H10 are all alkanes.

Physical PropertiesC1 to C4 are gases at room temp, C5 to C18 are colourless liquids, others are solids.The density of alkanes are significantly less than water (1.00g/mL), are non-conductors of electricity and are insoluble in water. The reason for their insolubility is that C-C bonds are non-polar, and C-H bonds are only slightly. This slight polarity is cancelled by symmetry in structure. Weak dispersion forces, relatively low boiling/melting points. Boiling/melting points increase as molecular weight increases, due to stronger dispersion forces (more electrons). Volatility decreases as molecular weight increases.

Alkenes

Chem notes David Lee BHHS 2007 1

C

C

C

C

C

C

C

C

C

C

C

C

mailto:[email protected]

Contain a double bond between a pair of carbon atoms. Homologous series, formula CnH2n, planar shape. There are different ways of representing structure:

(2)Full Structural Formula – shows planar geometry around double bond, and tetrahedral around other carbon atom(3)Intermediate type – infers tetrahedral shape(4)Condensed structural formula – no attempt to show structure, but enough information is provided

Isomers are different compounds with the same molecular formula but different structural formula. The double bond can be at different positions in the compound.e.g.

Physical PropertiesStraight-chain alkenes similar to alkanes. Densities similar to corresponding alkanes, insoluble in water.

AlkynesContain a triple bond between carbons. CnHn-2. As with alkenes, isomers are possible. They are non-polar, low boiling points and insoluble in water

Naming Alkanes, Alkenes and Alkyneso Stem telling length of carbon chain

C1 meth- C4 but- C7 hept-C2 eth- C5 pent- C8 oct-C3 prop- C6 hex-

o Look at the longest possible chain, then pick a prefixo Look for branches, and use a number to denote their position, starting from the closest

end e.g. 2,3 – dimethylpentane or 2 - dimethylpentaneo If double or triple bonds present, set this as priority (start counting closest to the bond)

and first state branches then double/triple bond e.g. 2– methyl – 1 – propene (methyl on second branch and double bond on first)

o If compound is cyclic, add a cyclo- before the name of the main branch e.g. 1,2,3 - trimethylcyclohexane

Saturated and unsaturated compoundsAlkenes and Alkynes – unsaturated, possible to attach more hydrogenAlkanes – saturated, max no. of H atoms that skeleton can hold

Functional GroupThe functional group of carbon compounds is the most reactive area of the compound. In alkenes and alkynes, the double/triple bonds are the functional groups. When a hydrogen atom is replaced with a halogen atom, e.g. OH, the halogen becomes the functional group.

Molecules with a particular functional group react similarly, regardless of the attached chains.


Alkanols are alkanes with one H replaced by an OH group. They are named with the ‘e’ replaced by an ‘ol’, and a prefix number to denote the position of the hydroxyl group. This group is the functional group, and provides high melting/boiling points due to polar bonds. Primary alcohols have one carbon bound to the carbon w/ OH group, secondary have two and tertiary have three. Extent of hydrogen bonding depends on exposure of OH group, most exposed in primary, highest boiling/melting points etc.

Construct word and balanced formulae equations of chemical reactions as they are encountered

Types of Organic reactionsSubstitution – replacement of one atom or group by anotherAddition – adding atoms or groups of atoms to alkenes or alkynes (bond breaks, new atoms are added on)Elimination – a small molecule breaks off and a double bond is formed in the original (reverse of addition)Condensation – two molecules react, forming a new compound and a small molecule (usually water)Hydrolysis – the action of water on a molecule results in two new products

Identify the industrial source of ethylene from the cracking of the fractions from the refining of petroleum

Ethylene is produced from natural gas or crude oil (mixtures of hydrocarbons, containing mainly alkanes and cycloalkanes and smaller amounts of unsaturated including alkenes), which is called feedstock. The feedstock is refined by fractional distillation to obtain alkenes since alkanes are susceptible to combustion and unreactive (not useful as starting material).

Ethylene is the most versatile, but not found in large quantities in feedstock. Produced from other hydrocarbons in ‘cracking’ (a process where hydrocarbons of higher mol mass are converted to lower mol mass via breaking of chemical bonds). There is greater demand for some fractions than others (e.g. gasoline > heavier hydrocarbons), and fractions from crude oil are not in optimum ratios, hence cracking. Note that air needs to be excluded to prevent combustion. Ethylene is simple and can be synthesised from many different hydrocarbons. Three ways:

1. Thermal cracking – requires very high temps and generally not used. End products hard to control since many places where bonds could break, early method. Accelerates reaction and drives equilibrium to reactants.

2. Catalytic cracking of fractions separated from petroleum. – material is passed over a catalyst at a temperature of about 500oC, and the particles adsorb onto the catalyst and have their bonds weakened, resulting in decomposition. E.g. C10H22(g) -> C8H18(g) + C2H4(g). Alkane splits further into smaller alkenes until propene/ethylene formed. Catalysts allow it to be carried out at lower temperatures. Zeolite (by mid 1970’s) is the main catalyst, and is a crystalline substance of Al, Si and O. Usually fine powder (higher surface area for action of catalyst) circulated through feedstock. Zeolite gives greater control over products under different conditions of temperature and pressure (thus increasing yields of desired products)i.e. C18H38(g) ----(zeolite catalyst)---> 4 C2H4(g) + C10H22(g)

3. Steam cracking of ethane and propane – ethane from natural gas deposits fed into furnaces with steam, heated between 750 – 900oC causing much ethane to be converted to ethylenei.e. C2H6(g) -> C2H4(g) + H2(g)

Propane can also be used:C3H8(g) -> C2H4(g) + CH4(g)


Dilute H2So4 cat w/ water

KMnO4, H+

Acidified

HCl, non aqueous solvent

Br2, non aqueous solvent

Identify that ethylene, because of the high reactivity of its double bond, is readily transformed into many useful products

Ethylene’s C=C double bond is highly reactive, allowing it to react with molecules to form many useful products

Reaction of alkenesCharacteristic reaction of alkenes is addition reaction. Two new atoms or groups of atoms are added across double bond, one to each carbon. The C=C is converted to a single bond and a saturated hydrocarbon is produced. General eqn:

H2C=CH2 + X-Y => XH2C-CH2Y

1) Addition of hydrogen to ethylene (hydrogenation) - ethylene to ethane by heating with hydrogen in presence of nickel, platinum or palladium

2) Dibromoethane - Used as a petrol additive - halogen reactions are useful for distinguishing between saturated and unsaturated hydrocarbons. E.g. A non aqueous solution of bromine (e.g. solvent carbon tetrachloride) when added to an alkene causes the solution to lose its colour as bromine becomes incorporated into the alkene:

[CH2=CH2(g) + Br2(l) CH2Br-CH2Br(l)] (petrol additive)


Alkanes do not react with NA bromine unless exposed to UV. In aqueous solutions, the reaction may be the same as above, but due to the presence of water products can include:

[CH2=CH2(g) + HOBr(aq) CH2OH-CH2Br]

Hydrogen bromide reaction:

[CH2=CH2(g) + HBr(g) -> CH3 – CH2Br] (What states?)

3) Chloroethene – Monomer for PVC

[2CH2=CH2(g) + Cl2(g) + ½ O2(g) 2CH2=CH-Cl(g) + H2O(g?)]

4) Styrene – produces from benzene and ethylene via the intermediate ethylbenzene

5) Ethanol – Used as a fuel in automobiles and as an industrial solvent

[CH2=CH2(g) + H2O(l) CH3-CH2OH(l) ]

6) Ethylene oxide and ethanediol – fumigant (former), manufacture of polymers (polyester fibres and PET) and antifreeze (latter)

[C2H4(g) + ½ O2(g) C2H4O(g)]

[C2H4O(g) + H2O(l) OH-CH2-CH2-OH]

Identify the following as commercially significant monomerso Vinyl chlorideo Styrene

By both their systematic and common names

Vinyl chloride – chloroethene CH2 = CHClMonomer for the production of PVC plastics which are widely used in applications such as electrical insulation, plumbing and garden hoses with various additives to change physical properties

Styrene – ethylbenzene C6H5CH=CH2 (also known as phenylethene)Production of polystyrene, most stiffened of common plastics due to large phenyl side group. Stable due to presence of C-C and C-H bonds only, minimal chain branching means it can be formed into clear objects. Tool handles, car battery cases, CD cases. Gas can be bubbled through to create foam (foam drink cups), making it soft and light.

Identify data … to compare the reactivities of appropriate alkenes with the corresponding alkanes in bromine water

Prac – Reactions of hydrocarbons with bromine waterRisk analysis:

Hazard Risk ControlBromine water Corrosive, and toxic, can

cause skin burnsWear safety gogglesUse small amounts to


minimise vapourCyclohexane Highly flammable

Eye and skin irritant with severe redness and pain

Wear safety glassesKeep away from hot surfaces, flames or sparks

Toluene Highly flammable, fire hazardEye and skin irritant with severe redness and pain

Keep away from hot surfaces, flames or sparksPolyvinyl gloves

Aim: To compare reactivities of an alkene (cyclohexene), alkane (cyclohexane), and an aromatic hydrocarbon (toluene) in bromine water

Method:1). Four semi-micro test tubes were half-filled with bromine water, cyclohexane, cyclohexene and toluene respectively, using eye droppers2). Bromine water was mixed with the other substances by placing a few drops of bromine water in each micro-test tube with a dropper3). The test tubes were tapped, and observations recorded

Results:Cyclohexane – noneCyclohexene – forms clear solutionToluene – none

The functional group reacting with bromine is the double bond present in alkenes, this decolourises bromine water. Addition reactions. These reactions are addition reactions:

Bromine w/ water

Bromine w/ cyclohexene (top)

Bromine water w/ cyclohexene (bot)

Toluene did not react as aromatic molecules have delocalised electrons which do something??? ************

Bromine with cyclohexane, this is substitution:

Requires UV light to break off hydrogen atom and allow reaction

Prac – Reaction of lycopene with bromine water


Aim: To determine the effect of bromine water in varying amounts on the spectrum of colours reflected by lycopene

Method: 1). Five semi-micro test tubes were filled halfway with tomato juice2). An eye dropper was used to place 1, 2, 3, 4 and 5 drops of bromine in each of the test tubes respectively3). The solutions were stirred with the stirring rod until colour appeared4). The colours and the corresponding amounts of bromine water in each test tube were recorded

Results:

Test tube no. No. of drops of Bromine Colour1 10 Blue2 8 Turquoise3 6 Green4 4 Khaki5 2 Orange

Varying amounts of bromine in tomato juice changes the number of delocalised electrons in lycopene molecules, changing the spectrum of colours absorbed and resulting in different reflected visible spectra

Identify that ethylene serves as a monomer from which polymers are made

Polymerisation is the process of bonding monomers together to form long chains.Polymers are macromolecules consisting of small repeating units called monomers joined by covalent chemical bonds. Polymers can be divided into two categories:

1. Natural polymers – naturally occurring polymers used by humans since ancient times (E.g. cellulose, silk, rubber)

2. Synthetic – more recent man-made polymers. Replacing natural since they do not corrode, are lightweight and relatively cheap. Celluloid was first commercially manufactured plastic, but highly flammable nature meant it was replaced

Ethylene serves as a monomer due to the reactivity of its double bond. It has a structure that can change to accommodate the additional bond needed to join repeating units together.

Identify polyethylene as an addition polymer and explain the meaning of this term

Polyethylene is an addition polymer, it is created through addition polymerisation.Def: The monomers add to the chain so that all atoms in monomer are present in polymer. It involves unsaturated monomers (a molecule containing a double or triple bond) joining together. One C=C is broken up and resulting molecules link up, since this provides molecules with extra bonding capacity. E.g. for polyethylene

Addition polymerisation requires a catalyst or initiator to start. Other polymers formed by addition are Polyvinyl chloride (PVC), polystyrene and Teflon.

Outline the steps in the production of polyethylene as an example of a commercially and industrially important polymer


General outline: Ethylene can be changed from gas to liquid under high pressure. This liquid ethylene can be heated in the presence of a catalyst to form polyethylene. Two forms of polyethylene can be produced, each with differing methods and varying properties:

o LDPE (reaction conditions 100 – 300oC, 1500 – 3000 atm) – Polymerisation consists of three stages

Initiation – organic peroxide catalyst. They produce free radicals (molecules with unpaired electron), such as H-O. which is a hydroxy radical. This causes the double bond in ethylene to break and form a bond with the radical. CH2=CH2 +R● RCH2-CH2●

Propagation - The resulting molecule contains an unpaired electron. Bonds to another ethylene molecule through the same process etc. (Chain propagation reactions). Backbiting, where the chain curls onto itself and the free electrons takes a hydrogen atom from an existing CH2 group, causes branching.

Termination – at various times, it is possible for two free radical polymers to react to form a covalent bond, ending propagation (chain terminating reaction).

o HDPE (50 – 75oC, <1 atm) – polymerisation process is same as above, but Ziegler-Natta catalyst (TiCl4, Al(C2H5)3 used. Ethylene molecules are added to chain on surface of catalyst, reducing backbiting and branching.

Comparison of structures and properties

LDPE HDPEHigher degree of branching, meaning less

dispersion forces between strands, making it softer and more flexible

Lower degree of branching, meaning more dispersion forces within strands,

making it harder and more rigidLess dense More dense

These products are plastics. Plastics are manufactured materials containing combinations of organic and inorganic elements. They are solid in the finished state but fluid at some stage, and able to be formed into shapes by application of heat and/or pressure.

Factors affecting the properties of polymers: o Length of chain (no. of monomer units) – those with longer chains are stronger since

greater dispersion forces between chainso Arrangement of chains relative to each other – when chains are unbranched, they are

lined up and closely packed creating crystalline areas resulting in stronger and less flexible plastics. Amorphous regions were alignment is more random, produce weaker and softer plastics. Polymer fibres drawn through a small hole aligns them and increases strength

o Function groups in monomer units – polar functional groups increase intermolecular forces between polymer molecules, increasing hardness.

o Cross-linking between polymer chains – covalent links between polymer chains makes polymers very hard and difficult to melt.

o Additives – few polymers are used in pure form, additives improve or extend properties. Additives can include pigments, plasticisers to soften, stabilisers to increase resistances to decomposition etc.


The variable chain length leads to many uses, with shorter lengths for food packaging and milk containers, to longer lengths (800,000 atoms per molecule) for artificial ice rinks.

Describe the uses of the polymers made from the above monomers in terms of their properties

Polymer Properties UsesLDPE 1). Resists water and chemicals

2). Electrical insulator

3). Easy and cheap to process

4). Waterproof

5). Non-toxic

1). Pipes for farm and industry2). Wire and cable sheathing for telephone, coaxial, submarine television and radar3). Shopping and garbage bags4). Milk and fruit juice packs, food containers

HDPE 1). Can take high pressures

2). High tensile strength

3). Chemical resistance

4). Durability and toughness

1). Coating in steel pipes in high pressure gas mains2). Fibres for ropes, fishing nets3). Moulded into containers to hold petrol, oil, detergents and acids4). Children’s toys, plastic buckets, playground equipment

Polyvinyl Chloride

1). Soft and pliableOR (depending on additives)2). Rigid

3). Resists burning

4). Low static electricity

1). Wallpaper, clothing upholstery2). Water pipes, guttering3). Coatings on materials to make flameproof 4). Flooring; tiles, roll flooring and carpet backing

Polystyrene 1). Rigid and electrical insulator---As foam:2). Chemically unreactive3). Low density---4). Resists high impact

1). Television backing, hairdryers, washing machines2). Food containers3). Marker buoys, surfboards4). Shoe heels, toys

Discuss the need for alternative sources of the compounds presently obtained from the petrochemical industry

Petroleum fractions have been the most convenient and economical raw material for synthetic polymers. However, alternatives are being sought since:

1) The current source is non-renewable, and the move to more renewable resources will allow us to continue manufacturing petrochemical products (supplies will run out)


2) Petrochemical products are (non-biodegradable?) and contribute to the degradation of the environment

o A solution is the use of biomass (organic material from living organisms). All living organisms produce biopolymers, which are naturally occurring polymers made entirely or in large part by living organisms

o They are advantageous since they are renewable, and can be used indefinitely with careful use, and are biodegradable since the bonds within the molecule can be broken down by bacteria and fungi, so they do not contribute to the degradation of the environment.

Explain what is meant by a condensation polymer

A condensation polymer is a polymer that was produced through the reaction of two different functional groups in which a small molecule (usually water) is eliminated and the two groups become linked together. Condensation reactions involve saturated molecules. Common groups are –COOH (carboxylic acid), –OH (alcohol) and –NH2 (amine) group. Condensation polymers do NOT require identical monomers

Describe the reaction involved when a condensation polymer is formed

e.g. Condensation polymerisation of nylon

PracticalAim: to produce nylon using interfacial polymerisation

Equipment: 2 x 100 mL beakersTweezersGlass stirring rod20mL of 1,6 – diaminohexane solution20 mL of 10% sebacoyl chloride in hexane

Procedure:1). 1,6 – sebacoyl chloride was added to a beaker2). The diamino hexane was run very carefully down the side of the beaker, so that the two solutions mix as little as possible3). The white material formed between the two layers was clamped using the tweezers 4). The material was drawn away from the beaker and onto the glass stirring rod, being careful to keep away from sides of beaker5). The material was wound onto the stirring rod, dried and examined

Results/Analysis:

A variety of monomers can be used to manufacture nylon, it is simply a generic name for a group of polyamide polymers, the common feature being the repeated –CONH- bond.

This experiment used interfacial polymerisation. It is thus named since the reactants bond together and form nylon at the contact surface between them.

PET (polyethene terephthalate) is also condensationChem notes David Lee BHHS 2007 10

Describe the structure of cellulose and identify it as an example of a condensation

polymer found as a major component of biomass

Biomass – organic material derived from living organismsGlucose C6H12O6 is a carbohydrate of form:

The presence of five hydroxyl groups allows glucose to form polymers such as starch, cellulose, and glycogen. Cellulose is a biopolymer, a polymer naturally synthesised by living organism. They are condensation polymers, since water is eliminated from a reactive functional group when glucose units join together.

o Many glucose units linked together (polysaccharides). o The glucose units come together, causing two hydroxyl groups to react, a hydrogen ion

dissociates from one hydroxyl and combines together with the other hydroxyl to form water.

o The leftover oxygen atom then forms a covalent bond between the molecules. These parts are called the functional units

o Cellulose is a linear polymer, producing a fibre-like material. The beta linkages result in flat, ribbon-like strands which are closely packed and have strong hydrogen bonds between them (cellulose strands). This gives cellulose its strength and rigid structure.

o It is the main component of plant cell wall and major structural component of woody plants and natural fibres

o This makes it the most abundant polymer known on earth.

Identify that cellulose contains the basic carbon-carbon structures needed to build petrochemicals and discuss its potential as a raw material

Cellulose contains three-carbon and four-carbon chains with attached hydrogen and hydroxl groups. Many polymers such as polypropylene are made from three-carbon and four-carbon monomers. If cellulose can be broken down and these chains isolated, it can be used to produce polymers. Large amounts occur naturally such as in plant cell walls, and large amounts left over from agriculture.

Use as raw material can be achieved by:1. Modification of existing biopolymer chains to meet specific applications (e.g. addition of

functional groups)2. Breakdown into smaller molecules which can then be used to build synthetic polymers e.g.

thermochemical (steam/acid) pre-treatment followed by hydrolysis using enzyme cellulase. This produces glucose which can be dehydrated to ethene. However, this is more expensive than using hydrocarbon sources.

Rayon is created from regenerated cellulose sourced from waste paper, straw, husks from wheat and corn and wood pulp. The fibres are chemically treated with sodium hydroxide and carbon disulfide to soften and break them down into smaller units.


A potential area is the use of biopolymer-based plastics is food wrappers and disposable containers, since they are used only once. An example is a US company which has made packaging products from corn starch.

