Double bond

Catalyst/ InitiatorEthylene moleculeActivated species

9.2.1 Fossil fuels provide both energy and raw materials such as ethylene for the production of other substances

Identify the industrial source of ethylene from the cracking of some of the fractions from the refining of petroleumIndustrial source: crude oil/ petroleum; either as by-product of petrol refining or ‘cracking’ of higher boiling point fractions. Cracking is the chemical process of breaking large hydrocarbons into smaller ones. Catalytic cracking is the process in which high molecular weight fractions from crude oil are catalytically broken into lower molecular weight substances (an alkane and alkene), like petrol, to increase output of these high demand products.

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

Ethylene has a double bond since it’s an alkene, which makes it reactive. Ethylene undergoes an addition reaction in which the double bond opens out to form two single bonds thus linking molecules together, i.e. (CH2=CH2)n → (-CH2 – CH2-)n. It forms useful products such as ethanol as a result.

Identify that ethylene serves as a monomer from which polymers are madeEthylene monomers polymerise to form low or high density polyethylene.Low density polyethylene is made from the gas phase process – high pressure, high temperatures and an organic peroxide initiator.High density polyethylene is made using Zeigler-Natta process – just above atmospheric pressure, temperatures around 60°C and zeolite catalyst.

Identify polyethylene as an addition polymer and explain the meaning of this termPolyethylene is called an addition polymer. This means that it forms by molecules adding together without the loss of any atoms. Each double C=C bond opens out to form single bonds with adjacent molecules thus linking molecules together.

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

1. Catalyst (for HDPE) or initiator (for LDPE) attaches to ethylene molecule → creates activated species.

Z + CH2=CH2 → Z – CH2 – CH2

2. Ethylene molecules attach to the species, expanding the chain.

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C6H5

Z – CH2 – CH2 + CH2=CH2 → Z – CH2 – CH2 – CH2 – CH2

3. Polymerisation stops when two activated species collide, forming a stable polymer.Z – (CH2 – CH2)x + Z – (CH2 – CH2)y → Z – (CH2 – CH2)x+y – Z

Note: Since activated chains of various lengths can collide, polymer molecules formed have different chain lengths and different masses, therefore, in a polymer sample there is a distribution of molecular weights, but the average is 46000.

Identify the following as commercially significant monomers: - vinyl chloride- styrene, both by their systematic and common names

Vinyl Chloride StyreneSystematic Name: chloroethene Systematic Name: phenylethene/ ethenylbenzeneCommon Name: vinyl chloride Common Name: Styrene

Describe the uses of the polymers made from the above monomers in terms of their propertiesPolymer Structure Properties Uses

Low density Polyethylene

Extensive chain branching; no stiffening side groups; no cross linking

Soft; flexible; low melting point; transparent; not strong

Cling wrap, carry bags; squeeze bottles

High density Polyethylene

No chain branching; no chain stiffening side groups; no cross linking

Hard; brittle; high melting point; translucent

Kitchen utensils; food containers; milk bottles; rubbish bins

Poly (vinyl chloride)

Considerable chain stiffening Cl side groups; polar C-Cl bonds produce strong intermolecular forces

Hard; inflexible; vulnerable to UV attack (inhibitor added to absorb UV light, preventing degradation)

Electrical wiring insulation; garden hoses; drainage and sewerage pipes

Polystyrene Large phenyl stiffening side groups; minimal chain branching; C-C and C-H bonds

Crystalline PVC - transparent; resistant to UV attack; hard; rigid

Expanded PVC – light weight; spongy; moulded easily; good insulator; soft

Disposable drink glasses; foam packing material

Identify data, plan and perform a first-hand investigation to compare the reactivities of appropriate alkenes with their corresponding alkanes in bromine waterBackground Information: Alkanes and alkenes are both non-polar molecules with weak dispersion (intermolecular) forces. Alkanes undergo substitution reactions (reactions in which an atom in a molecule is replaced by another atom or group of atoms). Alkenes undergo addition reactions.Aim: To compare the reactions of cyclohexene and cyclohexane in bromine water

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Chlorine replaces hydrogen of regular ethene molecule

Benzene ring replaces hydrogen of regular ethene molecule

Safety Precaution: Cyclohexane, cyclohexene and bromine water are toxic by all routes of exposure; protect yourself by using small quantities, wearing safety glasses and avoiding inhalation by using a fume cupboard.Method: In two test tubes, place ten drops of cyclohexene and cyclohexane. To each sample add 10 drops of bromine water. Shake vigorously. Observe for a colour change.Results: cyclohexene decolourises bromine water; cyclohexane shows no reactionConclusion: Cyclohexene undergoes an addition reaction: C6H10 + Br2 → C6H10Br2 (cyclohexane + bromine water → 1, 2 – dibromocyclohexane)Cyclohexane does not undergo a chemical reaction (however, in the presence of UV light, a substitution reaction occurs)

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9.2.2 Some scientists research the extraction of materials from biomass to reduce our dependence on fossil fuels

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

There is only a finite supply of crude oil; all reserves will be used up within the next few decades

There is pressure to reduce energy use and greenhouse gas emissions As oil supplies diminish, oil prices rise, one day to such an extent that oil will be too expensive a

fuel source When oil reserves run empty, if there were no alternatives, there would be no fuel or plasticFor these reasons and more, ethanol and cellulose are being researched as alternative fuel sources.

Explain what is meant by a condensation polymerCondensation polymers are chains of linked monomers that form when a functional group of one monomer reacts with the functional group of another monomer, joining and eliminating a small molecule (often water).

Describe the reactions involved when a condensation polymer is formedUsing glucose as an example,

… HO – C6H1004 – OH HO – C6H1004 – OH HO – C6H1004 – OH…→ O – C6H10O4 – O - C6H10O4 – O - C6H10O4 + xH2O

When two glucose monomer molecules react through two -OH hydroxyl groups, a H-OH (water) molecule is condensed out, leaving an -O- linking the two monomer molecules. The first two glucose molecules to join condense out an H-OH, and every glucose molecule added to the growing chain then condenses out another H-OH.

Describe the structure of cellulose and identify it as an example of a condensation polymer found as a major component of biomassCellulose is a long, linear molecule because of the alternating CH2OH groups on either side of the chain and C–O–C bond angles. Hydrogen bonds make cellulose difficult to break down, rigid and strong. The OH groups can not interact with water, making cellulose insoluble. Cellulose has potential as a biopolymer because it is a major component of biomass (organic material derived from living organisms, e.g. crops, animal waste).

Identify that cellulose contains the basic carbon chain structures needed to build petrochemicals and discuss its potential as a raw materialCellulose contains the basic carbon structures required by the plastics industry to be a source of chemicals or a biopolymer. Plastics made from cellulose biodegrade into fungi and bacteria. Cellulose → thermochemical pretreatment + hydrolysis → breaks down into constituent sugars → fermented → ethanol → polymerised to form useful productsHowever, the breaking down of cellulose requires much energy (as it is hydrogen bonded) and is more expensive than cracking crude oil→ form ethylene → hydrated to form ethanol. Nevertheless, research on cellulose needs to continue to see through its potential as a raw material.

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Use available evidence to gather and present data from secondary sources and analyse progress in the recent development of a named biopolymer. This analysis should name the specific enzyme used or organism used to synthesise the material and an evaluation of the use or potential use of the polymer produced related to its propertiesPolylactic acid (PLA) is a biodegradable thermoplastic derived from renewable plant material such as corn starch. Corn kernels are milled → extract starch → enzymes break down starch in dextrose → lactic acid bacteria converts dextrose to lactic acid → polymerised → PLAUses:

Plant pots, mulch film and disposable nappies because of its biodegradability and compostability.

Its transparency, rigidity and crack-resistance makes PLA suited for use as food containers and drink cups, since the food would be visible and containers would not break.

Advantages of PLA: Biodegradable; compostable Sustainable since it is made from renewable resources – corn Less greenhouse gas emissions, no toxic gases Requires less energy for production than conventional plastics

Disadvantages of PLA: Only biodegradable under ‘controlled composting environmental conditions’ which is not

readily accessible by consumers PLA breaks down into lactic acid, which demands a lot of oxygen. However, research is being

conducted into anaerobic digesters so that PLA can break down without oxygen.

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9.2.3 Other resources such as ethanol are readily available from renewable resources such as plants

Describe the dehydration of ethanol to ethylene and identify the need for a catalyst in this process and the catalyst usedC2H5OH → C2H4(g) + H2O(l) in the presence of heat and concentrated sulfuric acid which acts as a dehydrating agent and catalyst

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 usedC2H4(g) + H2O(l) → C2H5OH in the presence of heat and dilute sulfuric acid which acts as a hydrating agent and catalyst

Describe and account for the many uses of ethanol as a solvent for polar and non-polar substancesThe ethanol molecule consists of two parts – the non-polar alkyl (-CH2CH3) end and the polar hydroxyl (-OH) end.

Outline the use of ethanol as a fuel and explain why it can be called a renewable resourceEthanol has been proposed as an alternative fuel source because:

It undergoes complete combustion and unlike petrol, burns efficiently → does not release soot or carbon monoxide.

It is a renewable resource since it is made through fermentation of plant material, and so would reduce our dependency on non-renewable crude oil.

Ethanol is made from carbon dioxide, water and sunlight and plant material and when it is burnt it returns back to carbon dioxide and water, which can be reconverted into ethanol.

Assess the potential of ethanol as an alternative fuel and discuss the advantages and disadvantages of its useAdvantages of ethanol as a fuel:

At 10-20% concentration in petrol, it is a ‘petrol extender’ and vehicle engines do not need to be modified to utilise it.

It undergoes complete combustion [see previous dot point] It is ‘greenhouse neutral’ since the carbon dioxide released during fermentation and

combustion is used up in photosynthesis. i.e. Fermentation: C6H12O6 → 2C2H5OH + 2CO2 (g)

+ Combustion: 2C2H5OH + 3O2 (g)→ 4CO2 (g) + 6H2O(l)

= Photosynthesis: 6CO2 (g) + 6H2O (l) → C6H12O6 + 6O2

Disadvantages of ethanol as a fuel: Large areas of agricultural land need to be devoted to growing crops for fuel instead of for

food

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Non-polar alkyl end forms dispersion forces with non-polar substances, thus dissolving them

Polar hydroxyl end forms dipole-dipole or hydrogen bonds with polar substances, thus dissolving them

It does not release as much energy as petrol → vehicles receive fewer kilometres (-1367kJ in ethanol vs. -5464kJ in petrol)

Smelly wastes present environmental problems

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

Process information from secondary sources to summarise the use of ethanol as an alternative car fuel, evaluating the success of current usage In Australia, ethanol is currently added to petrol to form E10 Unleaded which is 10% ethanol/90% petrol blend. At this concentration, vehicle engines need not be modified; car manufacturers claim that higher concentrations corrode vehicle engines. Ethanol is more expensive and less efficient than petrol as a car fuel because and vehicles receive fewer kilometres because ethanol contains less energy. Furthermore, there are no reliable studies to show that ethanol produces less greenhouse gas emissions than petrol (although it does combust cleanly and produces less soot and carbon monoxide).

Describe conditions under which fermentation of sugars is promotedFermentation is a process in which glucose is broken down into ethanol and carbon dioxide by the action of enzymes in yeast. For fermentation:

Suitable grain or fruit is mashed up with water Yeast is added Air is excluded The mixture is kept at 37°C (body temperature)

Summarise the chemistry of the fermentation process

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Sugarcane Harvested and transported to mill

Clarified to remove impurities

Shredded and crushed to extract sugar-rich

juice

Mixed with enzymes and water to convert sugarcane into pure

sugars

Heated to turn starch into liquid

Cooled; yeast is added to convert sugars to ethanol and carbon

dioxide

Carbon dioxide is removed then

remaining mixtureis fractionally distilled

Pure ethanol

Present information from secondary sources by writing a balanced equation for the fermentation of glucose to ethanol

1. Starch/ sucrose are mixed with enzymes (a biological catalyst) to convert it into glucose.2. Glucose mixture is clarified to remove impurities and waste, cellulose.3. Yeast is added to convert mixture into carbon dioxide and ethanol (fermentation).

i.e. C6H12O6 → 2C2H5OH + 2CO2 (g)

4. Fermented further to produce ethanol at 15% concentration (any higher would kill yeast, ceasing further fermentation)

5. 15% ethanol mixture is fractionally distilled → 95% ethanol

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

1. Measure and record the mass of the conical flask.2. Add an aqueous mixture of glucose and measure and record the mass of the flask plus the

glucose mixture. Calculate the mass of the glucose mixture.3. Add 1 gram of yeast and secure flask opening with lid containing a pipe which leads to a test

tube containing limewater.4. Incubate overnight. 5. Measure mass of conical flask with fermented ethanol mixture. Calculate mass changes.6. Record observations of limewater.

Results: 5 grams loss, limewater turns milky as carbon dioxide is produced during fermentation.Note: Additionally, you can test for ethanol by adding potassium permanganate (KMnO4) to a sample of the final mixture, it should turn colourless.

Define the molar heat of combustion of a compound and calculate the value for ethanol from first-hand data The molar heat of combustion of a substance is the heat liberated when one mole of the substance undergoes complete combustion with oxygen at standard atmospheric pressure with products being carbon dioxide and water. Value for ethanol: -1367 kJ mol-1

Identify the IUPAC nomenclature for straight-chained alkanols from C1 to C81 – meth2 – eth

3 – prop4 – but

5 – pent6 – hex

7 – hept8 – oct

Formula for alkanols: CnH2n+1OHFormula for alkanes: CnH2n+2

Formula for alkenes: CnH2n

Identify data sources, choose resources and perform a first-hand investigation to determine and compare heats of combustion of at least three liquid alkanols per gram and per moleCalculating Heat Capacity: ∆H= -mC∆Ti.e. Heat Capacity (in joules) = mass in grams of water X 4.18 X temperature changeHeat of Combustion per gram: ∆H (in kilojoules) ÷ mass of alkanol usedHeat of Combustion per mole: Heat of combustion per gram X molar mass of alkanol

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9.2.4 Oxidation-reduction reactions are increasingly important as a source of energy

Explain the displacement of metals from solution in terms of transfer of electronsReactions between a metal and a solution containing the ions of a different metal are displacement or redox reactions. The metal is oxidised and dissolves and the ions of the metal in solution are reduced to form a solid metal deposit.Oxidation: loss of electrons (LEO)Reduction: gain of electrons (GER)

Identify the relationship between displacement of metal ions in solution by other metals to the relative activity of metalsDisplacement reactions are electron transfer reactions. In such reactions, a more active solid metal oxidises and will displace the ions of a less active metal in solution (which is reduced).Note: On the standard electrode potentials table, the metal higher up is oxidised.

Account for changes in the oxidation state of species in terms of their loss or gain of electronsOxidation → increase in oxidation stateReduction → decrease in oxidation stateElements in their naturally occurring state have an OS of 0.For ions, the OS is the charge/ valency of the ion (e.g. SO4

2- has OS -2)Examples:

1. Find the OS of the underlined species in Cr2O72-

Sum of oxidation states of constituent elements=overall charge of entire speciesTherefore, 2x + 7×-2 = -2 → x=6: the OS of Cr2 is 6.

2. If MnCl3 is converted to MnO2, determine whether oxidation or reduction has occurred.OS of Mn in MnCl3 is +3; OS of Mn in MnO2 is +4 → increase in OS → oxidation has occurred

Describe and explain galvanic cells in terms of oxidation/reduction reactionsA galvanic cell is an electron pump that produces electricity by pumping electrons out of the anode, where oxidation occurs, into an external circuit (a metallic conductor) and draws them back into the cathode, where reduction occurs.

Outline the construction of galvanic cells and trace the direction of electron flowA galvanic cell consists of:

Two different half cells, consisting of an electrode in electrolyte solution An external circuit, which allows the flow of electrons from the anode to cathode A salt bridge, which allows the migration of ions and maintains electrical neutrality

Define anode, cathode, electrode and electrolyte to describe galvanic cellsAnode: electrolyte at which oxidation occursCathode: electrolyte at which reduction occursElectrode: conductor of a cell (metal or carbon) which gets connected to the external circuitElectrolyte: a substance which in solution conducts electricity

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Perform a first-hand investigation to identify the conditions under which a galvanic cell is producedMethod: Set up a simple galvanic cell and observe the voltage. Observe what happens to the voltage when: the salt bridge is removed; and electrode is removed; two identical half cells are used.Results: 0.7 V with 2 different half cells; 0 V without salt bridge; 0 V without electrode; 0 V with two identical half cells.Conclusion: The essential features of a galvanic cell are: two different half cells; two different electrodes; a salt bridge.