Use available evidence to gather and present data from secondary sources and analyse progress in the recent development and use of a named biopolymer

o Biopol, PHA (polyhydroxyalkanoates) copolymers, a family of microbial energy reserves accumulating as granules within the cytoplasm of cells.

o PHA’s polyester thermoplastics properties similar to oil-derived polymers (e.g. melting temp 50-180oC) mechanical properties can be changed to range from elastic rubber or hard crystalline

plastic.

Simplest are PHB’sProduction is carried out by the following process:

o Alcaligenes Eutrophus, a bacteria widely found in soil and water, is fed a precise combination of glucose and propionic acid, producing PHB’s as energy storage

o The PHB’s are extracted and can be polymerised to create a plastic with properties similar to polypropylene Excellent flexibility and toughness Stable in air, humid conditions Biodegrades in microbially active environments, bacterial and fungi microorganisms

can utilise PHA’s as a source of energy by breaking it down using enzymes (depolymerases).

Potential useso Biocompatability - useful in several medical applications such as controlled drug release,

medical surtures, bone plateso Flexibility and toughness - structural materials in packaging products o Biodegradable – can be used in food packaging, natural breakdown reduces landfill

DevelopmentCurrent work by Metabolix, successfully engineered bio-factories to demonstrate economic production of a broad range of PHA’s. Demonstrated fermentation on a tonnage scale, cost to be under a dollar a pound.

Currently working to produce PHA’s directly in non-food crop plants.

Disadvantage: Current high cost of production as opposed to crude oil sources

Describe the dehydration of ethanol to ethylene and identify the need for a catalyst in this process and the catalyst used

Describe the addition of water to ethylene resulting in the production of ethanol and identify the need for a catalyst in this process and the catalyst used

Model the above two processes


Ethanol and is an alkanol. It is tetrahedral about each carbon and bent around oxygen atom (not shown here). Ethylene can be made through dehydration, heating with concentrated sulfuric acid or a porous ceramic catalyst (>350 in industry):

Reverse reaction is hydration, requires dilute aqueous sulfuric acid:

They are general, apply to any alkanol or alkene e.g. 1-pentanol to 1-pentene

Describe and account for the many uses of ethanol as a solvent for polar and non-polar substances

Risk Analysis

Hazard Risk ControlIodine Toxic – fatal if swallowed.

Corrosive, causes burns and damaging to lungs if

inhaled

Safety glasses, effective ventilation

Oxalic acid Poisonous if swallowed, inhaled or absorbed

through skin

Gloves, avoid generation of dust

Prac – Ethanol as a solventAim: To test the solubility of various materials in ethanol

Method:1). 20 mL of ethanol was poured into each of 10 test tubes using measuring cylinders2). A rice-grain amount of each solid was placed into successive test tubes, and a few millilitres of each liquid placed into the remaining test tubes using an eye dropper. 3). Each test tube was agitated by tapping and gently shaking4). Observations were recorded

Results:Solute Solubility

Sodium chloride NoNapthalene Slightly

Cyclohexanol YesGlycerol YesIodine Yes, dark red

Oxalic acid Yes, purpleBoric acid YesGlucose Yes

Wax NoUrea Yes

o Ethanol has a single hydroxyl group which is attached to an aliphatic

o Allows it to dissolve substances with polar covalent bonds, hydrogen bonds form


o Able to dissolve hydrocarbons and non-polar due to the formation of dispersion forces with the aliphatic group.

o Widely used as alternative solvent in dissolving medicines, cosmetics, food flavourings, alcoholic beverages, low toxicity so relatively safe

Ionic (sodium chloride): Unable to dissolve since strong ionic bonds holding atoms together, and intermolecular formed inadequately strong to break apart lattice

Polar covalent bonds (cyclohexanol C6H11OH, glycerol C3H5(OH)3, oxalic acid C2O2(OH)2, Boric acid B(OH)3, Glucose C6H12O6, Urea CO(NH)2) : Polar covalent bonds such as those in hydroxyl groups allowed ethanol to bond strongly and dissolve them

Macromolecules (Wax C24H50) : Though only held together with weak dispersion forces, its large size means a larger surface area of contact between molecules and thus more total dispersion forces. Ethanol was unable to dissolveNon-polar molecules (Iodine, Napthalene, heptane, pentane) – iodine is diatomic and has very weak dispersion forces holding together, but ethanol can form dispersion forces with iodine molecules and pull away. Napthalene is an aromatic hydrocarbon and does not attract strongly, but dissolves in a similar fashion, same for heptane and pentane which are short-chain hydrocarbons.

With 1,2,3 – propanetriol:

Outline the use of ethanol as a fuel and explain why it can be called a renewable resourceEthanol is a flammable liquid, burning with the reaction:

It is also easily transportable, and was used by hikers and campers. It has thus been proposed as an alternative fuel source, having already been used as an ‘extender’ in world war 2. The purpose of ethanol is to:

1. Reduce greenhouse gas emissions2. Reduce reliability on non-renewable fossil fuels

Engines would not need any modification to run 10-20% ethanol fuel, and is renewable since synthesised in sugar cane from carbon dioxide, water and sunlight. Burning produces carbon dioxide and water which can then be re-used to produce ethanol, so it follows an almost indefinite material cycle.

It has thus been promoted for motor cars to supplement and replace petrol.

Describe conditions under which the fermentation of sugars is promoted Summarise the chemistry of the fermentation process Present information from secondary sources by writing a balanced equation for the

fermentation of glucose to ethanol

o Fermentation requires a carbohydrate as starting material, such as glucose, sucrose or starch

o Disaccharides such as sucrose and polysaccharides such as starch need to first be broken down into monosaccharides such as glucose/fructose by enzymes in the mixture

o This carbohydrate is placed in the presence of yeasts, which produce enzymes that break it down to ethanol and carbon dioxide


o The optimum conditions are 37oC and anaerobic conditionso Reaction is exothermic, so temperature needs to carefully controlledo Fermentation can occur until 15% ethanol, then the yeasts cannot survive and

fermentation stops – extra ethanol can be added or distillation used to increase concentration

Solve problems, plan and perform a first-hand investigation to carry out the fermentation of glucose and monitor mass changes

Prac – Fermentation of glucoseAim: To investigate the fermentation of glucose

Method:1). One gram of beef extract and 25 grams of glucose and 7 grams of yeast were measured out using an electronic beam balance2). A test tube with barium hydroxide was weighed on the triple beam balance and its weight recorded3). The beaker was filled with 300mL of tap water and heated over a bunsen burner until 40oC4). The beef extract, warm water, yeast and glucose were quickly poured into the conical flask and weighed on the electronic beam balance5). The apparatus was set up as shown below:

6). After a week, the tubes and stopper were removed, then the fermented solution and test tube with barium hydroxide were weighed on the triple beam balance and analysed7). Experiment was repeated 2 times7). Steps 1-6 were repeated without the yeast, to act as a control

Results:

Result Loss in mass of conical flask and contents (g)

Gain in mass of test tube contents (g)

1 8.0 5.92 9.8 4.83 9.1 5.0

Average 9.0 5.2

Control:Result Loss in mass of conical

flask and contents (g)Gain in mass of test tube

contents (g)1 0 02 0 03 0 0

Average 0 0

The formation of carbon dioxide was evidenced by the formation of barium carbonate in the test tube. The control showed that the loss in mass, and thus creation of carbon dioxide, was caused by the yeast fermenting glucose. Equation:


Reaction of Barium hydroxide with carbon dioxide:

Nine grams of carbon dioxide was given off in this experiment (loss in mass of conical flask) the gain in the test tube was not used since some carbon dioxide escaped through holes around the stopper.

Through stoichiometry, 1 mole of glucose produced 2 moles of CO2 and ethanol each. Therefore 0.20 moles of ethanol produced

Mass of ethanol = molar mass ethanol x moles produced = 10g (3.3 % w/v)

This reaction did not go to completion, 6 grams of glucose left. This is most likely because the yeast were saturated in ethanol and could no undergo further fermentation, or were no left for sufficient time. Some discrepancy could have been caused by measurement error.

Process information from secondary sources to summarise the processes involved in the industrial production of ethanol from sugar cane

Industrial production of ethanol from sugar cane1. Feed preparation

o Crushing – sugar cane is crushed to remove high-quality sugars and molasses, which are used in fermentation

o Saccharification – bagasse, the other constituents of sugar cane (50% cellulose) undergo a multi-step hydrolysis process, using an enzyme and sulfuric acid as catalyst to produce glucose compounds

3. Fermentation – yeast and anaerobic conditions ferment glucose compounds to a certain concentration of ethanol (max 15%). (Note: Fermentation less efficient than hydration)

4. Purification of mixture – waste products are removed and distillation is used to concentrate ethanol

5. Addition of gasoline – varying amounts of gasoline are added to produce a commercial product, ‘gasohol’ or E10 is 10% ethanol 90% gasoline

Assess the potential of ethanol as an alternative fuel and discuss the advantages and disadvantages of its use

Note: it is more expensive to dehydrate ethanol than it is to purify ethylene from crude oil (this doesn’t really go here, I dunno where to shove it)

Ethanol is flammable liquid that is suitable as a fuel. It can be fermented from biomass such as sugarcane and corn, requires land for agriculture and infrastructure for a fermentation industry to be set up. If these requirements are met adequately, the advantages of ethanol as an alternative fuel are:

o Renewable – products of ethanol combustion can theoretically be completely recycled to produce more ethanol

o Intrinsic anti-knock properties – circumvents the need for toxic anti-knock agents due to presence of oxygen, increases octane of fuel


o Burns more cleanly due to presence of oxygen, , reducing toxic emissions (such as hydrocarbons)

o Reduction of net emission of greenhouse gases due to reuse of carbon dioxide by biomass used to produce ethanol

o Lower blends (<20%) do not require engine modificationo Predicted that household wastes will be recycled to produce ethanol – reduces dumping

emissions

Disadvantages are:o Large areas of arable land needed for agriculture, associated land degradation such as

erosion and fertiliser run offo Blends above 10% shown to be damaging to cars designed for gasoline, including

increased carbon deposits on pistons and corrosion of metallic engine componentso Greenhouse reductions hindered by use of fossil fuels and emission of toxic waste

products in transportation/production of ethanol

Process information from secondary sources to summarise the use of ethanol as an alternative car fuel, evaluating the success of current usage

Higher blends of ethanol require special engines suited, lower blends (<20%) do not require this. It is extensively used in some countries, such as Brazil and the US.

o In Brazil, a large portion of cars are currently “flex-fuel”, allowing them to use both ethanol and gasoline, 80% of cars produced in 2005 were flex-fuel

o Government subsidies and rising petroleum prices have successfully encouraged mass-uptake of ethanol use. Pure ethanol and 25% ethanol are available at nearly all gas stations

o Its use has had noticeable improvements on air quality due to more complete combustiono Brazil is approaching self-sustainability in areas of ethanol use due to its large areas of

arable land and tropical climate

The Bad:o There are situations where ethanol cannot replace fossil fuels, such as diesel fuels, and

they continue to be burnto The popularity of ethanol depends on government subsidies. Production of ethanol

requires a large investment of money and energy, and costs more than petroleum to produce

o Requires destruction of rainforest which acts as carbon sink, offsets greenhouse reductions

In Australia:o Ethanol costs more than petrol to produce, so subsidies are provided to encourage

addition of ethanol to petrolo Ethanol cars are in a small minorityo Public suspicion about fuel, since some independents add excessive ethanol causing

engine wear, and claims by manufacturers that blends above 10% will void warranties. Federal government has decided to limit mixtures to max 10% ethanol

o Skepticism about mass-implementation, no reliable studies showing improvements in air quality through ethanol industry and worry about associated environmental costs such as land degradation

o Ethanol uptake is less successful due to lack of arable land for feedstock growth, higher individual wealth and higher costs of labour

Evaluation:Chem notes David Lee BHHS 2007 17

Highly successive in some countries, but less successful in others with less land resources. Usage is hindered economically depending on economic situation, and for some is not economically feasible due to high energy and monetal cost.

Define the molar heat of combustion of a compound and calculate the value for ethanol from first-hand data

Molar heat of combustion is the amount of heat energy released when one mole of the substance undergoes a combustion reaction.

Prac – Heat of Combustion of AlcoholsAim: To measure the amount of energy produced when alcohols are burned and thus calculate their molar heats of combustion

Method:1). A copper calorimeter and spirit burner were weighed using an electronic beam balance2). A measuring cylinder was used to pour 100 mL of water into the copper calorimeter3). The apparatus was set up as show below:

4). The initial temperature of the water was measured using the thermometer5). The spirit burner was uncapped, lit using a matched then allowed to burn for 30 seconds6). The spirit burner was capped and the temperature of the water taken again

7). The mass of the spirit burner was weighed on the electronic beam balance8). Steps 3-7 were repeated for the other alcohols

Results:Molar heats of combustionMethanol: -410 kJ/molEthanol: -640 kJ/molPropanol: -980 kJ/molButanol: -1100 kJ/mol

Butanol produced the most soot since it had the longest carbon chain, increasing tendency for incomplete combustion:

Shorter chain fuels release less energy per molecule and react at an overall slower rate, meaning that the immediate availability of oxygen molecules is adequate to ensure complete combustion. Other fuels react at faster rates due to larger release of energy per molecule, and the rate of oxygen diffusion into the immediate environment of the fuel molecules is inadequate to prevent atoms


within the fuel reacting amongst themselves. These latter reactions occurs less readily than complete combustion reactions, but form eventually with adequate particle energy and collision.

o Butanol produced the most heat per gramo The larger the molecules, the more heat released per gramo This suggests that the breaking of a CH2 group is more exothermic than breaking off an

‘OH’ group, since one gram of lighter alcohols would contain more hydroxyl groups

Ethanol would be the best fuel source since it has the best balance of achieving complete combustion and burning with more heat per gram.

Safety considerationsDanger in transporting spirit burners – carried only while unlit so would not ignite if dropped and shatteredMethanol – vapours are toxic in larger doses, do not open spirit burner Ethanol – skin and eye irritant, wear safety glasses, do not open spirit b urner

Errors:Experimental results differed from theoretical results due to:

o Conduction and radiation of heat from copper calorimeter into surrounding environment, reducing heat of combustion values

o Inaccuracies in measuring equipment

This could be remedied by:o Burning for shorter periods of time, less radiation/conduction of heat

Explain the displacement of metals from solution in terms of transfer of electrons Identify the relationship between displacement of metal ions in solution by other metals to

the relative activity of metals

A displacement reaction is where a metal converts the ion of another metal to the neutral atom.

o Different metals have different reactivitieso Metals with higher electronegativity or lower reactivities will attract electrons more

strongly, or is the stronger oxidanto When a solid, pure metal is in contact with a solution of another metal’s ions, the metal

with the lower electronegativity, or higher reactivity will displace the other metal in solution, since it has weaker electron pull and loses an electron to the other metal to become a cation

o Electrons are thus transferred from one metal atom to the other, one becoming a neutral atom and depositing out of solution, and the other become a cation going into solution

e.g.A granule of zinc is dropped into a blue solution of copper sulfate, zinc gets covered with reddish-brown copper:

Oxidation half-reaction:

Reduction half reaction:

The anion is a spectator ion. The activity series can be seen on the table of standard potentials


Account for changes in the oxidation state of species in terms of their loss or gain of electrons

1. The oxidation of an element in its stable elemental state is 02. The sum of oxidation states in an element or compound = 0, and for a polyatomic ion the

charge of the ion3. An increase in oxidation state means a loss of electrons and vice versa.

Examining oxidation states is useful in determining whether a redox reaction has occurred during a chemical reaction. Some elements display multiple oxidation states.e.g.In Cu2O (2Cu+, O2-)Oxidation state of copper is +1

Outline the construction of galvanic cells and trace the direction of electron flow Describe and explain galvanic cells in terms of oxidation/reduction equations Define the terms anode, cathode, electrode and electrolyte to describe galvanic cells

A simple galvanic cell requires:o 2 solid metal pieces to serve as electrodeso 2 containers (e.g. beakers)o 2 different electrolyte solutionso A salt bridge with electrolyteo Connecting wire

1. Electrodes need to be matched with appropriate electrolyte, optimum metal ion same as metal of electrode. Electrolyte must have greater or equal reactivity than electrode, else displacement will occur

2. Salt bridge needs to be in contact with both solutions, and ions cannot form precipitate otherwise there will be no charge neutralisation (causing opposing emf which impedes current)

3. Conducting wire links both electrodes


The electrode are the conductors of a cell which get connected to an external circuitThe anode is the electrode where oxidation occursThe cathode is the electrode where reduction occursThe electrolyte is a substance which in solution or molten conducts electricity

Standard conditions are 1M solution and 25oC. Acidic conditions alter the potential difference.

Process:1. The electrode with lower reactivity attracts electrons from the other electrode through the

conductor2. A redox reaction results:

e.g.

3. The ion formed goes into solution, and the anion dissolves. The metal ions in solution around the cathode obtain the electron and plate onto the electrode

4. Ions flow from the salt bridge into the electrolyte solutions to neutralise charge and remove opposing potential difference

Detailed purpose of salt bridgeIf there is no salt bridge:

o As the redox reaction occurs, the electrolyte in contact with anode will have an increasing excess of positive ions as the electrode dissolves, and the electrolyte in contact with cathode will have an excess of negative ions as positive ions precipitate

o This imbalance of charges produces a potential difference against the direction of electron flow, eventually stopping it

The salt bridge:

o Allows ‘migration’ of charge, and positive ions in the salt bridge move into the negative solution, and vice versa

o This preserves electric neutrality, and eliminates negative potential difference

Cell diagramsType 1:Metal/metal ion electrodee.g.

Cu|Cu2+||Ag+|Ag| = phase separator|| = salt bridge

Reaction progresses from left to right

Type 2:Inert substance such as platinum or carbon and equimolar amounts of non-metal and its ion. e.g.

Pt(s) | I2(s)|I-(aq) || Fe2+

(aq)|Fe3+(aq) | Pt(s)

Reaction progress:

2I-(aq) I2(s) + 2e- (oxidation)

Fe3+ + e- Fe2+(aq) (reduction)


What about gas?

Perform a first-hand investigation to identify the conditions under which a galvanic cell is produced

Perform a first-hand investigation and gather first-hand information to measure the difference in potential of different combinations of metals in an electrolyte solution

Prac – Galvanic Cells

Aim: 1. to construct a galvanic cell called a Daniell cell and investigate conditions under which it

operates (A)2. to compare the effect of using different combinations of metals in electrodes (B)3. to compare the effect of different volumes of electrolytes under otherwise identical

conditions (A)

Method A:1). A zinc strip and copper strip were cleaned with sand paper2). A half cell consisting of a copper strip resting in a 250 mL beaker half-filled with copper sulfate solution was constructed3). A similar cell with a zinc strip and zinc sulfate solution was constructed, and these half-cells linked with a piece of filter paper soaked with potassium nitrate solution and folded4). The copper strip was connected to the positive terminal of a voltmeter, and te zinc strip to the negative terminal5). The reading and polarity were recorded before quickly disconnecting the voltmeter6). The beakers were then completely filled with corresponding electrolyte solutions, and the voltage measured again

Results:Beakers half: 0.20V Current: 0.5 mABeakers full: 0.25V Current: 0.8 mAAnode: ZincCathode: Copper

At anode (oxidation): Zn(s) -> Zn2+(aq) + 2e-

At cathode (reduction): Cu2+(aq) + 2e- -> Cu(s)

Overall: Zn(s) + Cu2+(aq) -> Zn2+

(aq) + Cu(s)

The salt bridge allows migration of ions into each beaker (cations into cathode solution and vice versa) to neutralise charge buildup and maintain cell voltage.

A larger volume of electrolyte solution means a large surface area of electrolyte in contact with electrodes. This increases the rate at which charged particles are removed from the electrodes, decreasing the internal resistance and thus the current. However, the voltage within the cell is caused by the intrinsic properties of the electrodes (difference in electronegativity), and thus is not altered significantly. It is affected slightly since a charge buildup generates a slight back emf, which reduces the voltage, and more electrolyte action reduces this. Note: Concentration of electrolyte does affect voltage.