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

Set up a beaker of sulfuric acid. Attach a piece of copper to a voltmeter and place this in the beaker. Attach different metals (Zn, Fe, Pb and Mg) to the other end of the voltmeter and place this end in the beaker as well. Measure the potential difference of various combinations of metals against copper.

OR an alternate method:Set up a standard galvanic cell

with a voltmeter and one half cell of copper in copper sulfate solution; let the other half cell be metal x in its corresponding solution. Measure and record the potential difference. Repeat process with different combinations of metal electrodes (Zn, Fe, Pb and Mg). The reliability of the investigation can be increased by the use of repeat trials for the various combinations you have chosen.Results: Greatest Potential Difference → Least Potential Difference: Mg → Zn → Fe → Pb

Solve problems and analyse information to calculate the potential E° requirement of named electrochemical processes using tables of standard potentials and half equationsProcess with worked example: Calculate the standard cell potential of Co I Co2+ II Ag+ I Ag

1. Determine which metal is oxidised and which is reduced by referring to the standard electrode potentials table; the metal higher up is oxidised.Co is oxidised; Ag is reduced

2. Write the reduction half equation and record E°.2Ag+

(aq) + 2e- → 2Ag(s) E°: +0.803. Write the oxidation half equation by reversing the appropriate reduction equation and record E°

and also reversing its sign.Co(s) → Co2+

(aq) + 2e- E°: +0.284. Add the two half equations to form a redox equation (the electrons should cancel out). Add the

E° potentials (a positive E° indicates a spontaneous reaction)2Ag+

(aq) + 2e- + Co(s) → 2Ag(s) + Co2+(aq) + 2e-

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2Ag+(aq) + Co(s) → 2Ag(s) + Co2+

(aq) E°: 0.80 + 0.28 = 1.08

Gather and present information on the structure and chemistry of a dry cell and evaluate it in comparison to the button cell in terms of: chemistry, cost and practicality, impact on society, impact on environment

Dry Cell Button CellStructure

Consists of a zinc anode, an aqueous paste of ammonium chloride and a mixture of powdered carbon, manganese, and a carbon graphite rod as the cathode.

Consists of a zinc anode, silver oxide cathode with a conductive substance C or Ag mixed with Ag2O, and an alkaline electrolyte solution of potassium hydroxide.

Chemistry O: Zn(s) → Zn2+(aq) + 2e-

R: 2NH4+

(aq) + 2MnO2(s) → Mn2O3(s) + 2NH3(aq) + H2O(l)

O: Zn(s) + 2OH-(aq) → Zn(OH)2(s) + 2e-

R: Ag2O(s) + H2O(l) + 2e- → 2Ag(s) + 2OH-

Cost and Practicality

Cheap; mass produced; useful in devices requiring a small current, e.g. torches, portable radios, clocks; leaking problems common because zinc casing erodes during operation

Expensive due to silver content; but practical because of its small size and constant voltage (1.5 V)

Impact on Society Made the use of portable appliances such as radios possible and feasible.

Has allowed for use of miniature appliances such as hearing aids and watches; its long life means it does not need frequent replacement → practical.

Impact on Environment

Small quantities of zinc, ammonium and carbon are harmless; but not rechargeable and large in size → increase in landfills.

Little environmental impact as it takes up little space in landfill; contents are less likely to leak and less toxic.

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9.2.5 Nuclear chemistry provides a range of materials

Distinguish between stable and radioactive isotopes and describe the conditions under which a nucleus is unstableIsotopes are atoms of one element that differ by having different numbers of neutrons but the same number of protons, hence have a different mass number. An isotope is unstable if:

Its atomic number is greater than 83 Its ratio of neutrons to protons places it outside the zone of stability

The nucleus of a large radioactive atom has high energy, and therefore is unstable; energy is reduced by expelling alpha, beta or gamma particles, and as a result the atom is more stable.

Describe how transuranic elements are producedTransuranic elements are artificial, man made elements with atomic numbers greater than 92. Early transuranic elements were made in nuclear reactors where existing elements were bombarded with slow thermal neutrons which were absorbed to produce new elements. Later transuranic elements were made in particle accelerators where heavy nuclei would be bombarded with positive particles at high speed; since both the target and particles are positively charged, fast speed is needed to overcome electrostatic repulsion.

Describe how commercial radioisotopes are producedRadioisotopes are made in the same way as transuranic elements, i.e. in nuclear reactors or particle accelerators; the only difference is that the target nuclei need not be heavy.

Process information from secondary sources to describe recent discoveries of elementsElements 104-112 (not 108) have been produced using particle accelerators. For example, Element 106, Seaborgium, is made in a particle accelerator involving fusion of an isotope of Californium 249Cf with an oxygen isotope:18

8 O + 249 98Cf → 267 106Sg + 4 1

0n

Identify instruments and processes that can be used to detect radiation Photographic film: radiation is monitored by measuring the extent of the darkening of

photographic film. Cloud chamber: as radiation passes through a supersaturated vapour, it ionises the air, forming

water droplets; the path of these droplets indicates the type of radiation: alpha – straight, dense tracks, beta – less dense, wavy tracks, gamma – faint, random droplets.

Geiger- Muller counter: as radiation passes the Geiger tube, it hits gas molecules and ionises them. Audible electrical pulses are produced as gas molecules are ionised, the rate of these pulses indicates the amount of radiation.

Scintillation counter: a flash of light is emitted as substances are irradiated with radiation.

Identify one use of a named radioisotope: in industry; in medicineTechnetium-99m is used in medicine for diagnosis of blood, heart, brain and thyroid abnormalities. Cobalt-60 is used in medicine for cancer therapy and in industry as a thickness gauge.

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

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Radioisotope Radiation Emitted

Half-life Uses in terms of Chemical Properties

Technetium-99m Gamma 6 hours Medical diagnosis – gamma rays are highly penetrable so can be detected at the body’s surface without invasion; short half life ensures it leaves the body quickly, leaving minimal damage; easily bonds with other chemicals, so it combines with tin to form a serum which is injected into the body and inside the body it binds with red blood cells to detect circulation disorders.

Cobalt-60 Beta; Gamma

5.3 years

Thickness gauge – beta and gamma rays penetrate through metal sheets (but only to a certain degree); relatively long half life makes it suited in machinery since the radioactive source does not need frequent replacement.Cancer therapy – gamma rays kill cancer cells because they contain high energy; half life is short enough to expel reasonable bouts of radiation at moderate intensity to kill cancer cells.

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

Field Benefits ProblemsMedicine Opens a wide range of non-invasive

diagnostic procedures Radiation therapy is, in most cases,

the most effective treatment for cancer

Harmful to people and life forms Causes tissue damage: skin burns,

nausea, radiation sickness Can cause cancer: leukaemia and lung

cancer Genetic damage: deformities in

offspringIndustry Monitoring equipment are more

sensitive and accurate Enables examination of weld and

structural faults in buildings and machinery which otherwise can not be detected

Radiation from equipment can stray and leak if not carefully monitored or stored in well shielded containers

Writing Nuclear Equations:x+y

a+bM ↔ xaP + y

bRSymbols: Neutron: 1

0nProton: 1

1pElectron/ beta particle: 0

-1eAlpha particle: 4

2He

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9.3.1 Indicators were identified with the observation that the colour of some flowers depends on soil composition

Classify common substances as acidic, basic or neutralCommon acids: vinegar, lime/ lemon juice, aspirin, vitamin CLaboratory acids: hydrochloric, sulfuric and nitricNeutrals: waterCommon bases: ammonia, washing soda, antacid tablets, oven/drain cleanersLaboratory bases: all hydroxides and oxides, e.g. NaOH, Mg(OH)2, Fe2O3

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, which is identified by colour changesIdentify data and choose resources to gather information about the colour changes of a range of indicators

Common Indicator Colour in Acid Colour in BaseLitmus Red (<5.0) Blue (>8.0)Phenolphthalein Colourless (<8.3) Pink (10.0)Methyl Orange Red (<3.1) Yellow (>4.4)Bromothymol Blue Yellow (<6.0) Blue (>7.6)

Perform a first-hand investigation to prepare and test a natural indicatorMethod: Heat red cabbage leaves in a beaker under a Bunsen burner. Stop burner after water turns purple and allow to cool. Pour purple water in three test tubes - one containing acid (HNO3), one with water and one with a base (NaOH).Results: Colour in: acid - red; water - purple; base - yellow Identify and describe some everyday uses of indicators including testing of soil acidity/basicity Pool Ranger Home Swimming Pool pH TesterTest: Collect pool sample in tube then add 5 drops of phenol red and shake. Compare tube colour with adjacent colour markings.Importance: Acid or basic pool water results from a buildup of bacteria and pollutants. Pools need to maintain a relatively neutral pH as acidic or alkaline water causes irritations to the skin and eyes. Aquasonic Home Aquarium pH Test KitTest: Collect water from aquarium in tube and add 3 drops of bromothymol blue. Compare test tube colour with chart provided. Add appropriate chemicals (if needed) to maintain pH between 6.0 - 7.8.Importance: pH is lowered when there is a buildup of bacterial wastes. A pH between 6.0 - 7.8 provides the ideal conditions for freshwater aquaria. Extreme changes in pH → reproductive abnormalities, dissolving of scale mucus membrane.Soil pH Test KitTest: Place a teaspoon of sample soil on test plate and add indicator liquid and barium sulfate white powder then stir. Compare colour of sample with colour card provided.