The current output needs to be measured quickly since the ions in salt bridge are being used up, and charge begins to build in half-cells opposing current flow.

Sources of error:


Electrodes not completely polished – reduces effective surface area of action for electrolyte and thus current

Method B:1). The copper half-cell in A was set up2). Another half-cell consisting of a magnesium strip in magnesium sulfate solution was set up3). The electrodes were connected to the terminals of a voltmeter, and voltage reading recorded4). Steps 2-3 were repeated with:

a. Aluminium in 1M aluminium nitrate solutionb. Tin strip in tin(II) nitrate solution

Results:

Test Half-cell Polarity (relative to Cu/Cu2+)

Total cell voltage

Zn/Zn2+ -ve 0.50Mg/Mg2+ -ve 1.1Sn/Sn2+ -ve 0.35Al/Al3+ -ve 0.90

Standard potential Cu – +0.34 V (Oxidation potential -0.34V)

Half Reaction (write in during exam)

Predicted voltage E0 (V) Experimental E0 (V)

Zn -0.76 -0.16Mg -2.36 -0.76Al -1.68 -0.56Sn -0.14 -0.01

The more active the metal, the greater the potential difference.

Note: Oxidation potential is the ability of a substance to oxidise in relation to hydrogen, reductional potential is ability to reduce in relation to hydrogen. E.g. Copper has reduction 0.34 meaning it has higher ability to reduce, but its oxidation is -0.34, it has less ability to oxidise.

Gather and present information on the structure and chemistry of a dry cell or lead-acid cell and evaluate it in comparison to one of the following:

o Button cell (Silver Oxide cell)In terms of:

o Chemistryo Cost and practicalityo Impact on societyo Environmental impact

Criteria Dry Cell (Leclanche Cell) Button Cell (Silver Oxide Cell)Structure

Error in diagram: ‘C’ in electrolyte should be ZnCl2

Note: Porous medium acts as ion bridge, and prevents direct contact of anode and cathode

Chemistry Oxidation Zn(s) -> Zn2+

(aq) + 2e-

Reduction2MnO2(s) + 2H+

(aq) + 2e- -> Mn2O3(s) + H2O(l)

OxidationAmalgamated zincZn(s) + 2OH-

(aq) -> ZnO(s) +H2O(l) + 2e-

ReductionSilver oxide


Hydrogen ions provided by ammoniumNH4

+<-> NH3(aq) + H+(aq)

OverallZn(s) + 2H+

(aq) + 2MnO2(s) -> Zn2+

(aq) + Mn2O3(s) + H2O(l)

Ag2O + H2O +2e- -> 2 Ag(s) + 2OH-

Electrolyte KOHOverallZn(s) + Ag2O -> ZnO(s) + 2Ag(s)

Cost and Practicality Adv:o Inexpensiveo Robusto Easy to store and useDis:o Short lifeo Voltage not as constant

as silver button (comparison)

o Cannot deliver very high currents

o Cannot be recharged

Adv:o Compacto Provides constant voltage over long

period of time, since solid reactants and products have fixed concentration

o Overall long operating life due to solid components

Dis:o Silver is expensiveo Not rechargeable

Impact on Society o First commercial battery, made portable electric devices possible

o Used widely in toys, torches, radios etc.

o Has allowed efficient powering of miniature devices e.g. watches, hearing aids

Environmental impact

o Manganese readily oxidised to stable manganese (IV) dioxide, becomes immobilised and not dangerous

o Small quantities of ammonium salts and zinc not harmful

o Not rechargeable, large volume in landfills so space is an issue

o Contains traces of mercury, causes problems with unsafe disposal

o Non-rechargeable so takes up space

Solve problems and analyse information to calculate the potential Eo requirement of named electrochemical processes using tables of standard potentials and half-equations

Eocell = Eo

cathode - Eoanode

Write down half-cell equations, then balance to get overall. These values are under standard conditions (298K, 1M electrolytes)A higher concentration of reactants relative to products increases spontaneity of reaction and thus emf.

Distinguish between stable and radioactive isotopes and describe conditions under which a nucleus is unstable

o The spontaneous emission of radiation by certain elements is called radioactivityo Some elements have all isotopes radioactive, some only one or someo These particles are thus referred to as radioisotopes


Stable isotopes Unstable isotopesNo radiation emission Emission of radiation

Z ≤ 83 Z > 83Ratio of neutrons to protons within zone

of stabilityRatio of neutrons to protons out of zone

of stability

Zone of stability:

A nucleus is unstable when its ratio of neutrons to protons is outside zone of stability. For light elements (Z < 20), 1.0. Stability ratio steadily increases as atomic number increases, up to 1.5 for Z = 83. Past this, all are unstable due to large size of nucleus.

Describe how transuranic elements are produced

Transuranic elements are elements with atomic number above Uranium (92)

o Some isotopes undergo fission when bombarded, others undergo nuclear reactions to form new elements

o When non-fissionable atoms such as Uranium 238 are bombarded with high speed particles, it absorbs the particle to become an unstable atom

o It then rapidly decays to form a new element

There are two machines that are used to produce high speed positive particles to produce transuranics.

o Linear accelerator – accelerates positive particles in straight line along axes of series of positive and negative cylinders, accelerating it. Often more than a kilometre in length

o Cyclotrons – accelerates positive particles by passing them through alternating positive and negative electric fields. A strong magnetic field is used to constrain particles to spiral path, reducing size of machine.

They can also be produced in nuclear reactors, a source of neutrons. Neutrons do not experience electric repulsion like positive nuclei, and thus speeds in nuclear reactors are adequate. These create transuranic elements with a proton deficiency.e.g.


Mode Emission Atomic mass Atomic number

α decay -4 -2

β decay 0 +1

Positron emission

0 -1

Electron capture

(absorption)

0 -1

γ emission 0 0

Neptunium, first discovered transuranic element obtained in by chemical separation of nuclear fission reactor products.This can be further bombarded to create plutonium:

Unstable

23 transuranic elements have been created thus far.

Process information from secondary sources to describe recent discoveries of elements

Recently discovered elements include:o Ununoctium (118, October 10 2006) – heaviest element discovered to date. It was

indirectly detected by a team of researchers working in Russia at Dubna University’s Joint Institute for Nuclear Research when they detected its decay products after bombarding californium-249 atoms with calcium-48 ions. Very unstable, half-life 0.89ms. Reaction:

o Ununpentium (115, February 2 2004) – Russian scientists at Dubna “…” and American scientists at Lawrence Livermore National Laboratory announced they produced 4 atoms of Uup which quickly decayed into Ununtrium (113) in about 100 ms. They bombarded Americium with calcium.

Note these elements are temporarily using IUPAC systematic element names, before they are officially named.

Identify instruments and processes that can be used to detect radiation

o Photographic film – photographic film darkens in the presence of radiation. Used in radiation badges worn by laboratory workers handling radioactive substances to determine radiation dosage

o Geiger-Muller tube – Radiation enters window and ionises gas particles inside Geiger tube (inert gas such as argon), and the resulting charged particles are accelerated to the two plates with a potential difference. They further ionise other argon atoms through collision creating a cascade effect. This creates a signal which is amplified and converted into an audio signal.

o Scintillation counter – ionizing radiation hits the scintillation crystal (depicted), the electrons are excited and emit photons which can be detected and amplified by a photomultiplier tube (depicted) to produce a reading.


o Cloud Chamber – contains supersaturated vapour of water or alcohol. Radiation (alpha or beta particles) ionises it, forming noticeable tracks. Alpha trails are broader and straight, whilst beta tracks are thinner and show more evidence of deflection

Describe how commercial radioisotopes produced

Radioisotopes can be produced by bombardment of high-speed particles. Radioisotopes are commercially produced in:

o Nuclear reactors (proton deficient, neutron enriched) – convenient source of electrons. Target nuclei are placed in reactor core and are then bombarded by neutrons to produce desired isotope. These are then separated chemically or physically from other substances within reactor.

Currently operating in Australia for this purpose is HIFAR reactor, managed by ANSTOe.g.Creation of technetium:

Technetium decays, releasing gamma ray inside body

o Cyclotron (proton rich, neutron deficient) – neutron deficient isotopes must be produced in a cyclotron. They are bombarded with a small positive particle such as a helium or carbon nucleus at high speed in order to overcome electrostatic forces of repulsion

National Medical cyclotrone.g.Creation of gallium-67

Identify one use of a named radioisotope:o In industryo In medicine

Describe the way in which the above radioisotopes are used and explain their use in terms of their chemical properties

Technicium-99m (nuclear reactor)Most widely used in medicine for diagnosis, such as locating brain tumours or studying other parts of the body by being attached to red blood cells.

o Short half-life of 6 hours means patient exposure is minimisedo Versatile chemistry and can be incorporated into range of biomolecules targeting

different organs


Iodine-131 – Produced in cyclotron or nuclear reactor, testing of thyroid function and treatment of thyroid ailments such as overactive thyroid or thyroid cancer (beta decay destroys some thyroid cells)

o Iodine-131 is naturally absorbed by cells in the thyroid glando Relatively short half life to minimise exposure (8 days)

Specific problems:o Ionising radiation of iodine-131 deals collateral damage to other cells o Radiation penetrates the body and can damage organic tissue near to the patiento Transport and production in nuclear reactors requires stringent safeguards, which is

problematic

Cobalt-60 – used to measure thickness of materials. With fixed geometry for source and detector, penetration of radiation emitted from radioisotope (beta particles in this case) determines thickness of material. Gamma ray producer

o Long half-life so does not need to be replaced frequently (5.3 years)o Low energy emissions, so absorption is significant and can be detectedo Low energy emissions, minimises safety procedures required

Sodium-24Used to detect leaks in water pipes or underground oil pipes. Dissolved into water source, can be subsequently detected in soil around areas of leakage.

o Dissolves easily into watero Relatively short half-life to minimise environmental damage (15 hours)

Use available evidence to analyse benefits and problems associated with the use of radioactive isotopes in identified industries and medicine

Benefits in medicine:o Created wide range of non-invasive diagnostic procedures otherwise impossible, such as

technetium-99m used to identify brain tumours, gallium-67 for cancers, an area very dangerous for surgery

o Allowed radiation therapy to treat many forms of cancer, e.g. iodine-131 for treatment of thyroid cancer

Benefits in industry:o Ability to make more sensitive, precise and reliable monitoring equipment e.g. cobalt-60

for measuring thickness of materialso Allows otherwise difficult activities such as detecting leaks in extensive water

distribution systems, sodium-24 can be dissolved and radiation detected near leaks

Problems:o Exposure of radiation doses to workers in medicine, industry and research can damage

tissues e.g. ionizing radiation of sodium-24 causes cancer by removing electrons from the biological molecule DNA.

o Extra safety precautions are required for sites with radioactive materials, such as proper storage facilities and protective clothing e.g. industries dealing with cobalt-60 and technetium-99m need to filter out the fine radioactive dust produced, which can pose a lung cancer risk

o Disposal of radioactive waste requires space, and can be problematic since isotopes such as Cobalt-60 remain radioactive a long time after they are no longer useful, may leak into environment without strict procedures


The Acidic EnvironmentDefinitions and properties of acids/basesAn acid is a substance that produces hydronium ions (H3O+) in solutionA base is a substance containing oxide or hydroxide ions (O2-, OH-) or which in solution produces the hydroxide ions. A soluble base is an alkali, i.e. one which dissolves or reacts in aqueous solution to produce hydroxide ions. Note that oxygen ion-containers are insoluble or only react.

Common properties of acids:1. sour taste2. sting or burn the skin3. conduct electricity in solution4. turns blue litmus red5. React with reactive metals to form salt + hydrogen gas6. Reacts with carbonates to form CO2, salt and water7. Reacts with metal oxides/hydroxides to form salt and water

Common properties of alkalis:1. have a soapy feel2. have a bitter taste3. conduct electricity in solution4. turns red litmus blue 5. React with amphoteric metals to produce hydrogen gas

Acids and bases react to form salt and water (there are exceptions). Note that the salt is in aqueous solution, separated as ions and not precipitated. E.g. reaction of sodium hydroxide with hydrochloric acid:

HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)

Can also be written as:

H+(aq) + NaOH(aq) - > Na+

(aq) + H2O(l)

OrH+ + Cl- + Na+

+ OH- -> H2O(l) + Na+ + Cl-

The chloride and sodium ions are spectator ions. The net ionic equation is simplyH+

+ OH- -> H2O(l)

A salt is an ionic compound formed when a base (alkali) reacts with an acid.

Hydrohalic acids such as HCl, HBr and HI lead to halide salts. Oxyacids (acids containing oxygen attached to an element) e.g. sulfuric acid, nitric acid, phosphoric acid form salts that end in –ate. Nitrous acid (HNO2) and sulphurous acid (H2SO3) create salts that end in ‘ite’. Anions formed from oxyacids are called oxyanions. Must be familiar with acid formulas.

Classify some common substances as acidic, basic or neutral

Acidic:Vinegar (acetic acid) – used in cooking (~3)Lemon juice (citric acid) (~2.5)Vitamin C (ascorbic acid) – vitamin supplement Hydrochloric acid – pH maintenance in swimming pools, clean bricks cement and tiles (~1)

Neutral:WaterSalts (e.g. sodium chloride, copper sulfate)Chem notes David Lee BHHS 2007 29

Milk

Basic:Baking soda (sodium bicarbonate) (NaHCO3) (~8.5)Oven and drain cleaners (sodium hydroxide) – sodium hydroxide also used in soap, and alumina (~ 13)Lime (calcium hydroxide) – making mortar (~ 11)Ammonia – used to make fertilisers (~ 12)

Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, and that the range is identified by change in indicator colour

Identify data and choose resources to gather information about the colour changes of a range of indicators

An indicator is a substance that in solution changes colour depending on the pH of the solution. There are many different indicators, and the range of pH over which these indicators change colour varies. Litmus is the most common and is extracted from lichens. The indicator changes colour in reaction with the pH of a substance, indicating acidity or basicity dependant on the range of the indicator.

Universal indicator is a mixture of several indicators and works over the whole range

Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity

o pH in soil testing – steps: small amount of moist soil placed in a well on a ceramic test plate White barium sulfate (neutral pH, insoluble) sprinkled on so colour change more

easily discerned 2/3 drops universal indicator added to barium sulfate Colour change of indicator compared to colour chart

Lime (CaO) or dolomite (CaCO3/MgCO3) added if too acidic CaO + H2O Ca(OH)2 Ca2+ + 2OH-

CO32- + H2O HCO3

- + OH-

Sulfur is added if too alkalineS + O2 SO2

SO2 + H2O H2SO3 H+ + HSO3-

o Testing home swimming pools which need to be neutral. Acidic water burns eyes, alkaline water causes skin rashes. Operation of electrochemical cell to produce chlorine makes pool more alkaline. A pool sample is collected into a vial and an indicator, usually


phenol red (yellow -> red) (6.6 – 8.0) is added and compared to a colour chart. HCl added if too basic.

o Monitoring wastes from laboratories that process photographic film, as photographic solutions are highly alkaline and discharges need to be neutral in order to not adversely affect the environment

o pH in aquariums – fish excrete ammonia which reacts with water, making it more basic. Universal indicator is used

Perform a first-hand investigation to prepare and test a natural indicator

Prac – Red Cabbage indicatorAim: to investigate the colour changes of an indicator extracted from red cabbage

Method:1). A handful of shredded red cabbage was boiled in water over a bunsen burner for about 5 minutes2). The resulting liquid solution was poured into a filter paper/filter funnel apparatus and collected in a beaker3). 20 mL of 1 M HCl solution was poured into a measuring cylinder4). 10 mL was poured into a small test tube which was labelled ‘0’5). The remaining 10 mL was poured into a beaker and diluted to 100 mL6). 20 mL of the resulting solution was poured back into the measuring cylinder, and steps 4 – 5 were repeated 5 times, each test tube being labelled a successive integer higher7). Steps 3-6 were repeated starting with 1 M NaOH solution and labelling from 14 down8). A few drops of red cabbage indicator were added to each test tube and observations recorded

Results:Red (0-1) Pink (2) Purple (3-4) Clear (5-6) Purple (7-9) Blue (10-11) Green (12-13) Yellow (14)

There are 6 discrete colour stages for this indicator. This suggests multiple molecules within the red cabbage indicator solution acting to produce colour changes. Each molecule has a specific colour when a proton is added or taken away. The molecules in this case are anthocyanins, and there are about 15 different ones in red cabbage indicator.

Solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic

See book

Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids

An acidic oxide is one which either reacts with water to form an acid, or reacts with bases to form salts (or both). E.g. carbon dioxide and diphosphorous pentoxide P2O5

CO2(g) + H2O(l) H2CO3(aq) 2H+ (aq) + CO3

2-(aq) (carbonic acid)

CO2(aq) + 2NaOH(aq) H2O(l) + 2Na+(aq) + CO3

2-(aq) (sodium carbonate)

Or alternatively:H2CO3(aq) + 2NaOH(aq) 2H2O(l) + 2Na+

(aq) + CO32-

(aq)

The latter is more correct, as the acidic oxide would react with water to form the acid first. It would depend on the relative concentrations of the oxide and the acid in solution, as it is an equilibrium reaction. However, since both create the same products, this is negligible.And similarly for P2O5


A basic oxide show basic character, and react with acids to form salts, but not with alkali solutions e.g.

CuO + H2SO4(aq) CuSO4(aq) + H2O(l)

CuO(s) + H2O(l) Cu2+(aq) + 2OH-

(aq)

Amphoteric oxides are those showing both acidic and basic character, and those that react with neither acids or bases are neutral oxides e.g. NO, CO, N2O

Analyse the positions of these metals in the periodic table and outline the relationships between position and acidity/basicity of oxides

Acidic nature of oxides increases from left to right.

Define Le Chatelier’s principle

If a system in equilibrium is disturbed, then the system adjusts itself so as to minimise the disturbance

Identify factors which can affect the equilibrium in a reversible reaction

Reversible reactions occur when products can react to generate reactants. When a reaction starts, forward reaction generates products from reactants. Backward reaction then generates products, which form at an increasing rate as product concentration increases. The equilibrium occurs at the point where formation of products is equal to the rate of reactant formation, no net change in concentration.

Factors:o Concentration species – increasing/decreasing concentration of a species will cause

reaction equilibrium to shift so that it decreases/increases the species concentration. This because it naturally results in more/less collisions or more/less decomposition to form more/less of that chemical. Note that reactions involving solids and liquids experience little effect, as concentrations remain almost unchanged (note: this does not include dissolved substances).

o Pressure in a gaseous reaction – an increase/decrease will cause a increase/decrease in concentration (and vice versa for volume). Depending on which side of the reaction has more particles, the equilibrium will shift in that direction in order to reduce number of particles and thus pressure (or vice versa). Note that increasing reaction by increasing concentration of gas not involved in reaction e.g. argon has no effect


o Temperature – If the temperature is lowered, the amount of energy in the system decreases and the exothermic reaction is favoured since less particles have sufficient energy to form products with a higher potential energy. And vice versa

o Catalysts – increases speed at which equilibrium is reached, does not alter equilibrium position as activation energy of both product and reactants formation is decreased

Notable exceptions:o When solid or liquid is involved in reaction – the concentration of these substances stays

constanto The addition of water to an aqueous reaction involving water – concentration of water

does not change significantly, but other substances more dilute

Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle

Prac – Degassing soft drinksMethod:1). A 300mL bottle of soft drink was taken and weighed on an electronic beam balance2). The drink was shaken vigorously, and the top slowly and carefully unscrewed to prevent spillage as the gas escaped3). The bottle was reweighed4). 6 grams of table salt was taken and added slowly to the contents of the bottle until no more fizzing was observed5). Bottle was reweighed6). A new bottle was heated over a bunsen burner with cap on, then the cap unscrewed slowly to prevent spillage before reweighing it7).Hydrochloric acid was added until the water stopped fizzing, then it was reweighed

Discussion:The dissolution and reaction of CO2 in water is multistep equilibria –CO2(g) CO2(aq) … (1)CO2(aq) + H2O(l) H2CO3(aq) … (2)H2CO3(aq) + H2O(l) HCO3

-(aq) + H3O+

(aq) … (3)HCO3

-(aq) + H2O(l) CO3

2-(aq) + H3O+

(aq) … (4)

Factors:o Temperature – solubility decreases as temperature increases, opposite to liquids and

solids. Increased average kinetic energy of CO2 molecules means they have greater overall tendency to escape from solution. Equilibrium shifts to left for all equations until new equilibrium is reached

o Pressure – solubility increases with increased pressure, more carbon dioxide dissolves to decrease pressure and act against change, equilibrium shifts to right.

o Dissolution of ions – dissolution of ions displaces carbonic acid ions and CO2 molecules from hydration shells and causes equilibrium to shift to left and increase CO2 gas concentration

o pH of water – increased pH means more hydroxide ions, which react with carbonic acid to neutralise it and produce water, resulting in more CO2 dissolved to produce acid to counteract change. If pH lowered, increased acidity means increased concentration of H3O+ ions, shifting equilibrium of (3) and (4) to left to decrease its concentration. This means increased concentration of the reactants on left, which has a cascade effect shifting all equilibrium to left and increasing CO2 gas concentration.