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Importance: Different plants prefer/grow best in different soils with varying pH, e.g. azaleas prefer slightly acidic, vegetables prefer slightly alkaline, no plants grow if pH < 4. Inadequate pH → death, impaired plant growth.

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9.3.2 While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur. The concentration of these acidic oxides have been increasing since the Industrial Revolution

Identify oxides of non-metals which act as acids and describe the conditions under which they act as acidsCO2 – carbon dioxide, SO2 – sulfur dioxide and NO2 – nitrogen dioxide are oxides of non-metals which dissolve in water to form acids and react with bases to form salts and water.

Analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and their acidity/basicity of oxidesAcidic character of elements increases across a period. Elements of the left (metals) form basic oxides; elements in the middle form amphoteric oxides (i.e. they display both acidic and basic character); elements on the right (non-metals) form acidic oxides; noble gases are inert so do not form any oxides.

Define Le Chatelier’s principleIf a system at equilibrium is disturbed, then the system adjusts itself as to minimize the disturbance.

Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principleCO2(g) + H2O(l) ↔ H2CO3(aq) ↔ H+

(aq) + HCO3-(aq) ∆H: negative

More CO2 → shift to the right → more H+(aq) + HCO3

-(aq)

Less CO2 → shift to the left → less H+(aq) → less acidic

Add H+ or HCO3- → shift to the left → more CO2(g) + H2O(l)

Add base → reacts with H+ → shift to the right Add heat→ shift to the left (since reverse reaction absorbs heat) [vice versa]

Note: If ∆H is negative, the forward reaction is releasing heat, if ∆H is positive, forward reaction is absorbing heat.

Identify factors which can affect the equilibrium in a reversible reaction Increase in concentration → reaction shifts to the side which uses up the added species Decrease in concentration (removed) → reaction shifts to the side which produces the removed

species Increase in volume → decrease in pressure → shift to the side which produces the most

gaseous molecules. Note: No change if both sides produce same number of gaseous molecules Decrease in volume → increase in pressure → shift to the side which produces the least

gaseous molecules Increase temperature → shift to the endothermic side that uses up the added heat Decrease in temperature → shift to the exothermic side that produces the removed heat

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogenSource Sulfur Dioxide Oxides of Nitrogen

Natural Geothermal hot springs Volcanoes Smoke from bushfires

High temperatures from lightning (which forms nitric oxide, NO)

Bacteria acting on nitrogenous material

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Bacterial decomposition in soil (which forms nitrous oxide, HNO2)Industrial Burning of fossil fuels (such as coal,

which contains sulfur) Extraction of metals from sulfide ores

(smelting)

Combustion of nitrogen and oxygen in vehicle chambers and power stations (which forms nitric oxide, as in lightning)

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

As bacteria decompose, it forms hydrogen sulfide (H2S) which oxidises to form sulfur dioxide: 2H2S(g) + 3O2(g) → 2SO2(g) + 2H2O(g)

Coal contains sulfur, so when it is burnt, it reacts with oxygen to form sulfur dioxide:S [in compound] + O2(g) → SO2 (g)

High temperatures from lightning or vehicle chambers causes nitrogen and oxygen to combine forming nitric oxide, which further oxidises to form nitrogen dioxide:O2(g) + N2(g) → 2NO(g) then 2NO(g) + O2(g) → 2NO2(g)

Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogenIndustrial Revolution → increased burning of coal + extraction of metals → increased SO2 emissionsElectricity + motor car → high temperature combustion → increased levels of nitrogen oxidesThere is a lack of reliable evidence before 1970 since there was a lack of technology that monitored atmospheric concentrations, and also the technology that existed at the time was insufficient in monitoring such low levels which existed then since the problems were just emerging.

Calculate volumes of gases given masses of some substance in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0°C and 100kPa or 25°C and 100kPaMole Ratio Equations:n = m ÷ fm i.e. number of moles = mass ÷ formula massn = V ÷ mV i.e. number of moles= volume ÷ molar volume which is 22.71 at 0°C/100kPa or 24.79 at

25°C/ 100kPaProcess with worked example: What volume of carbon dioxide gas measured at 0°C/100kPa can be reacted from 15.5g of NaOH to form NaCO3 and water?

1. Write a chemical equation2NaOH(s) + CO2(g) → Na2CO3 (s) + H2O(l)

2. Find the number of moles of the known substancenNaOH = 15.5 ÷ 39.998 = 0.3875…

3. Determine the mole ratio (i.e. moles of unknown : moles of known)CO2: NaOH = 1:2

4. Find the moles of unknown (i.e. multiply moles of known by mole ratio)nCO2 = ½ × nNaOH = 0.3875÷2

5. Convert moles of known into units asked forV= (0.3875÷2) × 22.71 = 4.4 L

Explain the formation and effects of acid rainSO2 dissolves in air to form aqueous droplets of sulfurous acid, which oxidises to form sulfuric acid.

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i.e. SO2(g) + H2O(l) → H2SO3(aq) then 2H2SO3(aq) + O2(g) → 2H2SO4(aq)

Similarly, NO2 dissolves in the air to form nitrous acid and nitric acid.i.e. 2NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq)

Sulfuric acid and nitric acid together form acid rain. Effects of acid rain:

Increases the acidity of lakes → detrimental effect on fish as it upsets their reproductive processes and strips mucus membrane on their scales

Damages pine forests and vegetation → corrodes leaves, changes soil concentration and pH Erodes marble, limestone and metal structures (carbonates in these materials react with acid →

weathering and erosion), e.g. CaCO3(s) + 2H+(aq) → Ca2+

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

Analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environmentBuildup in atmosphere of sulfur dioxide and oxides of nitrogen → formation of acid rain + photochemical smog → detrimental health effects on population (e.g. breathing difficulties) + detrimental effects of acid rain (see previous syllabus dot point). Therefore, the rate of emission of these chemicals needs to be monitored and regulated.

Identify data, plan and perform a first-hand investigation to gather data to measure the mass changes involved and calculate the volume of gas released at 25°C and 100kPaMethod: Weigh an unopened soda can on an electronic balance. Similarly, weigh another soda can that has been refilled with water – this is a control. Open the soda can. Leave both cans overnight. Reweigh add record the mass changes. Calculate the volume of CO2 released.

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9.3.3 Acids occur in many foods, drinks and even within our stomachs

Define acids as proton donors and describe the ionisation of acids in waterAn acid ionises in water to form hydronium ions and donates a proton to form a conjugate base:Acid (aq) → H3O+ + Conjugate Base

Identify acids such as acetic (ethanoic acid), citric (2-hydroxypropane-1,2,3-tricarboxylic acid), hydrochloric and sulfuric acid Acetic acid (vinegar), citric acid (in fruits and as a preservative), hydrochloric (in the stomach and made industrially) and sulfuric acid (in acid rain) are common acids.

Describe the use of the pH scale in comparing acids and bases The pH scale is used to compare the concentration of hydrogen ions in solutions of acids and bases

In a neutral solution, e.g. water, [H+] = [OH-] = 10-7 mol L-1 and so pH =7. In an acidic solution, [H+] > 10-7 mol L-1 and pH < 7. In a basic solution, [H+] < 10-7 mol L-1 and pH > 7.

pH increases as [H+] decreases

Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids In a strong acid, all acid molecules ionise, there are no neutral acid molecules.In a weak acid, only a small percentage of acid molecules ionise, most remain as neutral molecules.The concentration of an acid refers to its molarity; concentrated if it is above 5 mol L-1 and dilute if it is less than 2 mol L-1.

Identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+] To calculate pH given [H+]: pH = -log10 [H+]To calculate [H+] given the pH: [H+] = 10-pH

To calculate [H+] given [OH-]: [H+] = 10-14 ÷ [OH-]As pH increases by 1, the concentration of the hydrogen ions, i.e. [H+], decreases by a factor of ten, or ten fold. E.g. if pH = 1, [H+] = 10-1 = 0.1 but if pH = 2, [H+] = 10-2 = 0.01

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

HCl: strong acid → ionises completely → high [H3O+] → low pHCitric and Acetic: weak acids → only partially ionise → less [H3O+] → higher pH

Describe the difference between a strong and a weak acid in terms of equilibrium between the intact molecule and its ionsAn aqueous solution of a strong acid contains only hydronium ions and the anions of the acid; there are no neutral acid molecules, i.e. the ionisation reaction goes to completion.An aqueous solution of a weak acid is at equilibrium between the neutral acid molecules and hydronium ions and anions of the acid [see ionisation equations in table above].

Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals Using a pH meter/ probe is a non destructive means of measuring the pH of chemicals. Using indicators is a destructive means.

Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids Method: Using a pH probe measure and record the concentration of a strong acid, such as HCl. Rinse the probe with distilled water to avoid cross contamination. Repeat with a weak acid, such as acetic acid.

Gather and process information from secondary sources to explain the use of acids as food additives Acids are added to foods to:

Preserve food – acids lower the pH to a range outside one which microorganisms can survive in, thus enzyme reactions are inhibited or slowed down. e.g. citric acid is a preservative in jams; acetic acid preserves canned beetroot and pickled onions.

Enhance flavour and/or nutritional value – provides a tart, sour taste. e.g. phosphoric acid in soft drinks; citric acid and ascorbic acid are antioxidants.

Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition

Acid or Base Name Chemical Formula Where is it Found?Hydrochloric Acid HCl In the stomachLactic Acid CH3CH(OH)CO2H Milk and YoghurtAmmonia (base) NH3 Formed in the anaerobic decay of organic matterCalcium Carbonate (base) CaCO3 Limestone

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Acid Acetic Citric HydrochloricIonisation Equation

CH3COOH (aq) + H2O (l) ↔ H3O+

(aq) + CH3COO- (aq)

C6H8O7 (aq) + 3H2O (l) ↔ C6H5O7

3- (aq) + 3H3O+

(l)

HCl (g) + H2O (l) → H3O+ (aq) +

Cl- (aq)

pH 2.9 2.1 1.0Degree of Ionisation

1% 8% 100%

Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrationsExample:Calculate the pH of 2.0 mol L-1 solution of sulfuric acid[H2SO4] = 2.0H2SO4 (l) → 2H+

(aq) + SO42-

(aq)

∴ratio of H2SO4 : H+ = 1 : 2∴ [H+] = 2 × 2.0 = 4.0∴pH = -log10 4.0 = -0.6

Process with Harder Worked Example:[A neutralisation reaction] 50 mL of 0.100 mol L-1 hydrochloric acid is added to 75 mL of 0.050 mol L-1 sodium hydroxide solution. Calculate the pH of the resulting solution.

1. Equation states mole ratio. Calculate the moles of H+ from acid information.VHCl = 0.050; CHCl = 0.100C = n ÷ V∴nH+ = C × V = 0.100 × 0.050 = 0.005 moles

2. Calculate the moles of OH- from basic information.VNaOH = 0.075; CNaOH = 0.050∴nOH- = 0.050 × 0.075 = 0.00375 moles

3. Determine which is in excess and by how many moles.H+ is in excessExcess = (moles of H+) – (moles of OH-) = 0.005 – 0.00375 = 0.00125

4. Find the concentration of excess H+ or OH-.[H+] = moles ÷ total volume = 0.00125 ÷ (0.050 + 0.075) = 0.01Note: If OH- is in excess, calculate its excess, then calculate [H+] using [H+] = 10-14 ÷ [OH-]

5. Calculate pH.pH = -log10 [H+] = -log10 0.01 = 2

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9.3.4 Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refined

Outline the historical development of ideas about acids including those of: Lavoisier Davy Arrhenius

Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions

Scientist by Date

Observations Theory of Acids Example

Lavoisier (1780)

Non-metal oxides dissolve in water to produce acids

Acids contain oxygen CO2 (g) + H2O (l) → H2CO3 (aq)

Davy (1815)

Decomposed HCl and found that it did not contain oxygen

Acids contain replaceable hydrogen (ability to be replaced by metals)

Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g)

Arrhenius (1884)

When an electrical current was passed through acid, hydrogen gas evolved at the anode

Acids ionise in solution to produce hydrogen ions

CH3COOH (l) → H+

(aq) + CH3COO (aq)

The Bronsted-Lowry definition [see below] expanded our understanding of acids/bases as it proposes that the acidity of a substance depends on its properties relative to those of the solvent, and not just its structure. This concept furthers our understanding of acid-base equilibrium and pH calculations. Outline the Brönsted-Lowry theory of acids and basesBronsted-Lowry (1923) analysed the experiments of the scientists before them to propose their theory that acids are proton donors and bases are protons acceptors.e.g. HA + H2O → H3O+ + A- where HA is an acid or B + H2O → HB+ + OH- where B is a base

Describe the relationship between an acid and its conjugate base and a base and its conjugate acid Acid (aq) → H3O+ + Conjugate Base when the acid donates a protonBase + H+ → Conjugate acid when a base accepts a proton

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

Strong Base e.g. H2O, NaOH Weak Base e.g. NH3

Strong Acid e.g. HNO3, HCl

Neutral salt (neither the conjugate base of the acid nor the conjugate acid of the base significantly react with water to alter the pH)

Acidic salt (as cation reacts)

Weak Acid e.g. CH3COOH

Basic salt (as anion reacts) Neutral salt

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Identify conjugate acid/base pairs

Base

Conjugate Acid

CO32- HCO3

-

HCO3- H2CO3

NH3 NH4+

OH- H2OH2O H3O+

Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions Amphiprotic substances are those which can react as both a proton donor and a proton acceptor.e.g. water - [as a B.L. acid]: H2O → H+

+ OH-; [as a B.L. base]: H2O + H+ → H3O+

or HCO3- - [as a B.L. acid]: HCO3

- → H+ + CO32-; [as a B.L. base]: HCO3

- + H+ → H2CO3

Identify neutralisation as a proton transfer reaction which is exothermicNeutralisation reactions are exothermic proton transfer reactions.e.g. HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq) – a H+ proton from HCl reacts with OH- in NaCl to form water; heat is liberated → ∆H < 0 (approx. -56kJ mol L-1)

Describe the correct technique for conducting titrations and preparation of standard solutionsPerform 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 Preparing a Standard Solution

1. Measure the mass of the primary standard using a beaker and an electronic balance.2. Dissolve the standard in 100 mL of water.3. Transfer dissolved solution to a volumetric flask using a filter funnel.4. Rinse beaker with distilled water and transfer rinsed solution to volumetric flask. Swirl gently.5. Fill the volumetric flask until the bottom of the meniscus is in line with the calibration mark.6. Invert and shake to ensure a homogenous mixture.7. Calculate the concentration of the solution: n = mass ÷ molar mass then C = n ÷ V (0.250)

To be a primary standard, the substance must be: Of high purity and stability ∴not be volatile, not absorb moisture or react with CO2, as these

would create an impure substance Soluble in water Of accurately known concentration

e.g. Na2CO3 or NaHCO3 but not NaOHConducting a Titration

1. Collect 200 mL of the base solution (e.g. NaOH) in a beaker.

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Acid Conjugate BaseHCl Cl-

H2SO4 HSO4-

NH4+ NH3

H2O OH-

H3O+ H2O

2. Rinse the burette with 10 mL of the base solution, including running the solution through the stopcock.

3. Fill the burette and attach it to the retort stand using a burette clamp. Record the volume of base solution in the burette.

4. Rinse the pipette with 10 mL of the acid solution (e.g. HCl).5. Pipette out 25 mL accurately into a clean, dry conical flask.6. Add 2 drops of phenolphthalein indicator (or another indicator that will show a colour change

as the equivalence point is reached).7. Run the base solution from burette, at the same time, gently swirling the conical flask. As the

equivalence point gets closer, add solution drop by drop. Stop when the solution turns a pale pink colour (for phenolphthalein). [Carry out a rough titration first to get an estimate]

8. Repeat another 2 to 3 times for reliability.9. Determine the concentration of the acid solution, using CiVi = CfVf.

Titration Calculation Process with Worked Example:A student performed a titration of 25.00 mL of acetic acid of unknown concentration with 0.123 mol L-1 solution of sodium hydroxide. Her results are shown right:Calculate the concentration of the acetic acid solution.

1. Write an equation NaOH (aq) + CH3COOH (aq) → NaCH3COO (aq) + H2O (l)

2. Find the moles of one species. Average VolumeNaOH = (22.70 + 22.70 + 22.80) ÷ 3 = 22.73 Note: Trial 1 is a rough titration so do not include this result. nNaOH = 0.123 × 0.02273 = 0.002803. Use mole ratio to find moles of required species. Mole ratio of NaOH : CH3COOH = 1 : 1 ∴nCH3COOH = 0.002804. Find the concentration of the required species. C = n ÷ V = 0.00280 ÷ 0.02500 = 0.112 mol L-1

Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer based technologies Method: Carry out a titration reaction to determine the mass of acetic acid in vinegar with the aid of technologies such as a magnetic stirrer and computer programs that plot the pH as the titrant from the burette is added.

Qualitatively describe the effect of buffers with reference to a specific example in a natural systemA buffer solution is one that contains comparable amounts of weak acid and its conjugate base, so therefore is able to maintain an approximately constant pH even when significant amounts of strong acid or base are added.

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Trial Volume of NaOH used (mL)1 23.302 22.703 22.704 22.80

Example of a naturally occurring buffer: H2CO3 (aq) + H2O (l) ↔ H3O+ (aq) + HCO3

- (aq)

As rainwater falls, it forms a dilute solution of carbonic acid with a pH of 5.7 – 6.0: CO2 (g) + H2O (l) → H2CO3 (aq)

When it lands in the lake or river, comparable amounts of HCO3- from surrounding carbonate rocks,

such as limestone, provide a natural buffer and react with H3O+ in carbonic acid → shifts the equilibrium to the left → removes H+ → raises pH to 6.5 – 7.0:H3O+

(aq) + HCO3-(aq) → H2CO3 + H2O (l)

OH- in water then reacts with H3O+ → shifts equilibrium to the right → removes OH- → restores/ maintains slightly acidic pH:OH-

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

Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions Equipment:

Test tubes Universal Indicator pH chart A range of 0.1 mol L-1 salt solutions

Method: Add 5 mL a salt solution into a test tube. Add 3 drops of universal indicator and record the colour and corresponding pH with reference to a pH chart.Conclusion: From the Bronsted-Lowry theory of acids and bases, the ions in the salt solutions act as acids or bases when reacting with water; acidic solutions donate H+; basic solutions accept H+.

Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spillsAs acids are highly corrosive and bases are highly caustic, it is important to neutralise any spills of these substances if they occur. Amphiprotic substances like HCO3

- in NaHCO3 is suitable for neutralising both acidic and alkaline spills as it can act as a B.L. acid or base:In an acidic spill: H+ + HCO3

- → H2CO3

In an alkaline spill: HCO3- + OH- → H2O + CO3

2-

NaHCO3 is preferred as it is can be safely handled and stored, it is cheap, if excess is used it is not harmful.

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9.3.5 Esterification is a naturally occurring process which can be performed in the laboratory

Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds Alkanols contain a hydroxyl OH- functional group attached to a C atom. General formula: CnH2n+1OHAlkanoic acids are polar, weak acids that contain a carboxylic –COOH group attached to a C atom. General formula: CnH2n+1COOHAlkanols are a subset of alcohols and alkanoic acids are a subset of carboxylic acids. Both alkanols and alkanoic acids only contain C, H and O atoms.

Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures Alkanoic acids have higher melting and boiling points than their corresponding alkanols. Alkanols have relatively high boiling points as O – H bonds are able to form strong hydrogen bonds. But alkanoic acids, due to their –COOH group, have the ability to form 2 hydrogen bonds, creating even higher melting and boiling points as more energy is needed to break 2 bonds than 1.

Identify esterification as the reaction between an acid and an alkanol and describe, using equations examples of esterification Esters (alkyl alkanoates) are formed in a condensation reaction between alkanols and alkanoic acids.i.e. alkanol + alkanoic acid → ester (alkyl alkanoate) + watere.g. ethanol + propanoic acid → ethyl propanoate + water

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

Describe the purpose of using acid in esterification for catalysis As esterification is a moderately slow reaction, a few drops of concentrated sulfuric acid is added to speed up the rate of reaction. Sulfuric acid is a dehydrating agent so absorbs the water produced in the reaction which shifts the equilibrium to the right, thus producing more ester.

Explain the need for refluxing during esterification Refluxing is the process in which any alcohol vapour, which rises in the heat of the reaction, is condensed using a water cooling condenser to prevent the loss of reactants or products. Thus, the reaction is able to be heated to a higher temperature → pushes the equilibrium to the right (as reaction is endothermic) → produces more ester.

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+ + H2O↔

Number of Carbons

Alkyl Part Alkanoate Part

1 Methyl Methanoate2 Ethyl Ethanoate3 Propyl Propanoate4 Butyl Butanoate5 Pentyl Pentanoate6 Hexyl Hexanoate7 Heptyl Heptanoate8 Octyl Octanoate

Refluxing provides a safe alternative to performing the reaction in a closer container, which leads to a toxic build up of gases and possibly an explosion.

Outline some examples of the occurrence, production and uses of estersProcess information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmeticsEsters naturally occur in fruits and flowers and are responsible for their taste and scent. Industrially, esters are artificially produced for practical uses, e.g. ethyl ethanoate is a solvent in nail polish remover; the sweet scent is replicated for use in perfumes, such as apple or pear flavours. Producing esters industrially is cheaper than extracting them from their natural sources.

Identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux Equipment:

15 mL of acetic (ethanoic) acid 15 mL of ethanol Round bottom flask with water cooling

condenser 3 boiling chips

Concentrated sulfuric acid 500 mL beaker Wire gauze Tripod Bunsen burner

Method:1. Place 15 mL of ethanol and 15 mL of acetic acid into the round

bottom flask. Add 3 boiling chips (to ensure even boiling) and 10 drops of concentrated sulfuric acid.

2. Heat the mixture using reflux apparatus for 20 to 30 minutes, until 2 layers are visible. Allow to cool.

3. Add 100 mL of water to the mixture and shake. Allow the 2 layers to separate again.

4. Run this mixture through a separating funnel. Discard the aqueous layer.

5. Add 50 mL of Na2CO3 solution to the remaining mixture in the separating funnel. Shake gently, expelling gas that evolves. Allow the 2 layers to separate again.

6. Run the mixture through the separating funnel again. Discard the lower layer.

7. Repeat steps 5 and 6 to leave a pure sample of the ester, ethyl ethanoate.

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9.4.1 Much of the work of chemists involves monitoring the reactants and products of reactions and managing reaction conditions

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

the variety of chemical occupations a specific chemical occupation for a more detailed study

An analytical chemist working for Sydney Water will use AAS to monitor concentrations of metals, such as Pb and Hg, in water samples which eventually will be supplied to households. The principle of chemistry used by the scientist is that metal electrons move from one energy shell to a higher one by absorbing electromagnetic radiation of a wavelength specific to that metal, giving each metal a unique emission spectrum ∴ metals can be identified and their concentration can be determined even in the presence of other metals. It is important that the scientist monitor and regulate the concentrations of metals in drinking water sources as excessive concentrations can lead to poisoning and illness, e.g. lead poisoning can cause mental retardation, but trace element metals, like Zn, are essential in small concentrations for body functioning.

Identify the need for collaboration between chemists as they collect and analyse data Chemists specialising in various fields of chemistry, e.g. analytical, organic, industrial, need to

collaborate to solve problems requiring a broad spectrum of chemical knowledge. Chemists need to regularly exchange viewpoints and have an open minded but critical approach

to ensure that our scientific knowledge is constantly improving.