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen


Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen

Natural sources of sulfur dioxide:o Geothermal hot springs and volcanoes, release is unpredictable and changes with

volcanic activity. Natural levels of sulfur dioxide vary widely, but account for about ¼ of worldwide emissions

o Bacterial decomposition of organic matter, produces H2S which oxidises to form sulfur dioxide2H2S(g) + 3O2(g) 2SO2(g) + 2H2O(g)

Industrial sources are mainly:o Burning of fossil fuels – coal generally contains 0.5% to 6% sulfur as metallic sulfides in

sulfur in carbon-containing compounds, released as sulfur dioxide during combustion in power stations. Some sulfur remains in refined petrol which is released as sulfur dioxide in automobiles.e.g.iron sulfide in coal4FeS(s) + 7O2(g) 2Fe2O3(s) + 4SO2(g)

o Processing of fossil fuels - removal of sulfur from crude oil and natural gas releases some sulfur dioxide.

o Extraction of metal from sulfide ores – first step is to roast sulfide ore in air e.g. extraction of zinc roasting zinc sulfide 2ZnS(s) + 3O2(g) -> 2ZnO(s) + 2SO2(g)

Natural sources of nitrogen oxide/dioxide:o Lightning – high temperatures causes atmospheric nitrogen and oxygen to combine

forming nitrous oxide:O2(g) + N2(g) 2NO(g)

This slowly reacts with oxygen to form:2NO(g) + O2(g) 2NO2

o Denitrifying bacteria – converts nitrates in soil into nitrous oxide (N2O), increased use of fertiliser has increased emissions

Industrial sources:o Combustion – includes power stations and automobiles etc. High temperatures involved

causes atmospheric oxygen and nitrogen to react, and is released into atmosphere(equations same as above).

Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen

Nitrogen dioxide and sulfur dioxide are washed out by rain, so there is no significant buildup in atmosphere. Nitrous oxide however, has steadily increased by about 15%, from measurements made over the last century. There are problems associated with collecting evidence for sulfur and nitrogen oxides, namely:

o Concentrations of both are very low, below 0.1ppm, and only recently (since about the 1950’s) are instruments accurate enough to reliably measure the levels, so trends before this period could be invalid

o Sulfur dioxide and nitrogen dioxide form sulfate and nitrate ions which are changed chemically as they move around the hydrosphere, so measuring traces of these compounds is difficult


Most evidence comes from observed occurrences such as acid rain. There appears to be an increase from data but it is inconclusive due to lack of long-term trends and inaccuracies of earlier measurements.

Analyse information to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment

Concern about release into environment because it has detrimental effect on environment and can cause harm to people since:Sulfur dioxide:

o Sulfur dioxide irritates the respiratory tract, and can cause symptoms in people with asthma or emphysema in concentrations as low as 1ppm

o Forms acid raino Dry deposition causing environmental damage

Nitrogen oxides:o Nitrogen dioxide irritates the respiratory tract, and can cause extensive tissue damage in

concentrations 3-5 ppmo The action of sunlight on nitrogen dioxide, hydrocarbons and oxygen increases

photochemical smog, which includes ozone – poisonous substanceo NO and NO2 participate in ozone layer depletion (NO + O3 NO2 + O2)o Forms acid raino Contributes to global warming

Explain the formation and effects of acid rain

o Sulfur dioxide and nitrogen dioxide gases released dissolve in water to form sulfuric acid and nitric acid which is washed out of the atmosphere by rain, forming wet deposition acid rain. Reaction with hydroxyl radicals:SO2(g) + 2OH H2SO4(aq)

(OR)

2SO2(g) + O2(g) 2 SO3(g)

SO3(g) + H2O(l) H2SO4(aq)

--

2NO2(g) + H2O(l) -> HNO2(aq) + HNO3(aq)

2HNO2(aq) + O2(g) –(catalysed by impurities) 2HNO3(aq)

(OR)

4NO2(g) + 2H2O(l) + O2(g) 4HNO3(aq)

Dry formation:Incorporated into dust and smoke and falls to the ground

o Effects due to low pH include: Corrosion and tarnishing of metal and bridges, soiling and surface erosion of marble

and stone structures CaCO3(s) + 2H+

(aq) → Ca2+(aq) + CO2(g) + H2O(l)


Crown dieback in trees Leeching of leaf nutrients Killing of leaf tissue Leeching of Ca2+ and Mg2+ ions from soil as they are mobilised due to decreased pH,

reducing soil fertility Inhibits microbial activity

Increased acidity of lakes, killing aquatic life e.g. snails can only tolerate up to pH 6.0

Mobilisation of Al3+ ions in soil due to reduced pH. This flows into lakes and precipitates out, clogging fish gills and suffocating them

Calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at OoC, 100 kPa or 25oC and 100kPa

Equal numbers of molecules of different gases occupy the same volume in isothermal and isobaric conditions. At 0oC and 100 kPa, 22.71 L per mole and 24.79 L/mol for the other.

Gather and process information to write the ionic equations to represent the ionisation of acids

Define acids as proton donors and describe the ionisation of acids in water

Acids react with water in solution to form a solution containing hydronium ions and its conjugate base.

Identify acids including acetic acid, citric acid (2-hydroxypropane-1,2,3-tricarboxylic acid), hydrochloric acid and sulfuric acid

o Acetic acid (CH3COOH) aka ethanoic acid – present in vinegar

o Citric acid (C6H8O7) – occurs in citrus fruit, also widely used as a food additive for flavour or as a preservative

o Hydrochloric acid (HCl) – produced by stomach lining glands to break down food molecules, also made commercially to clean metals, brickwork, neutralising bases etc.

o Sulfuric acid (H2SO4) – synthetic acid manufactured to make fertilisers, synthetic fibres etc.

Phosphoric acid (H3PO4) weakPolyprotic acids have more than one ionisable hydrogen per formula unit.For others see book

Describe the use of the pH scale in comparing acids and bases Identify pH as –log10[H+] and explain that a change in pH of 1 means a tenfold change in

[H+]

pH scale is a scale of measurement for hydrogen ion concentration. pH = –log10[H3O+]

This obeys the significant figure rule

Water self ionises:H2O + H2O H3O+ + OH-

Kw = [H3O+][OH-] = 1.00 x 10-14 at 298K


Describe acids and their solutions with the appropriate terms weak, strong, concentrated and dilute

Describe the difference between a strong and weak acid in terms of an equilibrium between intact molecules and its ions

Note: In exams, define concentration and strength if used in question.

A strong acid is one in which all acid present in solution has ionised to hydrogen ions (no degrees of strength), no equilibrium is formed. A weak acid is one in which only some of acid molecules present in solution have ionised to form hydrogen ions, forming an equilibrium between intact molecules and ions. The fraction of molecules ionised is called the degree of ionisation (concentration of H+/concentration of acid originally).

Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids

Prac – Relative strength of acidsAim: To compare the relative strengths of different acids using a variety of methods

Method:1). 50mL of 0.1M hydrochloric, acetic, oxalic, citric and sulfuric acid were prepared in labelled 250 mL beakers2). A pH meter was used in each beaker and the reading recorded3). A pH strip was placed into each beaker for a short period and its colour compared with a chart4). A few drops of universal indicator were dropped into each beaker and the colour compared to a colour chart

Results:#####

The strongest was sulfuric acid. It is a strong acid and is diprotic, meaning that the concentration of H3O+ ions is twice the concentration of hydrochloric acid. HCl is strong but monoprotic, meaning concentration is identical to HCl concentration. Oxalic acid is diprotic and citric is triprotic, but both are weak and do not completely ionise. The tendency for their conjugate bases to re-bond with hydrogen ions limits the concentration of H3O+ in solution. Acetic is the weakest, being monoprotic and have a low degree of ionisation.

Oxalic – C2H2O4

Citric – C6H8O7

Acetic/ethanoic – CH3COOH

Safety:HCl - corrosive, vapour can burn mouth, throat and eyes Oxalic acid – corrosive to tissue, corrosive to respiratory tract if inhaledWear safety glasses, goggles, use lower concentrations and smaller amounts

Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules

An acid is stronger than another if it has a higher degree of ionisation.


Citric is a triprotic acid, acetic is monoprotic. The degrees of ionisation are:HCl – 0.010/0.01 = 1Citric acid – 2.74 x 10-3/0.01 = 0.274Acetic acid – 4.17 x 10-4 / 0.01 = 0.0417

HCl is the strongest acid and has the highest degree of ionisation.

Use available evidence to model the molecule nature of acids and simulate the ionisation of strong and weak acids

Gather and process information to explain the use of acids as food additives

Acids are added to food to:o Improve taste – e.g. carbonic acid in soft drinks, acetic acid in vinegaro Preserve food – increases acidity to point where bacteria can no longer survive e.g.

coating freshly cut fruit with citric acido Increase nutritional value – e.g. adding ascorbic acid (Vitamin C, an antioxidant)

Identify examples of naturally occurring acids and bases and their chemical compositionAcids:

o Ascorbic acid C6H8O6 – occurs widely in fruit and vegetables, essential to healtho Citric acid – found in citrus fruitso Lactic acid CH3CH(OH)CO2H – produced by anaerobic respiration in cells, found in

muscle tissue and milko HCl – stomach to break down food

Bases:o Ammonia NH3 – result of decomposition of proteins, or anaerobic decay of organic

matter found in fish urineo Carbonates – e.g. calcium carbonate (limestone), magnesium carbonateo Metallic oxides – e.g. Iron (III) oxide, copper oxide and titanium oxide, found in

minerals. Metals are extracted from them.

Calculate pH of strong acids given appropriate hydrogen ion concentrations

Outline the historical development of ideas about acids including those of:o Lavoisiero Davyo Arrhenius

o Lavoisier (1780) acids were substances that contained oxygen Disproved since some oxygen-containing compounds such as metallic oxides were

basic, and distinctly acidic substances such as hydrochloric acid contained no oxygen Wrong but stimulated research

o Davy (1815) suggested that acids were substances that contained replaceable hydrogen.


Bases were substances that reacted with acids to form salt and water. These definitions worked well for most of that century, but the definition made no

attempt to interpret the properties, only classify the substanceso Arrhenius (1884)

Interpreted acidic properties in terms of ionisation to form H+, and weak/strong in terms of degrees of ionisation

the conductivity of acid solutions and their reaction with many metals to form hydrogen gas evidenced that acidic solutions contained hydrogen ions

acids were substances that ionised in solution to produce hydrogen ions A base is a substance that in solution produced hydroxide ions He defined strong acids as those that ionised completely and weak as those that

partially ionised. General equations:HA(aq) H+

(aq) + A-(aq)

XOH(aq) X+(aq) + OH-

(aq)

Weaknesses were: Does not take into account role of solvent in ionisation of acid (ionisation

results from reaction of acid with solvent) Acid-base reactions can occur in solvents where there is no ionisation Not all acidic/basic substances (e.g. metallic oxides) ionised to produce

hydrogen/hydroxide ions

Outline the Bronsted-Lowry theory of acids and bases

An acid is a proton donor, a base is a proton acceptor. Gives the broadest definition of acid/base theory (it means that acids must have hydrogen). This definition :

o Does not restrict bases to those which ionise to produce hydroxide ions, such as in the case of metal oxides and ammonia

o Explains how neutralisation reactions don’t require dissolution of ions into aqueous solution e.g. NH3 + HCl in benzene (direct proton transfer)

o Exchange of proton relies on relative properties of both substances involved, accounting for the role of the solvent

o Shows that hydrolysis of salts to change pH were acid or base reactions o Provided basis for quantitative treatment of acid-base equilibria and pH calculations

Trace developments in understanding and describing acid/base reactions

?

Describe the relationship between an acid and its conjugate base and a base and its conjugate acid

When an acid loses a proton, the resulting ion is called a conjugate base. If the acid is not strong, this conjugate base can re-take a proton to reform the acid, resulting in an equilibrium reaction. Vice versa for bases

Identify a range of salts which form acidic, basic, or neutral solutions and explain their nature

A salt is an ionic compound containing a cation not H+ and an anion not O2- or OH-. In aqueous solution, salts completely dissociate into ions.

Type Acidic Basic NeutralSalt Ammonium nitrate

(NH4NO3)Sodium hydrogen

Sodium acetate (NaCH3COO)Potassium nitrite

Sodium chloridePotassium nitrate (KNO3)


sulfate (NaHSO4)Anything containing Al3+, Fe3+, HSO4

- or H2PO4

-

(KNO2)Sodium carbonate (Na2CO3)Anything containing F-, S2- etc.

Sodium sulfate (Na2SO4)

The pH of a salt solution depends on the nature of its ions, many cations/anions serve as acids or bases. Some generalisations:

o Neutral salts have anions which are the conjugate base of strong acids, and cations the conjugate acid of strong bases, since their reaction to accept/give protons is negligible

o Basic anions react with water to form hydroxide ions in solution. Reaction is equilibrium, occurs to small extent since conjugate acid is stronger than water, and conjugate base is stronger than basic anion

o Acid anions contain hydrogen atoms to react with water to form hydronium ions, derived from polyprotic acids. The anion resulting from hydrolysis of polyprotic acids is amphiprotic, and whether it is acidic or basic depends on the tendency for one hydrolysis reaction (proton donation or proton accept) to occur over the other

Some examples:(Basic anions)S2–(aq) + H2O(l) HS–(aq) + OH–(aq)

F–(aq) + H2O(l) HF(aq) + OH–(aq)

(Acidic cations)[Fe(H2O)6]3+(aq) + H2O(l) [Fe(OH)(H2O)5]2+(aq) + H3O+(aq)

Perform an investigation to identify the pH of a range of salt solutions

Self explanatory. You could use NaCl + KOH for neutral, sodium bicarbonate (NaHCO3) and sodium acetate (NaCH3COO) for basic, ammonium chloride (NH4Cl) for acidic.

Risk analysis:

Hazard Risk ControlAmmonium chloride Released vapour causes

coughing, shortness of breath

Wear safety gogglesUse small amounts to minimise vapour

Identify conjugate acid/base pairs

Acid Base Conjugate base Conjugate acidHCl Cl-

H2SO4 HSO4-,(and SO4

2-?)HNO3 NO3

-

NH4+ NH3

OH- H2OCN- HCNCO3

2- HCO3-

CH3COO- CH3OOH

Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions


Note: Hydrolysis is when a substance reacts with water

Amphiprotic substances can act as both a proton donor and proton acceptor. They react to both accept protons and donate protons. Their behaviour changes whether in aqueous solution or alkaline/acid solution.

e.g. HCO3- (hydrogen carbonate)

In aqueous solution:

HCO3-(aq) + H2O(l) H3O+

(aq) + CO32-

(aq)

HCO3-(aq) + H2O(l) H2CO3(aq) + OH-

(aq)

In basic/acidic solution:HCO3

-(aq or s) + OH- H2O(l) + CO3

2-(aq)

HCO3-(aq or s) + H+

(aq) H2CO3(aq)

Reactions go to completion since products cannot perform the reverse reaction (???). This also applies to HSO3

- (hydrogen sulfite) and HSO4-. Water is amphiprotic.

Identify neutralisation as a proton transfer reaction which is exothermic

Neutralisation reactions are proton transfer reactions, and involve the reaction between an acid and a base. They are exothermic and thus have a negative enthalpy change. The net ionic reaction in Arrhenius theory is:

OH-(aq) + H+

(aq) H2O(l)

Acids and bases not fitting in Arrhenius theory do not necessarily produce salt and water during neutralisation reactions e.g. neutralisation of ammonia. In LB theory, an acid and base react to form conjugate base and conjugate acid. The acid gives a proton to the base. Reactions between strong acids and bases form very weak conjugate acids/bases and go virtually to completion since the back reaction has almost no tendency to occur. Otherwise, reactions are equilibria.

Describe the correct technique for conducting titrations and preparation of standard solutions

Volumetric analysis is a form of chemical analysis where the concentration of a substance is determined.

Determining the composition of a solution require titration against another solution of known concentration, called the standard solution. The substance dissolved is a primary standard.

Equipment:

A primary standard:o Must be obtainable in very pure form and have known formula o Should not alter weight unintentionally during preparation/titration e.g. absorbing

moisture from airo Have a reasonably high formula mass to minimise weighting errorso Purified by drying in oven and cooling in dessicator to eliminate moisture and prevent its

absorptione.g. oxalic acid, sodium carbonate


Use of equipment:o Pipette – solution to be used is first drawn in above mark, then solution let out until

meniscus at mark, solution let out through gravity with tip against wall of containero Burette – first, rinse with portion of solution to be dispensed, overfilled then excess

allowed to run out

Preparation:o Accurately measure mass of primary standard e.g. electronic beam balanceo Rinse a volumetric flask and beaker with distilled watero Pour the primary standard into beaker and dissolve with distilled water, less than

intended volume of final solutiono Pour into volumetric flask, and repeat a few more timeso Use a pipette to add final few drops to complete solution

Titration curves:Strong acid strong base:

Equivalence point at pH 7, steep curve. Indicator used should have colour endpoint near equivalence point. Using indicator changing during equivalence point is inaccurate, too difficult to tell exact colour shade needed

Strong base/weak acide.g. NaOH, acetic acidEquivalence point in basic range since salt formed is basic, the anion is a conjugate base of weak acid and thus a weak base. Equilibrium reaction occurs, weak base reacts to reform conjugate acid, thus decreasing acidity since less H3O+

Strong acid/weak baseSimilar to above with equivalence point acid. Special case when CO2 formed during reaction e.g. HCl + Na2CO3, CO2 forms carbonic acid.

Weak/weakNot good since gradient around equivalence point is quite shallow, big volume difference between indicator endpoints and equivalence points, needs to fit indicator very well or use one which changes during equivalence, hard to distinguish.

Perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases


Prac – TitrationAim: To standardise HCl and NaOH solutions using titration

Method:Preparing standard:1). A clean 250mL beaker was placed on an electronic balance, zeroed, and had 2.650g of Na2CO3 added using a spatula2). Approx 100 mL of distilled water was added to the beaker, and solution stirred using stirring rod3). A 250mL volumetric flask was rinsed with distilled water*

4). The Na2CO3 solution was poured into the volumetric flask, and step 2 repeated5). The stirring rod/beaker were thoroughly washed using a wash bottle, the runoff dripping into the vol. flask6). Using a 25mL pipette, the volumetric flask was filled to the 250mL mark7). A stopper was placed on the flask and contents swirled to mix8). The pipette was rinsed with the unknown HCl*

9). 10 mL of unknown concentration HCl was poured using a pipette into a 50mL beaker washed with distilled water*, and a few drops of methyl orange added10). A burette was washed and filled with the Na2CO3 standard solution11). The standard solution was quickly drained into the beaker to find an approximate end point12). Steps 8-10 were repeated 3 times but more accurately*to ensure no cross-contaminationSee book for calculations

Qualitatively describe the effect of buffers with reference to a specific example in a natural system

Buffer solutions resist changes in pH. It contains comparable amounts of a weak acid/base and its conjugate base/acid. Take for example an acetic acid (CH3COOH) and sodium acetate (NaCH3COO) system (acidic buffer).

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO- (1)

Addition of sodium acetate would increase the concentration of CH3COO – ions on the right. The equilibrium shifts to the left, but due to the unchanged concentration of H3O+ ions, it stays enough to the right for dissociation to cause a net increase in CH3COO – ions. This net increase enhances buffering capacity. When hydronium ions are added to solution, the equation will shift to the left according to Le Chatelier’s principle to reduce the concentration of H3O+ ions. When hydroxide ions are added, the CH3COOH will react to form water and CH3COO-, reducing OH- concentration. In both cases, pH change is reduced.

Buffer in natural systemCarbonic acid / bicarbonate ion buffer system in mammalian blood:

H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+

(aq)

Maintains the blood at around 7.4 for optimum function, too high/low can result in death.

Analyse … to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills

o Many acids and bases are corrosive, can damage materials if spiltNeutralisation reactions can:

o Reduce or nullify corrosive properties of spill, minimising damage


o Utilise common, cheap, safely handled/stored materials and produces relatively harmless products e.g. sodium bicarbonate NaHCO3

Sodium bicarbonate is amphiprotic, so it can be used for both acidic and basic spills:

Vinegar – commonly used in cooking, contains acetic acid:

o Degree of reaction can be controlled by using different amounts of neutralising substance, excessive amounts are wasteful and some

Overall, it is a useful, convenient and safe technique if used appropriately

Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds

Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures

Strong intermolecular forces – higher BP/MP compared to similar mole mass (roughly similar dispersion forces)

-OH

carboxylic acid group

Forms One hydrogen bond between molecules

Two polar bonds between molecules (C-O polar and C – O – H hydrogen bond),

higher boil/melt pointNone ionised, neutral Small no. ionised, acidic

Identify the IUPAC nomenclature for describing the esters produced by reactions of straight-chained alkanoic acids from CI to C8 and straight-chained primary alkanols from C1 to C8

Alkyl (e.g. methyl) alkanoate (e.g. formate, acetate, propanoate etc.)

Identify esterification as the reaction between an acid and an alkanol and describe, using equations, examples of esterification

Esters – carboxylic acids combined with alcohols, equilibrium reaction

Standard equation:

o H2O molecule released (H of alkanol OH of acid)o OR’ (R’ dummy variable) of alkanol C of acid

e.g. (acetic acid + methanol)

(formic acid + ethanol) Chem notes David Lee BHHS 2007 44

Describe the purpose of using acid in esterification for catalysis

o The acid acts as a dehydration agent, removing water from the reaction (Le chatelier argument), thus increasing yield

o Acts as a catalyst, lowering activation energy to speed reactiono Only small amounts of acid required o Most common is sulfuric, others include tosic, scandium (III) triflate

Explain the need for refluxing during esterification

Reflux – the backflow of reactants into the reaction vesselo Reactants and ester products can be volatileo Reaction is carried out at high temperatures to speed reaction, causing evaporation of

alcohol and ester and thus loss of reactants/productso Vapour is also dangerous as it is flammable and toxico To avoid loss and prevent diffusion, a cooled condenser is placed over the reaction

vessel, covering ito The vapour condenses here and runs back into the reaction vessel, which

Increases yield and saves on resources Allows the reaction to be carried at higher temperatures (faster) Prevents flammable gases from escaping

Outlines some examples of the occurrence, production and uses of esters Identify and describe the use of esters as flavours and perfumes in processed foods and

cosmetics

Esters occur naturally, and are identified as fragrances and flavours in fruit and flowers e.g. orange (octyl acetate). Animal fats such as butter, or oils such as linseed are also esters, as are waxes.These fats and oils can be used to make soap

(Fat or oil) + sodium hydroxide → Salt of carboxylic acid + glycerol

The salt of carboxylic acid is the cleaning factor of soapArtificially produced esters:Aspirin:

o Acetylsalicylic acid active ingredient, discovered by Hoffman in 1897o

Salycilic acid + acetic anhydrideo Most widely used pain-relieving drug, e.g. headache, and prevents blood clots

Ethyl acetate:o Solvent in industry, nail polish removero Acetic acid and ethanol

C2H5OH + CH3COOH → C4H8O2 + H2O

Octyl acetate:o As a food flavouring such as in sweets, ice cream etc.o Acetic acid and octanol

C8H17OH + CH3COOH → C10H20O2 + H2O

Perform a first-hand investigation to prepare an ester using reflux

Prac – EsterificationChem notes David Lee BHHS 2007 45

1). 15 mL of acetic acid, 15 mL of butan-2-ol and 10 drops of conc. Sulfuric acid was placed into a 100 mL round-bottom flask2). The apparatus was set up as shown3). The bunsen burner was lit and reflux allowed to occur for about 15 minutes until two layers clearly visible4). A separating funnel was assembled on a retort stand and 100 mL of distilled water poured inside5). Contents of round-bottom flask were transferred into funnel and the mixture shaken6). 50mL of sodium carbonate solution was added and gently shaken, occasionally inverted and tap opened to release gas7). Once the layers separated, the lower layer was discarded8). Step 6-7 was repeated9). The remaining liquid was poured into a conical flask and a teaspoon of CaCl2 added, shaken gently then allowed to stand for a few minutes

CH3COOH(l) + C4H9OH(l) (H2SO4 conc.) CH3COOCH2CH2CH2CH3(l) + H2O(l)

(diagram representation here)

H2SO4 (see above)Na2CO3: neutralise remaining acid, excess forms distinct layer so easily separatedCaCl2: anhydrous, absorbs water in mixture to leave ester

Safety:1). Acetic acid + butan-2-ol flammable, apparatus should be secured to ensure no tipping; tipping could pose fire risk due to contact with bunsen flame2). Acetic acid corrosive to skin, avoid spilling3). Condenser should have adequate water to prevent organic vapour escaping, flammable and respiratory irritant

Chemical Monitoring and Management

Outline the role of a chemist employed in a named industry or enterprise, identifying the branch of chemistry undertaken by the chemist and explaining a chemical principle that the chemist uses

Industry: Australian chemical manufacturing companyBranch: Analytical chemistry – quantitative and qualitative analysis of substances present within materialsRoles:

o Occasional monitoring of ethylene quality, waste water (pH, suspended solids, hydrocarbons etc.), gaseous emissions (particulates, toxic pollutants) to ensure reliability of results by other chemists

o Check proper operation of equipment/calibrate instrumentso Train shift workers to use instrumentso Look for ways of improving processes

Principle: Adsorption in gas-solid chromatographyo Gas chromatography – a liquid or gaseous mixture is vaporised into a stream of helium

flowing over a stationary phase such as a solido If stationary phase is solid, the components of injected mixture adsorb (stick onto) its

surfaces to differing extents, and desorb at different rateso This causes different substances to pass through at different rates


o A detector is able to quantitatively measure each substances as it passes outo Can be used to determine chemical composition of substances

Gather, process and present information from secondary sources about the work of practising scientists identifying:

o The variety of chemical occupationso A specific chemical occupation for a more detailed study

Many areas a chemist can work in, 13 divisions recognised by Royal Australian Chemical institute including:

o Environmental chemistry (detailed) – determining how substances interact in the environment, monitoring concentrations of substances especially in air, water and soil

Environment monitoring, employed by Environmental Protection Authority, mining companies, local government – qualification can be BSc and postgrad qualification in fields such as scientific communication/management. Could collect data on air/water quality, then analyse and assess this information. Require strength in chemical analysis, and instrumental analysis. May work in a team, providing environmental advice to external bodies via reports.

o Physical chemistry – study and measurements of physical aspects of compounds and reactions e.g. reaction rates, structure of substances, nature of chemical bonding

o Pharmaceutical chemistry – discovery, testing, synthesis and commercial development of chemicals for use as medicines

o Industrial chemistry – chemistry of industrial processes such as manufacture of ammonia, sulfuric/nitric acids and others

Identify the need for collaboration between chemists as they collect and analyse data

o Chemistry is diverse, chemists specialise in particular branches as range of knowledge too large

o Some real-world problems require expertise from more than one brancho Collaboration is required for proper tackling of problemo Also, work of one chemist may have implications in another area, requires active

communication skillso This facilitates efficiency of scientific progress and scientific worko Can be achieved by:

Publishing of papers Collaboration between laboratories Direct voice communication

Describe an example of a chemical reaction such as combustion, where reactants form different products under different conditions and thus would need monitoring

e.g. Petrol combustion in motor vehicles (mostly octane)Different situations/products:

o Complete combustion – excess oxygen, products carbon dioxide and water

o Incomplete combustion – insufficient oxygen, products carbon monoxide, carbon dioxide, soot, unburnt hydrocarbons, water. Ensure adequate oxygen supply to fuel

o Nitrogen oxides – reaction of oxygen with atmospheric nitrogen due to high temps forms nitric oxide and nitrogen dioxide (see acidic environment). Rhodium-platinum catalyst converts more polluting gases into less harmful ones


o Sulfur oxides – some sulfur compounds in fuels:

Monitoring can ensure minimum possible toxic chemicals released, important since:o CO affects judgement/perception as levels as low as 10 ppm, can cause death by

asphyxiationo Soot contributes to particulate pollution, bad for asthma suffererso Nitric oxide affects respiratory systems and is generally toxic, excessive production by

motor vehicles can affect health of population o Sulfur oxides and nitric oxides contribute to acid raino Motor industry can use information to build more efficient engines (more complete

combustion)

Identify and describe the industrial uses of ammonia

o Fertilisers – reacted with sulfuric acid to form ammonium sulfate or nitric acid to form ammonium nitrate fertiliser. Application to soil provides good source of nitrates essential for crop growth, improving yields. 2NH3 + H2SO4 (NH4)2SO4(s)

o Conversion to nitric acid – nitric acid is used in making explosives, dyes, fibres and plastics

o Neutralisation of acid – petroleum industry uses ammonia to neutralise acid components of crude oil and protect equipment from corrosion

o Water treatment – addition of ammonia and chlorine to water produces more stable disinfecting residual than chlorine alone

Identify that ammonia can be synthesised from its component gases, nitrogen and hydrogen

Describe that synthesis of ammonia occurs as a reversible reaction that will reach equilibrium

Identify the reaction of hydrogen with nitrogen as exothermic Explain why the rate of reaction is increased by higher temperatures Explain why the yield of product in the Haber process is reduced at higher temperatures

using Le Chatelier’s principle Analyse the impact of increased pressure on the system involved in the Haber process Explain that the use of a catalyst will lower the reaction temperature required and identify

the catalyst(s) used in the Haber process

Nitrogen and hydrogen combine to form ammonia, which in turn decomposes to reform reactants (reversible reaction). Equilibrium reached when rate of forward/reverse reactions the same.At higher temperatures, the average kinetic energy of the reactants is higher. Thus:

1. Larger fraction of molecules have adequate energy to overcome activation energy and react upon collision

2. Molecules move faster, more collisions between moleculesThese increase the rate of reaction, and apply to both forward and backward reactions. However, it affects forward reaction more.


Reaction is exothermic, higher temperatures shifts equilibrium to left to reduce temperature change.

Number of moles of gas on each side of reaction is different. According to Le Chatelier’s principle, increasing pressure shifts equilibrium to right since there are less moles of gas, decreasing pressure and thus minimising pressure change. Vice versa

Catalysts: iron-iron catalyst, small amounts of K2O, Al2O3

o Hydrogen and nitrogen molecules adsorbed onto surface, increasing collisionso Allows reaction via new chemical pathways with lower Ea

o Reduces activation energy, allowing molecules with lower kinetic energy to react, thus lowering the required temperature

o Reaction is faster, equilibrium unaffected

Explain why the Haber process is based on a delicate balancing act involving reaction energy, reaction rate and equilibrium

Balancing these factors to maximise yield and reaction rate is required to maintain adequate production and use of resources:

o Higher temperatures means faster reaction rate but lower yield (equilibrium more to reactants), and higher energy costs associated with temperature maintenance

o The reverse is true for lower temperatures, their advantages conflict, optimal temperature currently used is 700K

o Higher pressures increases yield but places more stress on reaction vessel, current optimum 2.5 x 104 kPa

o Catalysts can speed up reaction rate and lower reaction energy, lowering temperature required for same production rate

Explain why monitoring of the reaction vessel used in the Haber process is crucial and discuss the monitoring required

o Temperature and pressure of reaction vessel – keep within range for optimum production rate, excess temperatures can damage catalyst and lower yield

o Ratio of incoming reactants – maintain stoichiometric ratio and prevent buildup of one reactant, slowing reaction

o Impurities in incoming gases – O2 can cause explosion, CO/CO2 can poison catalyst and reduce its lifespan

o Rate of ammonia removal – inadequate rate of removal shifts equilibrium to reactants, reducing yield

Describe the conditions under which Haber developed the industrial synthesis of ammonia and evaluate its significance in that time in world history

o Haber developed process in 1908, before WW1o Growing population in early 20th century required large amounts of fertiliser to feed

populationo Growing militancy of Germany required more ammonia for explosiveso Haber process able to meet these demandso Germany originally obtained nitrates as saltpere/guano from Chile, but advent of WW I

caused allies to set up a naval blockade, preventing imports, but Haber process allowed Germany to be self-reliant

o This prolonged their resistance against the Allies, increasing the length of World War 1 and resulting in loss of many more lives

o Significant impact in world historyChem notes David Lee BHHS 2007 49

Deduce ions present in a sample from a range of tests (Maybe need mixtures – go to conquering chem for good summary)

Cations:Cations:Confirmation tests:Ba2+ - apple green flame

Ca2+ - brick red flame

Cu2+ - blue-green flame, dissolves in ammonia to form deep blue solution

Fe2+ - decolourises acidified dilute KMnO4 solution

Fe3+ - deep red solution with SCN-

* Fe2+ can form white or green which may tun brown due to oxidisation over

time (older samples)* Copper forms CuI with iodine, not used in initial procedureNote: Silver chloride is also a white precipitate, but silver sulfate is insoluble and colourlessAnions:

HNO3 Ba(NO3)2 Pb(NO3)2 AgNO3 OtherCarbonate

CO32-

Bubbles1 White precipitate soluble in

HNO3

- White precipitate, soluble in

HNO3

pH between 8 and 11

Sulfate SO4

2-- (acidified)*

thick white precipitate

(acidified) *white

precipitate

- -

Phosphate PO4

3-- (ammonia)*

white precipitate soluble in

HNO3

- Yellow precipitate soluble in

HNO3

(acidified) (NH4)2MoO4,2

yellow precipitate (may need warming)

Chloride Cl- - - - (acidified) white

precipitate

dissolves in ammonia, darkens in sunlight

1 CO32-

(aq) + 2H+(aq) CO2(g) + H2O(l)

* Sulfate weaker lewis base than phosphate SO4

2-(aq) + H3O+ HSO4

-(aq) + H2O(l), equilibrium enough to left so adequate sulfate to produce

noticeable precipitation with barium/lead (CHECK – how would this effect lead?)In basic conditions, enough phosphate for precipitation2 12(NH4)3MoO4 + PO4

3- + 3H+ (NH4)3PMo12O40

Perform first-hand investigations to identify the following ions:o Phosphateo Sulfateo Carbonate


o Chlorideo Bariumo Calciumo Leado Coppero Iron

Some safety information:o Barium compounds are toxic (see barium chloride below) (barium sulfate is mostly

harmless)o Silver chloride – eye, skin and respiratory irritanto Lead nitrate – poisonous if swallowed, causing spasms, nausea etc. Can be absorbed

through skin to cause irritation, redness over short periods

Describe and explain the evidence for the need to monitor levels of one of the above ions in substances used in society

Copper is an essential trace element for many organisms, such as humans. However, excess amounts can be toxic and detrimental to the environment:

Sources of Copper Effects on People Effects on aquatic ecosystems

Corrosion of copper plumbing

Over consumption – Liver, kidney damage (liver is storage point for copper)

(Chuttani et al., 1965)

Kills plants as well as algae (bioaccumulation)

Copper sulfate crystals to control algal growth

Over consumption - Vomiting, diarrhea, nausea

(USEPA, 1980)

Reduces survivability of aquatic invertebrates

(WSDOE, 1992)Contact with skin –

eczema, edema of eyelids (Patty, 1963)

Reduces survivability of fish (Holland et al, 1960)

Death (NRC, 1977)

Perform first hand investigations to measure the sulfate content of lawn fertiliser and explain the chemistry involved

Analyse information to evaluate the results of the above investigation and to propose solutions to problems encountered in the procedure

Hazard Risk ControlHydrochloric acid (if used) Corrosive. Can cause

permanent eye damageKills tissue

Wear safety glasses, use lower concentrations

Barium Chloride Eye, skin and respiratory irritant


Hot plate (if used) Can cause burns upon contact

Gloves, keep body parts away

Prac – Sulfate content of fertiliser1). 1.0g of ammonium sulfate was placed into a measuring cylinder, and 2 ml of water was added; mixture was shaken to dissolve2). 10 mL of 1.0M barium chloride solution added to form white precipitate


3). This precipitate was allowed to settle overnight, then the clear liquid decanted4). 10mL of water was added to precipitate to further dissolve chloride and ammonium ions, mixture shaken 5). Steps 3-4 repeated twice6). Two pieces of filter paper were weighed and then shaped into a cone7). The barium sulfate slurry was evenly poured into this cone, then the cone placed on a source of heat for an hour to evaporate remaining water8). Barium sulfate + filter paper was weighed, results recorded and analysed

ResultsMass of ammonium sulfate: 1.000 Mass of BaSO4 precipitate: 0.655 g

Mass of SO42- =

27% sulfate ions (w/w) in fertiliser

EvaluationThe sources of error are:1). Loss of barium sulfate due to slight solubility – using increased concentrations of reactants would decrease its effect on results, but increases contamination by adsorption. Use between 0.005 and 0.05 mol/L of initial sulfate concentration 2). Adherence of barium sulfate to walls of the measuring cylinder3). Incomplete drying of precipitate, contains water while weighed – repeated cycles of drying/cooling/weighing until constant mass obtained4). Remaining chloride and ammonium ions not completely removed by washing – more repetitions, although this would cause greater loss of BaSO4 due to dissolution5). Contamination by adsorption of substances in solution during precipitation – form precipitate slowly by slowly mixing reactants, and forming in hot solutions maximises particle size and reduces adsorption6). Small size of barium sulfate crystals – some may have fallen through double filter paper, use of a sintered glass funnel to trap it would result in more accurate resultsORWeighed amount of agar added as coagulantORForm in hot solution to maximise particle size + small amount of HCl

Describe the use of AAS in detecting concentrations of metal ions in solutions and assess its impact on scientific understanding of the effects of trace elements

Interpret secondary data from AAS measurements and evaluate the effectiveness of this in pollution control

Atomic absorption spectroscopy – used to measure low concentrations of elements in ppm range, mainly metals. Each element has unique emission spectrum, so by measuring, studying and using spectra we can determine qualitatively and quantitatively the elements present in sample (by looking at spectra, measuring intensity)

Practical arrangement + workings:o Measures how much light of a specific wavelength is absorbed by sample being studiedo Sample to be analysed fed into flame, vaporises and converts molecules and ions into

atoms


o Atomic emission lamp producing a specific emission spectrum matching element to be studied is passed through the flame

o Electrons absorb energy and are excited to higher energy levelso Light is passed through prism to concentrate light of desired wavelength onto detectoro Wavelengths characteristically absorbed by element in question shows obvious drops,

indicating absorbanceo Absorbance linearly proportional to concentrationo This value is compared to a function (absorbance vs concentration) produced by

measuring known concentrations to find conc.