Describe an example of a chemical reaction such as combustion, where reactants form different products under different conditions and thus would need monitoringCombustion reactions can produce solely carbon dioxide or a mixture of poisonous carbon monoxide and pollutant soot (C), depending on the amount of oxygen present. Using methane as an example:

Oxygen Supply

Bunsen Burner Hole

Flame Colour

Rate of Combustion Reaction Equation

Sufficient Open Blue Complete combustion; maximum energy released

CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)

Limited Partially open

Mauve Incomplete combustion 2CH4 (g) + 3O2 (g) → 2CO (g) + 4H2O (l)

Blocked Closed Yellow Incomplete combustion CH4 (g) + O2 (g) → C (g) + 2H2O (l)

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9.4.2 Chemical processes in industry require monitoring and management to maximise production

Identify and describe the industrial uses of ammonia Fertilisers (ammonium nitrate) – important for growing food and crops for growing populations;

ammonia is reacted with nitric acid to form ammonium nitrate i.e. NH3 + HNO3 → NH4NO3

Explosives – during WWI, Germany called for more explosives

Identify that ammonia can be synthesised from its component gases, nitrogen and hydrogenHydrogen + atmospheric nitrogen ↔ ammonia i.e. 3H2 (g) + N2 (g) ↔ 2NH3 (g)

Identify the reaction of hydrogen with nitrogen as exothermic∆H = -92 kJ mol-1 (exothermic)

Describe that synthesis of ammonia occurs as a reversible reaction that will reach equilibriumThe synthesis of ammonia is a reversible reaction, meaning that ammonia is formed at the same time as it is being decomposed. Equilibrium is reached when the synthesis (forward) reaction proceeds at the same rate as the decomposition (reverse) reaction.

Explain why the rate of reaction is increased by higher temperaturesIncreasing reactants’ temperatures → increase in their kinetic energy → brings them closer to their activation energy →higher chance of successful collisions → increased reaction rate

Explain why the yield of product in the Haber process is reduced at higher temperatures using Le Chatelier’s principleBy Le Chatelier’s principle, if the temperature is increased, the equilibrium will shift to the side that uses up the heat. Since the synthesis (forward) reaction of ammonia is exothermic (heat is produced), the equilibrium will shift to the left to oppose the change→ decomposition of ammonia → reduced yield of ammonia.

Explain why the Haber process is based on a delicate balancing act involving reaction energy, reaction rate and equilibriumIncreased temperatures → reduced yield but also an increased reaction rate ∴a balanced, compromised temperature is needed to maximise yield and have a moderate reaction rate

Explain that the use of a catalyst will lower the reaction temperature required and identify the catalyst(s) used in the Haber process The addition of an iron catalyst, such as magnetite Fe3O4, lowers the activation energy, which enables a faster reaction rate. Thus, the reaction proceeds at a moderate rate at lower temperatures. The catalyst has no effect on temperature so does not affect the equilibrium.Analyse the impact of increased pressure on the system involved in the Haber process A high pressure (250 × atmospheric pressure) increases both the reaction rate and the yield. Higher pressure → more moles of gas in the closed container → increased chance of successful collision → faster reaction rateHigher pressure → shifts equilibrium to the right (as least gaseous molecules are produced on this side) → more ammonia (increased yield)

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

Temperatures need to be monitored as excessively high temperatures permanently damage the catalyst and low temperatures compromise the optimum yield and reaction rate

Pressure needs to be monitored as high pressure could be explosive Incoming gases lower the reaction’s efficiency and in the case of oxygen, could lead to

explosions if it reacts with hydrogen gas

Gather and process information from secondary sources to describe the conditions under which Haber developed the industrial synthesis of ammonia and evaluate its significance at that time in world historyHaber (and Bosch) discovered that a temperature of 400°C, high pressure of 20MPa and an iron catalyst were the ideal conditions needed to optimise the synthesis of ammonia from hydrogen and nitrogen.Significance:

The synthesis of ammonia was necessary for the production of explosives for Germany in WWI after the British cut off Chilean guano (bird dropping) supplies

Haber’s contributions helped Germany’s war efforts, and even prolonged the war The synthesis of ammonia facilitated the manufacture of fertilisers for food production for

growing populations

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Unknown CationsBa2+, Ca2+, Pb2+, Cu2+, Fe3+, Fe2+

Add Cl- by adding dilute 1 mol L-1 HCl

No precipitate forms

Add H2SO4No precipitate forms

White precipitate forms

Cations:Cu2+, Fe3+, Fe2+

Cations:Ba2+, Ca2+

Add NaOH Conduct flame test

Brown precipitate forms → Fe3+Grey green precipitate forms → Fe2+Blue precipitate forms → Cu2+Apple green flame → Ba2+ Orange red flame → Ca2+

9.4.3 Manufactured products, including food, drugs and household chemicals, are analysed to determine or ensure their chemical composition

Deduce the ions present in a sample from the results of testsPerform a first hand investigation to carry out a range of tests, including flame tests, to identify the following ions:

Phosphate Sulfate Carbonate

Chloride Barium Calcium

Lead Copper Iron

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Unknown AnionsCO32-, SO42-, PO43-, Cl-

Add 2 mol L-1 HNO3Bubbles seen

No bubbles

Silver chloride precipitate forms → Cl-

Anion: CO32- Acidify solution with HNO3 and add Ba(NO3)2

Anions: PO43-, Cl-

No precipitate formsWhite/ pale blue precipitate forms

Anion: SO42-

Add NH3 → makes solution basic, then add aqueous Ba(NO3)2

White precipitate forms → PO43- Acidify by adding HNO3, then add aqueous AgNO3

Gather, process and present information to describe and explain evidence for the need to monitor levels of one of the above ions in substances used in societyPhosphate

Low concentrations are essential in waterways for normal aquatic plant growth High phosphate concentrations → algal bloom → algae grows to cover lake/river surface → kills

fish + uses up all the phosphate → then, algae decays anaerobically → low oxygen levels in water → death of marine life in the water

Monitoring the amount of phosphate entering and already present in waterways, scientists can guard against algal blooms

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

Process:1. A hollow cathode lamp emits light of a specific wavelength in a metal’s emission spectrum.2. The sample to be analysed is fed into a flame which vaporises molecules and ions, converting

them into atoms.3. A monochromator selects a particular wavelength and directs it to a photomultiplier/ detector.4. The photomultiplier measures the intensity of the light. By comparing the intensity absorbed

with and without a sample, the concentration of a particular element can be calculated.5. Measure the absorbance of several standard solutions and plot a calibration graph.

Principle of Chemistry:An electron in a metal atom moves from one energy shell to a higher one by absorbing electromagnetic radiation of a specific wavelength ∴the wavelength is unique to the metal → the metal has its own unique absorption or emission spectrum. This fact is critical to the success of AAS as elements can be distinguished, and their concentration can be found, even in the presence of other elements.

Trace Elements: Elements present in very small but essential amounts Help enzymes function to maintain optimum nutrition and functioning in soils and organisms Prior to AAS, trace elements went unnoticed as analytical methods were not sensitive enough to

measure their minute concentrations Alan Walsh’s discovery of AAS has made it possible to monitor and maintain safe but essential

amounts of trace elements → nutrient soils for maximum yield of crops and produce + optimum functioning of organisms

Gather, process and present information to interpret secondary data from AAS measurements and evaluate the effectiveness of this in pollution control

AAS’ ability to measure very low concentrations of a wide range of elements makes it an effective means of monitoring pollution.

Water and soil samples can be analysed to identify the presence of specific trace elements [refer to principle of chemistry] and their concentrations.

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Identify data, plan, select equipment and perform first hand investigations to measure the sulfate content of lawn fertiliser and explain the chemistry involvedIonic equation: Ba2+ (aq) + SO4

2- (aq) →BaSO4 (s)

Method: 1. Dissolve 0.5 g of fertiliser in 25 mL of deionised water – use deionised water as sulfate

contaminants in normal water may alter the results.2. Add 5 drops of HCl – HCl breaks down insoluble sulfate complexes.3. In another beaker, add 25 mL of Ba(NO3)2. 4. Heat both solutions until near boiling – heating increases the kinetic energy → increases rate of

successful collisions → increases particle size.5. Add Ba(NO3)2 drop by drop to the hot fertiliser solution, constantly stirring – adding Ba(NO3)2

slowly allows BaSO4 precipitate crystals to form.6. Allow to cool for 30 minutes and then add 5 mL of acetone – acetone acts as a coagulating

agent. 7. Filter through a sintered glass funnel.8. Wash the precipitate 3 to 4 times and then dry using a vacuum pump and more acetone –

washing the precipitate several times minimises mass loss; drying using a vacuum pump and acetone ensures no water is mistaken for BaSO4.