Why useful:o Relies on absorption rather than emission, nearly 100% atoms in ground state absorb as

opposed to <0.1% excited which emit, can measure down to ppbo Can measure concentration of only one element at a time

Uses:o Detection of heavy metals e.g. copper, aluminium in waterways, since only found in trace

quantitieso Concentrations of micronutrients in soilso Small amounts of contaminants in medicines and foods e.g. mercury, leado Elements in organisms e.g. blood and urine samples for mineral deficiencies

Trace element: Element present in concentrations < 100 ppm

Impact on scientific understanding:o Allowed scientists to detect minute concentrations of trace elements, allowing them to

recognise its importance e.g. legumes were unable to grow in arid parts of Victoria until AAS tests showed molybdenum deficiency

o Helped to demonstrate the importance of trace elements in living organisms, such as in maintaining enzyme function e.g.

Evaluation on pollution control:Useful since:

o Allows very small ppm concentrations of elements to be measured, (up to 0.01 ppm), allowing detection of pollutants which may still be harmful in trace amounts

Describe the composition and layered structure of the atmosphere

Layer of gas about 200 – 300 km thick surrounding the earth. Many different gases and distinct layers which different characteristics:

Overall gas composition: 78% N2, 21% O2, 0.93% argon, trace amounts of CO2, neon, methane etc.Layer Altitude above

surface (km)Most common gases Description

Troposphere 0 – 15 N2, O2, H2O, CO2, Ar Contains most of earth’s gases, organisms inhabit this zone,

weather eventsStratosphere 15 – 50 N2, O2, O3, Contains ozone layer (25 km),

temperature increases with altitude and gives stability

Mesosphere 50 – 80 Coldest layer (down to -100 celsius)


Thermosphere/Ionosphere

>80 Ions (O2+, NO+), O Temp rises with altitude, ionic

and atomic gas particles, important in radio

communications since radio waves reflect off

Identify the main pollutants found in the lower atmosphere and their sources

Pollutant SourceCO Burning fossil fuels, forest fires

Airborne lead Lead smelters, leaded fuelsCFC’s Foaming agent, refrigerant-air conditioner coolant,

propellantSO2 Combustion (fuel impurities), metal extraction from

sulfide ores, chemical manufacturingOxides of nitrogen (NO +

NO2)Combustion (vehicles and power stations)

Particulates Combustion, bush fires, industrial processes such as mining

Describe ozone as a molecule able to act both as an upper atmosphere UV radiation shield and a lower atmosphere pollutant

Allotrope – a different physical form of the same element in the same phase Ozone is:

o An allotrope of element oxygeno Naturally present in atmosphere; only 0.02 ppm at ground level in clean air, 2 – 8 ppm in

stratosphereo Detrimental in lower atmosphere – poisonous to many organisms; causes breathing

difficulties, fatigue and headache in humanso Beneficial in upper atmosphere - filters our short wavelength UV light which can damage

living tissue

Produced in upper atmosphere through UV light:3O2(g) -- (UV light) 2O3(g)

Describe the formation of a coordinate covalent bond Demonstrate the formation of coordinate covalent bonds using Lewis electron dot

structures

Coordinate covalent bonds occur when shared the electrons come from one atom. Once formed, identical to regular covalent bond.

Compare the properties of the oxygen allotropes O2 and O3 and account for them on the basis of molecule structure and bonding

Property O2 O3 ReasonBoiling point (oC)

-193 -111 Polar bonds between molecules means

intermolecular forces


stronger (O3)Odour None Sharp, irritatingColour None Pale blueDensity Slightly denser than

air1.5 times denser than

airOne more O atom per

moleculeReactivity Highly reactive with

many metals and non-metals

Very highly reactive, attacks double bonds

on alkenes

The O-O bonds in O3

are less strong than the double covalent bond

in oxygenSolubility in water

Sparingly soluble More soluble than O2 O2 is non-polar, O3 is bent and thus is polar

Oxidation ability

Lower Higher O involved only in coordinate covalent has greater electron

affinity

Compare the properties of the gaseous forms of oxygen and the oxygen free radical

Free radical – a neutral species with an unpaired electron which can be formed by splitting a molecule into two neutral fragments

Property O2 OReactivity Less reactive (full outer

valence shell)Very reactive (unpaired electron)

Oxidation ability

Lower Higher (unpaired electron, high tendency to take electrons to

complete valence shell)

Identify the origins of CFCs and halons in the atmosphere

o CFCs (contain chlorine, fluorine and carbon) – developed as refrigerant in 1930’s to replace ammonia, also used as propellant, solvent and foam blowing agent. Through use, gas released into atmosphere

o Halons (carbon and halogens) – were used in fire extinguishers, recently use has been drastically reduced

Identify and name examples of isomers (excluding geometric and optical) of haloalkanes up to eight carbon atoms

Model isomers of haloalkanes using model kits

Haloalkane – hydrocarbon with one or more hydrogens replaced by halogen atoms (encompass CFC’s)

Discuss the problems associated with the use of CFC’s and assess the effectiveness of steps taken to alleviate the problem

Write the equations to show the reactions involving CFC’s and ozone to demonstrate the removal of ozone from the atmosphere

Identify alternative chemicals used to replace CFC’s and evaluate the effectiveness of their use as a replacement for CFC’s

Problems:Removal of stratospheric ozone


o CFC’s released into the troposphere are not washed out by rain (non-soluble) and not destroyed by sunlight/oxygen at low altitudes

o Diffuse into the stratosphere and short wavelength UV breaks a chlorine off e.g.

o Chlorine reacts with ozone

o ClO reacts with free oxygen atoms

o Net result is conversion of O3 and O to two O2

o Chlorine molecule unchanged at end, can continue to react and remove ozoneo This occurs on average a few thousand time before chlorine radical reacts with another

chemical which removes it e.g. methane

The problem:o Removal of stratospheric ozone reduces filtering of short length UV radiation, meaning:

Increased incidence of sunburn and skin cancer Increased damage e.g. brittleness to synthetic materials such as PVC Increased risk of eye cataracts Reduced plant growth for some species (e.g. rice) due to UV interference

with photosynthesis mechanismso CFC’s are greenhouse gases, and enhance global warming

Alleviation:o Agreements to phase out use of CFC’s (e.g. Montreal Protocol, cease use in developed

by 1996)o Agreements to phase out use of halons by 2010o Assistance to poorer countries to phase out CFC useo Replacement with safer alternatives

HCFC’s – contain C-H bonds decomposable by radicals and atoms in troposphere and are decomposed to significant extend. However, ozone-destroying capacity is still significant (phase out by 2030 – Montreal Protocol). Only useful as temporary substitute

HFC’s – no C-Cl bonds, do not form Cl atoms in atmosphere, no ozone-destroying capacity. Useful as permanent substitute, but more expensive

Air being used to replace as foaming agent

Assessment:o Adherence to agreements will ensure ozone layer returns to pre-CFC state since damage

is reversibleo The pace of CFC withdrawal means it will be many decades before the above happens,

meaning the effects of increased UV radiation will be felto Relies on co-operation of countries, will be less effective if countries withdrawo Replacements such as HFC’s are more expensive than CFC’s, may be a burden on lower

countries until better alternatives found

Overall, it is an effective long-term solution but prolongs problems in the short term. It relies heavily on cooperation of countries and this may be a downfall.

Analyse the information available that indicates changes in atmospheric ozone concentrations, describe the changes observed and explain how this information was obtained


Evidence/information CHECK:o Measurements of total ozone in a column of atmosphere have been conducted since 1957o In springtime of 1980-1984, a severe depletion of ozone above Antarctica was detected

by the British Antarctic surveyo By 1985 it was approximately 30%, and in some places it had been completely destroyedo A net decrease of 3% per decade was recorded for the period 1978-1991, factoring in

natural variations

Method of collection:Data collected by range of ground and airborne instruments

o Dobson spectrophotometer (up to 48 km, groundbased) Developed in 1924, only source of long-term data Can be used to measure both total column ozone and profile ozone, currently

used to calibrate measurements by other methods Measures the intensity of four different wavelengths of UV radiation

reaching it; two are strongly absorbed by ozone, the others are not The ratio between the two intensities is determined and used to calculate total

ozoneDisadvantages/advantages

Strong affected by aerosols and pollutants Measures only over a small area

o LIDAR (10 – 50 km, groundbased) Relies on absorption of laser light by ozone Telescope used to collect UV light scattered by two laser beams, one which is

absorbed by ozone (308 nm) and one which isn’t (351 nm) By comparing these values, a profile of ozone concentration vs altitude can

be measured

o Balloons (up to 40km, airborne) Various instruments can be mounted onto balloons Electrochemical concentration cells measure current produced by chemical

reactions with ozone Photospectroscopy utilises film or electric sensors sensitive to UV light to

measure wavelengths affected by ozoneDisadvantages/advantages

Can provide many days of continuous coverage Inexpensive Unpowered, flight path cannot be controlled

o TOMS (from space) Observes incoming solar energy and backscattered UV radiation at six

different wavelengths Gas molecules in the atmosphere scatter some EM radiation back, while

some is absorbed by ozone By comparing the intensity of backscattered radiation to incoming radiation,

amount of ozone can be obtainedDisadvantages/advantages

Provides global coverage Constant, accurate coverage Coverage in variety of weather and geophysical conditions More expensive than other methods


Identify that water quality can be determined by considering:o Concentrations of common ionso Total dissolved solidso Hardnesso Turbidityo Acidityo Dissolved Oxygen and Biochemical oxygen demand

Water quality is the chemical, physical and biological characteristics of water, with respect to its intended purpose.

o Concentrations of common ions1. Cations include Na+, Mg2+, Ca2+, K+, Al3+ and heavy metals such as Hg2+

, Pb2+, Cd2+

and arsenic. Common anions are Cl-, SO42- and HCO3

-

2. Hardness (e.g. calcium, magnesium), detract from taste and appearance (iron), Choke marine life(Al3+), toxic to humans (Hg2+, Pb2+), promote eutrophication (nitrates, phosphates) which forms undesirable decomposition products e.g. ammonia and detracts from taste/appearance

3. Water percolating through soil or underground estuaries, agricultural runoff4.

a. AAS used to analyse cationsb. Gravimetric analysis used for anions e.g. precipitating Cl- as AgCl

o Total dissolved solids1. Mass of solids dissolved in unit volume of water (ppm)2. High TDS reduces crop and plant growth, >500 ppm is not suitable for human

consumption, >1000 ppm is unsuitable for irrigation3. Underground aquifers, flowing through farming/grazing areas with disturbed soil4.

a. Evaporation, but it produces very small masses of solid, easy to lose through turbulent bubbling/spitting (use large volumes of water, 1 litre round-bottom flask is suitable)

b. Conductivity – most solid dissolved substances are ionic and conduct electricity; can be measured by conductivity meter to determine dissolved solids

o Hardness 1. Concentration of Ca2+ and Mg2+ ions in water (ppm CaCO3)2. Hard water forms a precipitate with soap, reducing cleaning power. Under high

temperatures such as in kettles, Ca2+ forms an insoluble precipitate with sulfate and carbonate ions, reducing kettle efficiency

3. (see below)4.

a. Titration with EDTA, which forms stable complexes with these ions. Indicator Eriochrome Black T. In solution is blue but forms red coloured complex with Mg2+. Endpoint when it turns blue, indicating no more Mg2+ in solution

o Turbidity 1. Measure of suspended solids in water2. Undesirable appearance and taste, reduces sunlight penetration for plant

photosynthesis, can absorb IR and raise water temperature


3. Clay, silt, plankton, industrial wastes4. Secchi disk, visually

o Acidity1. pH of water2. At extreme ranges, it can reduce survivability for aquatic organisms 3. Decomposition of organic matter, acid rain, exposure of sulfide ores in mining,

fertiliser run-off4.

a. Universal indicator solution or paperb. pH meter

o Dissolved Oxygen (ppm), Biochemical oxygen demand1. Concentration of dissolved oxygen in water / capacity of organic matter to

consume oxygen2. Regular levels of O2 (~10 ppm) indicate high quality water, since low levels

indicate high BOD and other factors such as:a. Heat pollution which reduces O2 solubilityb. Excess of organic wastes such as sewerage which take up O2 in aerobic

bacterial decompositionc. Eutrophication by excessive growth of aquatic plants

3. Dissolved from atmosphere, produced by photosynthesis by aquatic plants, algae / Aerobic organisms such as bacteria, fish, worms

4.a. (dissolved oxygen) Titration (see later)b. (BOD) Addition of nutrient to sample and incubating at 20oC in sealed,

air-free container in dark for 5 days, then measuring residual dissolved oxygen. Difference is BOD

Identify factors that affect the concentrations of a range of ions in solution in natural bodies of water such as rivers and oceans

o Pathway from rain to water body – rainwater collects ions before it runs into natural bodies of water Bushland contains small amounts of nitrates and phosphates from natural nutrients on

surface Rainwater soaking into ground collects Ca2+, Mg2+, sulfate and chloride from soil and

rocks it flows through Percolation into deep underground aquifers results in collection of Fe3+, Mn2+ among

otherso Human activity

Removal of natural vegetation or irrigation can increase salinity and thus NaCl in rivers

Agricultural fertilisers contribute nitrates and phosphates through runoff or dumping Discharge of sewage increases nitrates/phosphates, and various ions such as Cl-

Acid rain caused by industry is better able to leech certain cations e.g. Ca2+ and Mg2+ from soil

Motor car emission can increase leado Frequency of rain – more rain means more dissolved ions entering water bodieso Bushfires – bushfires unlock nutrients and ions such as nitrates from plants, water picks

this during runoffo Water temperature – higher water temperature increases evaporation and thus increases

concentrations of all ions in solution


Perform first-hand investigations to use qualitative and quantitative tests to analyse and compare the quality of water samples

Prac – Qualitative analysis1). A sample of catchment water was taken and visually inspected for colour and hydrocarbons (hydrocarbons produce rainbow effect on surface)2). Some water was poured into a small conical flask and shaken for one minute, then observed for bubbles which indicate detergent3). Temperature of the water was recorded with a thermometer4). A pH meter connected to a data logger was used to measure the pH of the solution5). A turbidity tube was lowered into the sample until the bottom disappears, and reading recorded6). Silver nitrate was added, and a white precipitate indicated the presence of chloride ions7). A fresh sample was taken and Heavy metals (See below)8). The above steps were repeated for tap water and distilled water

Results:

Distilled water Tap water Creek waterColour Clear Clear Slight yellowPO4

3- No No NoSO4

2- No No SlightHydrocarbons No No No

Turbidity (NTU) <10 <10 <10Detergents No No Slight

pH 4.5 6 6Heavy metals No No Trace

Cl- No Slight LotsTDS 0 160 560

Bad odour mainly due to decaying vegetation in creek water.Low pH can be due to ammonium fertiliserHigh pH can be due to limestone and dolomiteHigher turbidity means lower intensity and quality of light for photosynthesis, lower DO and thus aerobic aquatic organisms

Prac – Dissolved oxygen in water1). A 200mL sample of water was taken in a conical flask, poured slowly to avoid aeration and temperature measured with a thermometer2). Using a graduated pipette, 2mL Mn(OH)2 and 2 mL alkaline NaI was added3). The flask was stoppered and the solution inverted to mix4). 4 ml of 1M sulfuric acid solution was added, then flask stoppered and inverted again to mix5). 10 drops of starch indicator solution was added; this detects the presence of I2

6). A clean burette was filled with 0.02 M Na2S2O3, and the initial reading recorded7). The water sample solution was titrated with the sodium thiosulfate, swirling constantly until the blue colour completely disappeared8). Steps 1-7 were repeated with tap water and another sample from same spot in creek

Results/Analysis4 Mn2+

(aq) + O2(aq) + 8OH-(aq) 2Mn2O3(s) + 4H2O (brown ppt)

Upon addition of acidMn2O3(s) + 6H+

(aq) 2Mn3+(aq) + 3H2O(l)

2Mn3+(aq) + 2I-

(aq) I2(aq) + 2Mn2+(aq)


Redox titration:I2(aq) + 2S2O3

2- -- (starch indicator) 2I-(aq) + S4O6

2-(aq)

Overall reaction:O2(aq) + 4S2O3

2-(aq)+ 4H+

(aq) 2S4O62-

(aq) + 2H2O(l)

From data: 1mL 0.02M S2O3 = 0.527 mg/L DO@ 15oC saturation = 10.1 mg/L

Creek TapVol. of S2O3

2- titrated (ml) 17.4 17.6[DO]mg/L, % saturation 0.1698 (90.79) 9.2752 (91.83)

Very high creek saturation, probably due to recent precipitation which disturbs water and causes aeration.

Gather, process and present information on the range and chemistry of the tests used to:o Identify heavy metalso Monitor possible eutrophication of waterways

Heavy metals are transition metals plus lead which can be toxic to humans if ingested in higher concentrations

Precipitation test (qualitative)o Water sample acidified and Na2S addedo Precipitate formation indicates presence of one or more of Pb2+, Ag+, Hg2+, Cu2+, Cd2+¸

As3+

ORo If no precipitate forms, water made alkaline then Na2S addedo Precipitate formation indicates Cr3+, Zn2+, Fe2+, Ni2+ …o If precipitate forms when acidified, precipitate is filtered off and remaining solution

made alkaline to check for above ions

Chemistry:o In solution, sulfide and hydronium react according to:

When acidified, reaction proceeds enough to right for only minute concentrations of sulfide left in solution

o Sparingly soluble ions such as lead can precipitate, while others such as zinc cannoto Non-precipitation eliminates presence of these sparingly soluble ionso In alkaline solution, enough sulfide remains to cause noticeable precipitation of ions such

as zinc and iron

o AAS

Eutrophication is the enrichment of water bodies by nutrients such as phosphate and nitrate in excessive amounts (from agriculture, environment or discharged effluents which are decomposed by aerobic bacteria), leading to algal blooms. This increases the BOD and decreases DO, reducing survivability of aquatic life. Algal blooms block out sunlight for plant photosynthesis, further reducing DO and aquatic life. Also develops unpleasant smells and unsightly appearance.

Colorimetry (quantitative) (phosphate)


o Method of measuring phosphate (a nutrient) in water, which corresponds with increases algal growth

o Relies on absorption of light by a coloured solutiono Instrument used are colorimeters which use filters to select wavelengths of light, or

spectrophotometers which use monochromators to more precisely select wavelengtho Measured quantity of ammonium molybdate with catalyst added to sample and carefully

mixed, forming a coloured yellow complex with phosphate (phosphomolybdate)o Measured quantity of ascorbic acid added, which reduces yellow Mo(VI) to intense blue

Mo(V)o Colour absorbance is compared to that of standard solution to determine concentrationo Absorbance is proportional to concentration

Oxygen probe and data loggero Aerobic oxygen decomposing organic waste into products including phosphate and

nitrate use oxygeno Lower DO levels indicate higher microbial activity and thus higher nutrient levels in

watero Oxygen probe lowered into sample, produces a current due to redox reaction and

datalogger converts this into a reading

Advantages of data logger:o Can take readings over long period of time at regular intervals (reliable since larger

volume of data)

Rest is possibly useless:o An oxygen probe consists of a cathode (platinum), and anode (usually silver) surrounded

by KCl solutiono A thin membrane permeable to O2 allows dissolved oxygen in near the cathodeo Since there is no easily reduced cation, the following reaction occurs:

(cathode)

o When electrodes are kept at constant voltage, constant separation and constant surface area, the rate pf electrolysis is proportional to the concentration of DO

o This produces a current which is recorded by a datalogger and converted into a reading

Describe and assess the effectiveness of methods used to purify and sanitise mass water supplies

At the catchment:o Preservation of natural environment – activities such as land clearing and

development cause increased TDS and turbidity; these activities prohibited around catchments. Agriculture and industry can contribute heavy metals or nutrients from fertilisers to water, catchments are distanced from these areas

This minimises the treatment that needs to be done at treatment plants

At treatment plant:o Screening – large debris such as rubbish which can interfere with treatment is removed

Clarification - These processes remove turbidity and colour to give water optical clarity


o Flocculation/Sedimentation – aluminium sulphate or ferric oxide added to cause precipitation of fine suspended particles, otherwise kept apart by surface charge. The particles build as smaller particles adsorb onto them, coagulating to form large lumps. In this process, dissolved particles can also become physically trapped. Resulting mixture is settled in sedimentation tank and sludge removed

Is able to removed dissolved and suspended particles

o Filtration – water is pushed through sand/anthracite filters which trap small particles Cannot remove extremely small particles Fast enough to produce adequate volume of water for big cities

i. Reducing dissolved organic carbon – improves taste and odour, allows water to be safely chlorinated and makes it easier to further treat for domestic use

ii. Removes iron and manganese ions – improves taste, and eliminates stains on laundries, fixtures and coatings on pipes

iii. Reducing phosphate concentrations – reduces algal bloomsiv. Reduces suspended particles – improves aesthetic appeal and makes it easier

to treatv. Adsorbing organic matter – anthracite is able to remove odour and taste from

water

Sanitation – this removes anything harmful to human beingso Disinfection – addition of a disinfecting agent to kill microorganisms, enough added so

residual amount remains to prevent further infection during distribution Addition of chlorine gas or liquid sodium hypochlorite, strong oxidising

agents – cheap and effective, but can fail e.g. Sydney Giardia/Cryptosprodium scare in Sydney. Not effective against viruses

Addition of ozone – stronger oxidant than chlorine and more effective against bacteria and viruses but more expensive

(include membrane filters in a question)Assessment:1). Sydney water uses some of these methods – water out of tap is potable, clear and free of smell, parasites such as Giardia and Cryptosporidium are kept at safe level2). Water analysis has shown membrane filters produce water with less TDS and other contaminants than traditional methods – method trialled in London with success

Describe the design and composition of microscopic membrane filters and explain how they purify contaminated water

o Thin film of synthetic polymer (e.g. polypropylene) or ceramic with microscopic holes of approximately uniform size. Holes made by beam of ions through sheet of polymer, then washing in alkaline solution. Changing ion changes pore size. Various forms include a folded sheet which can be placed into pipes and used, or a bundle of hollow tubes – all of which are very thin, so that water can flow through more quickly

Diagram:o Reverse osmosis filters consist of cellulose

acetate or polyamid attached to another polymer

o Water forced through the tiny holes by gravity, vacuum or pressure pumps, which traps larger particles and allows smaller particles through


o These trapped larger particles are water contaminants such as bacteria and suspended matter, whilst water molecules are small enough to pass freely through

o This results in purer watero The particles filtered depends on the size of the holes:

Microfiltration – inorganic and biological particles included supsnded solids, protozoans and bacteria

Ultrafiltration – can remove fine suspended particles, viruses and water borne parasites such as Giardia

Reverse osmosis – can filter out viruses, bacteria, antibiotics and other chemicals

Present information on the features of the local town water supply in terms of:o Catchment areao Possible sources of contamination in catchmento Chemical tests available to determine levels and types of contaminantso Physical and chemical processes used to purify watero Chemical additives in the water and the reasons for the presence of these additives

Warragamba dam – managed by SCACatchment areaLocated 65 km SW from Sydney, narrow gorge on Warragamba river. Covers 9050 square km and collects water from Coxs and Wollondily water systems, part of Sydney basic catchment area.

Possible sources of contaminationo Agriculture (fertilisers), industry (ions including heavy metals), commercial and

residential allotments (rubbish)o Soil erosion which increases dissolved ions and turbidity o Particulates from fire outbreaks

Chemical testso AAS for heavy metalso Winkler method for DO/BODo pH testing using data loggero Chloride ion titration

Physical and chemical processes to purify waterMain treatment occurs at Sydney Water treatment plants:

o Screening – sieve-like devices removes solid objects such as fisho Coagulation – iron (III) chloride added to cause flocculationo Filtration – water pushed through sand/anthracite filters to remove particulate matter,

cleaned by backwashing with water and airo Microbial Treatment – Chlorine in the form of chlorine gas, liquid sodium

hypochlorite, and calcium hypochlorite tablets to kill microorganisms. Enough added so residual remains throughout distributions. Sometimes, ammonia is added after chlorine in fixed ratio to form monochloramine, which is less reactive and lasts longer

o Other chemical treatment Sodium silicofluoride and hydrofluosilicic acid are added to the water as

mandatory medication. 1ppm which reduces tooth decay and avoids fluorosis.

Lime and carbon dioxide, which react to form calcium bicarbonate, are added at some treatment centres with soft water to increase resistance to pH change, increase hardness and reduce corrosivity.


KMnO4 added to oxidise manganese, converting to insoluble form and filtered out

Gather and process information from secondary sources to identify and analyse the chemical composition of an identified range of pigments

Analyse the relationship between the chemical composition of the metallic component(s) of each pigment in the periodic table

Pigment Colour Composition MetalCadmium Yellow Yellow CdS Cadmium

Haematite Red Fe2O3 IronAzurite Blue Cu3(CO3)2(OH)2 Copper

Cobalt green Green CoO CobaltZinc white White ZnO ZincCinnabar Red HgS Mercury

Pigments all contain transition metals or cadmium and mercury - elements with valence electrons in d-orbitals. These elements have non-degenerate d-orbitals, which have energy differences corresponding to the visible spectrum of light. They absorb certain frequencies of incoming white light, causing the reflected or transmitted light to be coloured as certain frequencies are missing.

Explain that colour can be obtained through pigments spread on a surface layer (e.g. paints) or mixed with the bulk of material (e.g. glass colours)

Paint – reflection of lightGlass – selective transmission of light

Describe paints as consisting of the pigment and a liquid to carry it Explain why pigments used needed to be insoluble in most substances

Paints consist of a pigment dispersed in a liquid called a vehiclePigment:

o Solid material composed of very fine particles insoluble in vehicleo Provides:

Colour Hiding power (opacity) Strength and adhesion Gloss

o Must remain suspended in paint and not settle to form hard sedimento May be organic, inorganic or metallic

o Lakes are organic pigments sourced from plants and animals e.g. cochineal (crimson) from insects, tyrolean purple from marine snail

o They are soluble, don’t get depth and intensity of inorganic pigments, and are broken down by UV radiation and fade

Vehicle:o Can consist of:

A binder which controls flow properties of coating and hardens on exposure to air and gives bulk, gloss and toughness to paint film (natural oils, synthetic organic compound, gum). Drying oils are polyunsaturated, polymerised due to opening of C=C causing chain cross-linking. They physically lock in pigments and are effective binders e.g. linseed/walnut oil

Thinner which is a solvent to facilitate application


Drier – a metal soap which speeds up drying of alkyl-based painto Affects drying time and thicknesso Needs to:

Wet particles of pigment Sufficient viscosity to hold particles in suspension when drying Capacity to form tough adherent film on surface

Pigments need to be insoluble since they are not easily removed or affected by rain or perspiration (makes them moisture resistant). Pigment needs to be opaque and reflective so colour is vibrant (in many cases)

Outline the processes used and the chemistry involved to prepare and attach pigments to surfaces in a named example of medieval or earlier work

Madonna and Child with saints (By Sano di Pietro, medieval)

Support: Oak and pineGround (preparatory layer):Linen glued to poplar with animal glue to protect from environmentGypsum (CaSO4) or Gesso coats the linen to provide smooth white base

Coarse Gesso – coarse and thick, first layerSmooth Gesso – smooth hard second layer made by soaking/slaking plaster in water

Underdrawing was made to outline the paintingPaint:Pigments included: Red Lead (Pb3O4), Azurite (hydrated CuCO3, found in oxidised zones of copper ore deposits), Yellow orpiment (As2S3), soot (carbon)Binder was. Egg tempera method used, finely ground pigment mixed with separated egg yolk, linseed oil and water. Layers applied thinly to prevent shrinking and cracking. Egg tempera holds painting together and gives bulkResin (amber and natural tree resins) added as varnish

Other:Attachment of gold leaf – graffito, gold leaf painted over and then design scratched to reveal it – gilding.Mordant gilding – thin lines of adhesive painted on surface, gold leaf cut and adheres to adhesive, any excess is brushed away

Identify the sources of the pigments used in early history as readily available minerals Identify minerals that have been used as pigments and describe their chemical

composition with particular reference to pigments available and used in art by Aborigines

Early pigments derived directly from naturally occurring coloured earth and soft rocks e.g.

Mineral Colour Chemical formulaCerussite White, grey PbCO3

Stibnite Lead-grey Sb2S3

Cinnabar Red, brownish red HgSOrpiment Lemon yellow As2S3

Ochres (natural earth of silica and clay), and kaolin were used extensively by the aboriginals:

Mineral Colour Chemical formulaRed Ochre (haematite) Red Fe2O3 anhydrous


Yellow ochre (limonite) Yellow Fe2O3 hydrateBrown ochre Brown Fe2O3, partially hydrated

Kaolin white (Al2O3.SiO2.H2O)Umbers (mixtures of manganese oxide and iron oxides) were natural brown clay pigments – also used.

Outline the early use of pigments for,o Cave drawingso Self-decoration including cosmeticso Preparation of dead for burial

Cave Drawings:Naturally obtainable pigments were used in the art of early cave drawings. Lascaux cave paintings in France dating to 15,000 BC, and Aboriginal rock art such as Mimi and Bradshaw showed that humans used:

o Yellow and red ochres (naturally occurring tinted clays)o Black manganese dioxide (MnO2) (ore pyrolusite), or carbon from fires, or charred boneso Kaolin (clay mineral) or chalk (CaCO3) (sedimentary rock) for white

They were ground into powder then mixed with blood, wax, saliva or cave water to create paint. They were applied using the fingers or brushes made of kangaroo fur/a feather.

Self-decoration including cosmetics:Aboriginals, such as Koori, used the above pigments as body paint for decoration, dance and rituals. In PNG, Mount Hagen men painted their faces with black charcoal and white river clays.

The Egyptian, Roman and Greek cultures used pigments as cosmetics, mostly mixed with water or saliva. Egyptians used kohl [made up variously of stibnite (Sb2S3), black manganese oxide (MnO2), lead and CuO] as mascara after being wetted with saliva. Orpiment was used as yellow eye shadow. All three cultures used white lead (PbCO3) for face paint. Greeks used vermillion (chemically same as cinnabar): produced heating mercury and sulfur together in flask – taken out and ground. Henna was also used in Egypt to dye fingernails, palms and soles, and hair.

Preparation of dead for burial:Cro magnon man used red ochre in burial sites, place on chest and head, probably associating it with life-giving blood.Egyptians painted tombs with green, obtained from ground malachite (isolated from mineral by grinding, washing and sifting). It symbolised rebirth and was significant to the afterlife. Sarcophaguses were painted with Lapiz Lazuli, which were believed to give the dead person god-like powers, helping them in the afterlife. Yellow and red ochre were used for paintings illustrating the deceased’s family and slaves.

Identify the chemical composition of identified cosmetics used in an ancient culture and use available evidence to assess the potential health risk associated with their use

In Egypt:Pigment Composition Use HazardCinnabar

VermillionHgS Rouge, lipstick Mercury is not harmful in pure form, but

harmful if allowed to combine with oxygen and hydrogen in water and air. Toxic by ingestion and inhalation in large doses.

Results in numbness, staggered walk, tunnel vision and brain damage

Galena Contains (PbS) Eyeshadow Harmful in small amounts. Most is removed


in urine, but there is risk of buildup. Damages nervous system, causes mental

retardation and deathYellow

OrpimentAs2S3 Eyeshadow Highly toxic, can cause: vomiting; diarrhea;

nausea; numbnessMalachite CuCO3.Cu(OH)2 Eye paint Can cause anemia, liver and kidney damage,

and stomach/intestinal irritation in high doses

Describe an historical example to illustrate the relationship between the discovery of new mineral deposits and the increasing range of pigments

Titanium oxide (TiO2):Historical example: titanium dioxide. Before it was available, the principal white pigments used were calcium carbonate from shells, bones, limestone and marble. These then moved to lead white, used by Ancient Greeks and Egyptians, and European Ease paintings in C19th. ▪ 1791: discovered in ilmenite in England▪ 1795: discovered in rutile – named “titanium” in Germany▪ These Mineral sands were first mined in Australia, by 1960 mined in NSW, Qld, Vic▪ Ilmenite is changed to synthetic rutile (more TiO2) by Becher process or newer chlorination process to produce white pigment/very fine crystalline rutile.▪ Mixed with other white pigments to produce different grades of whites. ▪ Now used widely in paint pigments, sunscreens, and cosmetics

These different pigments provided new shades of colours or new colours altogether.

Describe the development of the Bohr model of the atom from the hydrogen spectra and related energy levels to electron shells

Explain why excited atoms only emit certain frequencies of radiation Explain what is meant by ‘n’, the principle quantum number Identify that, as electrons return to lower energy levels, they emit quanta of energy which

humans may detect as a specific colour

In 1901, Planck formulated the relationship between frequency and energy. He proposed EM radiation was transmitted in discrete units or quanta, called photons. He showed that energy was proportional to frequency.If a hydrogen lamp is subjected to high voltage, the excited electrons will emit a violet light which can be separated into component wavelengths using a prism. Lymann (ultraviolet), Balmer (visible) and Paschen (IR). Rutherford’s model did not explain this, it implied an electron would move smoothly towards nucleus, releasing continuous spectrum

Bohr applied Planck’s concept of energy quantisation to explain the hydrogen spectrum. He proposed that electrons move around the nucleus in a circular orbit

attracted by electrostatic forces without radiating energy, and that an atom could only have a restricted set of discrete energy values (only orbits of certain radii/energies)

Bohr calculated a set of allowed energies using the hydrogen spectrum, and the principle quantum number ‘n (a cardinal)’ denotes the energy level of a particular orbit. Energy is emitted or absorbed Chem notes David Lee BHHS 2007 68

by an atom when an electron moves from one stationary state to another. The difference in energy between initial and final states is equal to the difference in energy between the initial and final energy levels. The lines in the hydrogen spectrum represent a drop from higher energy to lower energy (return to ground state). Energy of photon is exactly equal to energy difference between levels. Some of this radiation corresponds with radiation in visible spectrum.

Solve problems and use available evidence to discuss the merits and limitations of the Bohr model of the atom

Merits:o Explained the observation that excited atoms generated discrete spectrao Predicted with reasonable accuracy the emission spectrum of hydrogeno Incorporated the idea of ‘quanta’ into a model of the atom

Limitations:o When applied to other atoms, predictions failed to agree with experimental resultso Could not explain closely spaced emission lines, or the further energy splitting by

magnetic fieldso Did not explain why certain radii were permitted, or why moving electrons did not lose

energyo Did not explain the different intensities of these lines

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

Simplified sodium emission spectrum shows primarily two yellow spectral lines at 589 and 589.6 nm wavelengths, a ‘doublet’. This causes problems for Bohr model since:

o Closely spaced doublet could not be predicted by Bohr model, the energy transitions between principal energy levels could not account for this energy transition

o One yellow line is more intense than others, and both these lines are far more intense than other emission lines, could not be explained by Bohr

o Magnetic field caused further splitting, could not be explained by Bohr

o The quantum model of the atom, can account for the spectrum through orbitals and orbital splitting

Identify Na+, K+, Ca2+, Ba2+, Sr2+ and Cu2+ by their flame colour Perform first-hand investigations to observe the flame colours above


Na+ - yellow K+ - violetCa2+ - orange-redBa2+ - apple greenSr2+ - redCu2+ - blue-green

Flame test procedure:1). Piece of nichrome wire is cleaned by repeatedly dipping into HCl and heating it to red heat in flame, this is to eliminate sodium which can give intense yellow colour masking other colours2). Wire was dipped into aqueous solutions of samples and placed back into flame3). Flame colour observed

Risk analysis:

Hazard Risk ControlHCl Corrosive. Can cause


Wear safety glasses, have sodium bicarbonate to neutralise spills

Bunsen flame Wire heated in hottest part of flame, can cause burns

Protective gloves

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

The flame is a source of energy that can be absorbed by electrons in the atoms to be excited and move to higher energy states. When they return to ground state, they release energy in the form of EM radiation, which corresponds with frequencies in visible spectrum, producing colour.

Distinguish between the terms spectral line, emission spectrum, absorption spectrum and reflectance spectrum

o Spectral line – a discrete wavelength of light emitted by a radiant source, light emitted by an atom is quantised, not continuous

o Emission spectrum – a set of discrete spectral lines corresponding to the energies emitted by an excited atom, represented by bright lines against dark background

o Absorption spectrum – the specific wavelengths of light absorbed by an atom, appearing as dark lines across a continuous spectrum. Complementary to emission spectrum

o Reflectance spectrum – the wavelengths of light reflected by an element

These spectra can also be represented as intensity vs wavelength plots (quantitative)Each spectrum is unique to an element

Outline the use of infra-red and ultraviolet light in the analysis and identification of pigments and their chemical composition

IR and UV are utilised in spectroscopy techniquesSpectroscopy: the production, measurement and interpretation of electromagnetic spectra interacting with substances (emission, reflectance, absorbance)Quantitative and qualitative analysis in a wide range of applications

Many techniques for analysis, includes:Infrared (700nm to 1 mm):Chem notes David Lee BHHS 2007 70

Source of IR is commonly a heated ceramic e.g. silicon carbide rod. Qualitative methods:

o IR reflectography (used for paintings, quantitative, USE THIS ONE SPARINGLY) Radiation penetrates through most outer pigments and reflects off white

background Carbon from deposited graphite, charcoal or black ink which the artist has

used to make a preliminary outline absorbs the radiation, forming a black image

IR also absorbed by copper-containing green pigments, can provide information about their use

Emission is detected by thermocouple Method is non-destructive (near infrared doesn’t damage it) and can be used

on whole artwork

Quantitative methods:o IR absorption spectroscopy

A double-beam spectrometer is used. One beam passes through the sample and another through the reference

A mono-chromator is used to select particular frequencies of IR to pass through the sample and reference

The molecules in a compound ‘vibrate’, and absorb radiation of same frequency of their natural vibrational frequency, causing less radiation to be transmitted

The pattern of absorption is unique for each compound due to mass of atom, and length/strength of bonds

A detector compares the energy that is transmitted by the sample with that of the reference – an absorption spectrum is plotted from the difference

This absorption spectrum can be compared with standard absorption spectra for certain functional groups and compounds, allowing identification of pigments used

o ATR (Attenuated total reflectance) (Reflectance type) Used to examine samples in liquid or solid state without need for further

preparation For a solid sample, a crystal (e.g. KBr) is clamped tightly to the solid, and an

IR beam is directed through the crystal to the sample Some is absorbed and some reflected The beam is reflected internally within

the crystal and is collected by a detector upon exit

The resulting reflectance spectra is compared with standard sample graphs for compound identification, allowing determination of pigments used

Technique similar for liquids, but liquid is poured onto crystal instead of clamped

Used for analysing coatings, pastes and paints

Ultraviolet (300 nm to 10 nm):UV spectrophotometer, with radiation source a tungsten lamp or deuterium discharge tube (and halogen lamp for the visible region in UV/visible). o Ultraviolet visible spectroscopy (absorption + reflectance)


UV light is shined onto sample, and the reflected beam is collected by a detector (photomultiplier tube)

Absorbance is directly proportional to concentration, thus this can be used quantitatively

Compared with reflectance spectrum of non-absorbing substance such as silica (SiO2) (this is the reference)

OR for absorption spectra of pigments in solution: Solution is placed into a silica cell The absorption spectrum of a reference cell containing only solvent is

determined The UV light is shined onto sample, and a detector compares the radiation

passing through the reference and sample

Absorbance spectrum of sample is determined, and this is used to determine the pigment, its concentration and chemical composition by comparing it with absorbance spectra for different compounds

o UV fluorescence (qualitative, USE SPARINGLY) Different materials can fluoresce when exposed to UV light, the colour being

dependant on chemical composition and age ZnO (zinc white) fluoresces a pale yellow, while green malachite is a dirty

mauve

Can be used for paintings, pigments dissolved in solution etc.

Explain the relationship between the absorption and reflectance spectra and the effect of infra-red and ultraviolet light on pigments including zinc oxide and those containing copper

o The reflectance and absorption spectra are complements of each other when sample is opaque

o Represented as plots of wavelength against intensity of reflection/absorbanceo A material can either reflect or absorb a particular wavelength o Thus, the non-reflected wavelengths must be absorbed and vice versao Putting the two spectra together gives you the original radiation spectrum

IR:o Far IR radiation can change the colour of zinc oxide from white to yellow in the presence

of oxygen due to increased temperature, this is reversible by decreasing tempo Red copper (I) oxide and malachite permanently change to black copper (II) oxide since

the heat causes breakdown

UV:o ZnO (zinc white) fluoresces a pale yellow, while green malachite is a dirty mauve

(CuCO3.Cu(OH)2), since they absorb the radiation but the electrons don’t return immediately to ground state, going down transitional stages and thus emitted different wavelengths to the originally absorbed one

Gather, process and present information about a current analytical technology to:o Describe the methodology involvedo Assess the importance of the technology in assisting identification of elements in

samples and in compounds, ando Provide examples of the technology’s use


Laser microspectral analysis

o A laser beam is directed at a sample of the pigment, vaporising a tiny amounto This vapour passes between two electrodes which spark and excite the vapour particleso An emission spectrum is obtained by a spectrophotometer when electrons return to

ground state, emitting radiationo This can be compared with the unique emission spectrum of different elements to obtain

the chemical composition and thus the pigments presento The intensity of the spectra is proportional to concentration

Importance:o Laser beam is very intense, pure and focused, allowing the technique to detect trace

amounts of elements in samples o It is useful for determining authenticity, and in restorationo Requires minimal preparation, so can be used for a variety of solid samples

Examples of use:o Analysis of paintings for validity (time period etc.)o Used to analyse chemical compositions of paintings so correct restorative chemicals are

used

Define the Pauli Exclusion Principle to identify the position of electrons around an atom

No two electrons in an atom may have identical sets of four quantum numbers

Identify that each orbital can only include two electrons

Only two electrons which have opposite spins (anticlockwise, clockwise). These produce magnetic fields, and result in slightly different energy levels for each electron. Note: This arrangement produces a lower potential energy than if they were paired, since the electron-electron repulsion is minimised

Define the term sub-shell

A sub-shell is an energy sublevel within a principle energy level (or shell). S, P, D, F. Number of subshells = principle energy level. S – spherically symmetrical around nucleus, P – dumbbell shaped with 3 orientated in perpendicular planes through nucleus

Outline the order of filling of sub-shells


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

When an electron absorbs a quanta of energy corresponding to the difference between its current energy level and another energy level, it moves to the higher energy level to expend this energy.

Explain the relationship between the elements with outermost electrons assigned to s, p, d, f blocks and the organisation of the periodic table

o Elements with similar outer shell electron configurations occur in the same groupo Periods correspond to principle energy levelso The periodic table is divided into blocks depending on the element’s outer shell

Explain the relationship between the number of electrons in the outer shell of an element and its electronegativity

Electronegativity – a measure of the ability for an atom to attract electrons to itself in a chemical bondElectronegativity increases across a period – more protons in nucleus meaning increased nuclear charge. Electrons added do not fully shield each other from effect of increased protons, meaning electrostatic force on electrons around the nucleus is strongerElectronegativity decreases down a group – the size of an atom increases, and the attractive force of the nucleus on valence electrons is diminished

Analyse information about the relationship between ionisation energies and the orbitals of electrons

Describe how trends in successive ionisation energies can be used to predict the number of electrons in the outermost shell and the sub-shells occupied by those electrons

The ionisation energy is the energy required to remove an electron completely from the nucleus’ electric field.

In the s-subshell, ionisation energy increases as the number of electrons increases. The addition of an electron a similar distance from the nucleus corresponds with the addition of protons to the nucleus, increasing nuclear charge. The electrons do not shield each other enough to cancel this increased nuclear charge, causing a net increase of electrostatic force towards the nucleus, increasing the ionisation energy.

In the p,d,f-subshells, the ionisation energy increase as the electrons are added into different orbitals (See above), but decreases when an electron is first paired in an orbital. This is because the electrons repulse each other enough to overcome the effect of increased nuclear charge, raising each other to higher potential energies and decreasing ionisation energy.

It decreases down a group as valence electrons are further from the nucleus, and thus at higher potential energies.


Use Hund’s rule to predict the electron configuration of an element according to its position on the periodic table

Hund’s rule – every orbital in a subshell must be singly occupied by one electron with identical spin before any one of these orbitals in the same subshell are doubly occupied

Identify the block occupied by the transition metals in the periodic table Define the term transition element

Transition element - those that form at least one ion with a partially filled sub-shell of d electronsZn is not a transition element since its ion Zn2+ has a completely filled d-subshell. Scandium ion (Sc3+) has no electrons in d-subshell.

They have similar properties since their outer subshell is only s1 or s2

The partially filled d-subshell explains:o Colour of metal ion complexes – d-d transitionso Magnetic properties of metal ions – unpaired electrons in d-orbitals causes

paramagnetism, fully paired causes diamagnetismo More than one stable oxidation state – due to the similarity of the s and d-subshell energy

levels, the electrons in both levels can be lost without high energy cost, meaning the transition metal forms multiple oxidation states

*In other elements, the energy levels are too separated to have commonly observable multiple oxidation states.

Explain why transition metals may have more than one oxidation state Account for colour changes in transition metal ions in terms of changing oxidation states Explain, using the complex ions of a transition metal as an example, why species

containing transition metals in a high oxidation state will be strong oxidising agents

Oxidation number is the charge an atom would have in a chemical bond if the bonded electrons belonged to the more electronegative element.

The number of oxidation states for the first period of transition elements progressively increases to a peak at VIIB then decreases, since:

o As more electrons are added to a d-orbital, the d and s subshells progressively separate in terms of energy due to electron repulsion

o However, more electrons also means more possible oxidation stateso The “sweet spot” is Mn, with the number decreasing to the left due to lack of electrons,

and decreasing to the right due to separation of energy levels

CHECK – why do transition metals need to be in oxidised states for this to occur?


CHECK – what about return to ground state, wouldn’t it release the light again?Transition metals ions in compounds have d orbitals with slightly different energy levels and are incompletely filled. These energy differences correspond with the energy of visible light, and electrons in the d subshell can absorb photons to become excited, meaning the complementary light spectrum is able to pass through and be observed. E.g. transition metals which absorb the red end of the spectrum will appear blue. In different oxidation states, the transition metal has a different arrangement of filled and unfilled 3d orbitals, causing differences between d energy levels and thus causing different wavelengths of visible light being absorbed.

Transition metals in a high oxidation state have a high deficit of electrons, which reduces the orbital radii and decreases electron shielding. This results in a high oxidation potential and makes them strong oxidising agents. E.g. Cr2O7

2- and MnO4-

Write electron configurations of the first transition series in terms of subshells

As per Hund’s Rule. Note: The d-orbitals are in their most ‘stable’ configuration with a complete set of unpaired or paired electrons. When the set is almost complete, an electron from the 4s orbital will transit to make it complete (Chromium and Copper).

Perform a first-hand investigation to observe the colour changes of a named transition element as it changes in oxidation state

Prac – Oxidation states of VanadiumEquipment:100 mL of 1M NaOH3 grams of ammonium vanadate75 mLs of 2M H2SO4

Granulated zinc4 large test tubes and text tube rack250 mL conical flask and rubber stopper

Method:1. 3 grams of ammonium vanadate was weighed on an electronic beam balance 2. 100 mL of 1M NaOH and the 3 grams ammonium vanadate were added to the conical flask3. The mixture was swirled to dissolve4. 75 mLs of 2M H2SO4 was added to the solution to acidify it5. 20 mL of the resulting solution was poured into a test tube using a 50 mL measuring cylinder6. 8 granules of zinc were dropped into the conical flask which was then stoppered7. The solution was swirled gently until it became blue, then step 5 performed8. The solution was swirled again until it became green, then step 5 performed9. The flask was swirled vigorously until the solution became violet, then step 5 performed

Results and analysis:(NH4)VO3(s) NH4

+(aq) + VO3

1-(aq)

(+5 YELLOW to +4 BLUE)Oxidation: Zn(s) Zn2+

(aq) + 2e-

Reduction: VO31-

(aq) + 4H+(aq) + e- VO2+

(aq) + 2H2O(l)

Redox: Zn(s) +2VO31-

(aq) + 8H+(aq) 2VO2+

(aq) + 4H2O(l) + Zn2+(aq)

(+4 BLUE to +3 GREEN)Oxidation: Zn(s) Zn2+

(aq) + 2e-

Reduction: VO2+(aq) + 2H+

(aq) + e- V3+(aq) + H2O(l)


Redox: Zn(s) +2VO2+(aq) + 4H+

(aq) 2V3+(aq) + 2H2O(l) + Zn2+

(aq)

(+3 GREEN to +2 VIOLET)Oxidation: Zn(s) Zn2+

(aq) + 2e-

Reduction: V3+(aq) + e- V2+

(aq) (CHECK IT)Redox: Zn(s) +2V3+

(aq) 2V2+(aq) + Zn2+

(aq)

Risk analysis:

Hazard Risk ControlSulfuric acid Corrosive. Can cause



Sodium Hydroxide Corrosive. Can cause permanent eye damageKills tissue


Ammonium vanadate Eye contact causes redness and swellingSkin contact causes itching and pain

Safety goggles, gloves, do not make airborne

Rules for balancing redox half equationsIn acidic solution:1. Write down the reactant and product [Cr2O7

2- Cr3+]

2. Balance the number of atoms [Cr2O72-

2Cr3+]3. Balance the number of atoms of oxygen by adding water to the side with least oxygen

[Cr2O72-

2Cr3+ + 7H2O]]

4. Balance the hydrogen by adding H+ ions [Cr2O72-

+ 14H+ 2Cr3+ + 7H2O]

5. Balance the charge by adding electrons then add states [Cr2O72-

(aq) + 14H+(aq) + 6e- 2Cr3+

(aq) + 7H2O(l)]

CHECK – what about 2HI I2 ?

In alkaline solution:1. Balance as though in acid solution2. H+ ions are removed to form H2O by adding same number of OH- ions to both sides3. Simplify

e.g. [MnO4

-(aq) + 4H+

(aq) + 3e- MnO2(s) + 2H2O(l)][MnO4

-(aq) + 4H+

(aq) + 4OH-(aq) + 3e- MnO2(s) + 2H2O(l) + 4OH-]

[MnO4-(aq) + 2H2O(l) + 3e- MnO2(s) + 4OH-]

choose equipment, perform a first-hand investigation to demonstrate and gather first hand information about the oxidising strength of KMnO4

Prac – Oxidising strength1). 20 mL of 0.01 M KMnO4 and 20 mL 1M H2SO4 were added to the same 50 mL measuring cylinder2). 8 medium test tubes were set up on a test tube rack 3). About 5 mL of 0.5 M solutions of KI, KBr and KCl droppered into the first 3 using an eye dropper, and zinc powder, a short magnesium strip, copper pieces, tin pieces and iron nails placed into the five remaining test tubes corresondingly3). The eye dropper was rinsed, then used to drop a few drops of the acidified KMnO4 solution into each of the test tubes


4). A centimetre of toluene was poured into these three test tubes not containing metals, to help indicate redox reactions as halides discolour it5). Any changes were observed5). The contents of the test tubes containing toluene were disposed of in an organic waste bin

Results:Constant Reactant: KMnO4 solution (MnO4

- is pink)MnO4

- + 8H+ + 5e- Mn2+ + 4H2O (E0 = 1.51 V)E0 (V) (for

halide/metal)Variable Reactant

Colour KMnO4

solutionProduct state Reaction E0

(V)0.62 KI Colourless I2 (toluene -> pink) 0.891.09 KBr Colourless Br2 (toluene ->

yellow/brown)0.42

1.40 KCl No change* 0.11-0.76 Zn Colourless Mn2+ 2.27-2.36 Mg Brown MnO2 3.870.35 Cu Colourless Mn2+ 1.160.14 Sn No change0.77 Fe Colourless Mn2+ 0.74

Some change would be expected to occur (under standard conditions), but the already low reaction E0 + the non-standard state of the reactants meant there was too small a tendency of reaction for observable change

Relative oxidation strengths:

Mg2+ < Zn2+ < Sn2+ < Cu2+ < I2 < Fe2+ < Br2 < Cl2 < MnO4-

Risk analysis:

Hazard Risk ControlPotassium permanganate Stains skin and clothing,

concentrated solutions are corrosive


Toluene Very flammableSkin irritant

Keep away from hot surfaces, flames or sparksPolyvinyl gloves

Sulfuric acid Corrosive. Can cause permanent eye damageKills tissue


The permanganate ion is a strong oxidant due to the presence of manganese in a high (+4) oxidation state. This gives it a high electronegativity, thus making it a strong oxidant and in turn a weak reductant.

Some example reaction formulae,(Iodide and permanganate) Reduction: MnO4

- + 8H+ + 5e- Mn2+ + 4H2O (E0 = 1.51 V)Oxidation:5I- (5/2) I2(aq) + 5e- (E0 = -0.62 V)Overall:MnO4

- + 5I- + 8H+ Mn2+ + 4H2O + (5/2) I2(aq)

explain what is meant by a hydrated ion in solution


When an ionic solid dissolve in water, the ions dissociate and are surrounded by water molecules in a process called hydration. The charge of the ion attracts the polar water molecule. These ions are hydrated. For many metal cations, a covalent bond is formed between the oxygen atom and the cation due to vacant orbitals in the metal.

The general formula is [M(OH2)n]m+ where ‘m’ is the charge of the ion and ‘n’ is the no. of water molecules surrounding the ion. Note that it is OH2 since the “O” is donating the electrons.Note that many other ligands are possible.

describe hydrated ions as examples of a coordination complex or a complex ion and identify examples

describe molecules or ions attached to a metal ion in a complex ion as ligands explain that ligands have at least one atom with a lone pair of electrons

A complex ion (a lewis acid) is where a central metal ion is surrounded by ligands. Ligands (a lewis base) are atoms, ions or molecules which donate one or more electrons in a coordinate covalent bond with the central metal ion. Ligands therefore need a lone pair of electrons which it can donate to an empty orbital of the central ion, in order to bond.

In a hydrated ion, the ligands are water molecules.

identify examples of chelated ligands

Ligands can bond using different numbers of electron pairs. Monodentate ligands bond using the electron pair of a single donor atom. Others have multiple atoms with unpaired electrons and can bind simultaneously, and are polydentate or chelated ligands. Note: Memorise ‘en’ and ‘EDTA’

Example of chelated ligand bonding:


discuss the importance of models in developing understanding of the nature of ligands and chelated ligands using specific examples

Models are significant in understanding the nature of ligands. They allow us to specify mechanisms for the formation and bonding of ligands, and other phenomena where the mechanism cannot be directly observed but only inferred (e.g. colours). They are not necessarily true, but are adequate explanations for current use.

1). Valence Bond TheoryAssumes:

o Ligands bond to central ion in completely covalent coordinate bonds

Bonding in the formation of complexes depends on:o Orbitals available for coordinate covalent bond formationo Tendency of ions or groups to share a pair of electronso Number of ligands that can be placed around the central iono The geometry assumed by the ligands

1. The electron pairs from the ligands are placed in empty orbitals of the central ion2. These orbitals are hybridized, that is, they are mixed to form new orbitals3. Some ligands cause unpaired electrons in d-orbitals to pair with other unpaired electrons, and

then use the newly empty d-orbitals. Others cannot do this, and use completely vacant orbitals. 4. Those ligands utilising d-orbitals previously occupied have inner spin complexes, the ones not

have outer spin complexes5. Inner spin complex are paramagnetic, and move in the direction of a magnetic field. Outer spins

are diamagnetic, and move opposite to the direction of a magnetic field6. The structural geometry can be determined using VSEPR rules

The advantages are:o Accounts for magnetism of coordination complexeso Accounts for shape of complexes through hybridisation

Disadvantages:o Cannot show details such as the energy changes involved, rate at which reactions occur

and the mobility/flexibility of bonds involvedo Cannot account for colours of transition metal compounds

2). Crystal field theory Assumes:

o Bonds between ligand and central ion are completely ionico Ligand and central ion are infinitesimally small, non-polarisable point charges

The ligand is treated as a point negative charge, and when it approaches the central metal ion, the electron clouds of both get disturbed, resulting in changed energy states for orbitals such as the d-orbitals. Different ligands and different metal complex geometries result in different degrees of energy separation. The resulting separation of the d-orbitals is in the energy range of visible light, and different energies absorb different wavelengths, resulting in different observed colours.

Advantages:o Explains the colours of transition metal complexes

Disadvantages:o Neglects any covalent contribution


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

Some examples:

process information from secondary sources to give an example of the range of colors that can be obtained from one metal such as Cr in different ion complexes

The colour depends on the central metal ion (element and oxidation state) and the surrounding ligands.

Prac – Colours of Chromium complexes1). Place 5 mL of chromium (III) nitrate solution into five large test tubes2). Record the colour of the solution in the first test tube0

3). 4 rice grains of sodium sulfate was added to the next test tube, and gently warmed over a bunsen burner1

4). The colour of the solution was recorded5). 4 rice grains of sodium chloride was added to the next test tube, and gently warmed over a bunsen burner2

6). Colour was recorded7). A dropping pipette was used to drop 4 drops of 1 M sodium hydroxide solution into the next test tube, then colour recorded3

8). More sodium hydroxide was added to the test tube until the precipitate disappeared, then colour recorded4

9). Ammonia solution was added to the last test tube so a precipitate formed, then more added until the precipitate disappeared10). Colour recorded5

0. [Cr(H2O)6]3+ (blue)

1. [Cr(H2O)5SO4]+ (green)

2. [Cr(H2O)4Cl2]+ (green)

3. Cr(H2O)3(OH)3 (grey-blue ppt)4. [Cr(OH)6]3- (green)5. [Cr(NH3)6]3+ (mauve)

Errata:1). Added an actual evaluation in the summary of ethanol usage, changed it2). The diagram for dry cell is incorrect (page 23-24)3). Some content changes made to reflux (page 45)4). Defined allotrope (page 54) 5). Added risk analysis (page 69)6). Big error!! Explanation of why transition metals compounds are coloured was incorrect before. See last sentence of paragraph below table (page 64)7). Should be bromine water not iodine in the lycopene prac (page 7)8). Ester structure clarified. The way I wrote them as word equations before was misleading (page 44)9). Completed the IR Absorption spectroscopy description (page 71)10). Made risk analysis more specific for bromine water prac (page 6)11). Scintillation counter description before was INCORRECT (page 27)


12). Modified the explanation explaining multiple oxidation states of transition metals13). Ester page updated with more information (page 44)14). Made the uses of ethylene page more readable15). Error! Iron sulfide in coal (production of sulfur dioxide) – equation is incorrect, see correction page (34)16). Expanded description of Arrhenius acids, including limitations, and also Bronsted-Lowry theory (page 39)17). Clarified conjugate base/acid explanation (page 39)18). Changes to history of pigment use – I got kohl composition wrong 19). Amphoteric is NOT amphiprotic! Amphoteric means it has both acidic and basic PROPERTIES. Amphiprotic means it can both accept and donate protons which ARE properties of most acids/bases but do not encompass all of them


chemistry hsc full notes best notes

Documents