9. Weigh the synthesised BaSO4. 10. Repeat steps 8 and 9 until a constant mass is reached. 11. Determine the amount of sulfate present.12. Calculate percentage mass of sulfate in fertiliser sample, using:

fm (SO4) ÷ fm (BaSO4) × mass of BaSO4 weighed × 100

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

Measures taken that ensure reliability Sources of error that compromise reliabilityAdding HClHeating the solutionsAdding BaNO3 very slowlyAdding acetoneWashing the precipitate 3 to 4 timesUsing a vacuum pumpReweighing until a constant mass was reached

Precipitate adheres to walls of apparatus or is spilt so is lost when transferring solutionsContamination by substances absorbed from the solution during precipitationIncomplete drying, leaving water in the final BaSO4 precipitate

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9.4.4 Human activity has caused changes in the composition and the structure of the atmosphere. Chemists monitor these changes so that further damage can be limited

Describe the composition and layered structure of the atmosphereLayer Altitude

(km)Composition Temperature

Range (°C)Troposphere 0(sea

level) – 15

Contains 80% of the atmosphere’s massMajor constituent gases: nitrogen (78.08%), oxygen (20.95%) and argon (0.93%) and water vapour (0.5-5%)Minor constituent gases/pollutants: methane, carbon monoxide, carbon dioxide, sulfur dioxide, nitrogen oxides and ozone (0.02 ppm)

15 – -50

Stratosphere 15 – 50 Same percentage composition of gaseous molecules as the troposphere but greater spacing between gaseous particles → low pressure. Water vapour concentration drops – 5 ppm.

-50 – 0

Mesosphere 50 – 85 As altitude increases, spacing between molecules becomes greater → pressure decreases exponentially. Major constituent gases: nitrogen, oxygen and carbon dioxide.

0 – -100

Thermosphere 85 – 500 Gases are widely spread apart and include electrically charged ions, such as O2

+ and NO+, and atoms, such as O, and free electrons which would not be stable in the lower atmospheric layers.

-100+

Identify the main pollutants found in the lower atmosphere and their sourcesPollutant Main Source(s)

Carbon monoxide Incomplete combustion of fossil fuels in cars, e.g. incomplete combustion of petrol: C8H18(l) + 6O2(g) → CO2(g) + CO(g) + 5C(s) + 9H2O(l)

Nitrogen oxides High temperature combustion, e.g. power stations, internal combustion engine:N2(g) + O2(g) → 2NO(g) then 2NO(g) + O2(g) → 2NO2(g)

Sulfur dioxide Smelting of sulfide oresCFCs Refrigerants, air conditioning, halon fire extinguishers

Describe ozone as a molecule able to act both as an upper atmosphere UV radiation shield and a lower atmosphere pollutantIn the stratosphere, ozone acts as a radiation shield by absorbing, and thus filtering out, short wavelength ultraviolet UVB and UVC radiation from the Sun. A balanced cycle of ozone formation and decomposition reactions converts dangerous UVB and UVC radiation to heat energy and maintains a steady concentration of ozone in the stratosphere:

1. Ozone is formed when oxygen molecules (O2) absorb energy from short wavelength UVC radiation forming energised oxygen atoms: O2(g) + UV light →2O(g)

2. These free radicals rapidly combine with oxygen molecules to form ozone:

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Coordinate covalent bond

O2(g) + O(g) → O3(g) [∆H = -106 kJ/mol – exothermic]3. Ozone decomposes back into oxygen molecules and oxygen atoms as it absorbs the longer

wavelength UVB and remaining UVC radiation: O3(g) + UV light → O2(g) + O(g)

Describe the formation of a coordinate covalent bondDemonstrate the formation of coordinate covalent bonds using Lewis electron dot structuresA coordinate covalent bond is formed when one atom donates both a pair of electrons to another, resulting in a complete valence electron shell in both atoms. E.g. in the formation of ozone:

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

Oxygen OzoneStructure

Melting/ Boiling Point

Lower m.p. and b.p. than ozone, due to: Lower molecular mass Weak dispersion forces Non-polar molecule

Higher m.p. and b.p. than oxygen, due to: Higher molecular mass Polar molecule Strong intermolecular forces ∴more

energy is needed to overcome these forces

Reactivity Less reactive than ozone, due to its double covalent bond which requires lots of energy to be broken

More reactive than oxygen, due to its coordinate covalent bond which requires much less energy to be broken

Compare the properties of the gaseous forms of oxygen and the oxygen free radicalOxygen free radical has unpaired electrons in its incomplete valence electron shell → very unstable and highly reactive → more reactive than oxygen and ozone

Identify the origins of chlorofluorocarbons (CFCs) and halons in the atmosphere CFCs are compounds containing chlorine, fluorine and carbon only, they do not contain hydrogen. Until 1996, CFCs were released into the atmosphere through their use in refrigerants, fire extinguishers and foam plastics as they were relatively cheap, non-toxic and inert.Halons are compounds containing carbon, bromine and other halogens. They were used as BCF fire extinguishers in cars and boats until 1994.

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

1. Count the number of carbons to determine the alkane, e.g. methane, ethane.2. Name, in alphabetical order, any groups attached to the alkane, e.g. Br – bromo, Cl – chloro, F –

fluoro, I – iodo.

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3. Count how many of each group is present, then denote the number using prefixes, e.g. 2 – di, 3 – tri, 4 – tetra.

4. State the position of each group along the carbon chain, numbering using the lowest numbers.Note: Use hyphens (-) between numbers and words and commas between numbers and do not have any space between words, e.g. 2-bromo-3-chloropentane

Discuss the problems associated with the use of CFCs and assess the effectiveness of steps taken to alleviate these problemsThe reaction of the chlorine atom from CFCs with ozone destroys many ozone molecules resulting in a thinning of the ozone layer. Ozone depletion → more UV radiation reaching the Earth’s surface which results in:

Increased incidence of skin cancer and sunburn Increased risk of eye cataracts and irritations if its concentration in the lower atmosphere is

greater than 0.3 ppm Lowering of immune response → increased risk of disease and illness UV radiation interferes with photosynthesis → reduced plant growth

To alleviate these problems, countries have phased out the use of CFCs and halons and provide alternatives like HFCs and HCFCs [refer to next syllabus dot point].

Present information from secondary sources to identify alternative chemicals used to replace CFCs and evaluate the effectiveness of their use as a replacement for CFCs

HFCs and HCFCs are alternatives replacing CFCs in refrigerants, air conditioners and aerosols. The C-H bonds in HCFCs and HFCs are reactive so they readily react and decompose to a

significant extent in the troposphere ∴ do not diffuse into the stratosphere to damage the ozone.

HCFCs contain C-Cl bonds which form chlorine atoms which contribute to ozone depletion so are only a temporary substitute for CFCs.

HFCs do not contain C-Cl bonds ∴ no Cl available to react with ozone ∴ does not contribute to ozone depletion.

HFCs are more expensive and less efficient than CFCs.

Analyse the information available that indicates changes in atmospheric ozone concentrations, describe the changes observed and explain how this information was obtainedTotal Ozone Mapping Spectroscopy (TOMS)

Satellites situated above the atmosphere Measures ultraviolet radiation backscattered off the earth and compared to incoming radiation Measures total column ozone Covers the entire planet and produces contour maps of total ozone over large areas of the

Earth’s surfaceDobson Spectrometer

Situated on the ground Measures the difference between UV levels coming into the earth from the sky at a frequency

which is absorbed by the ozone and a frequency which is not absorbed by ozone Measures total column ozone and profiles of ozone in the atmosphere Provides measurements for a small area only

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Present information from secondary sources to write the equations to show the reactions involving CFCs and ozone to demonstrate the removal of ozone from the atmosphereWhen CFCs, like CCl3F, reach the ozone layer, they undergo the following reactions:

1. Ultraviolet light breaks CCl3F into a chlorine atom and CCl2F: CCl3F(g) + UV light → CCl2F(g) + Cl(g)

2. The chlorine atom reacts with an ozone molecule to form chlorine monoxide and oxygen:Cl(g) + O3(g) → ClO(g) + O2(g)

3. Chlorine monoxide reacts with an oxygen atom to form a chlorine atom and an oxygen molecule: ClO(g) + O(g) → Cl(g) + O2(g)

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9.4.5 Human activity also impacts on waterways. Chemical monitoring and management assists in providing safe water for human use and to protect the habitats of other organisms

Identify that water quality can be determined by considering: Concentrations of common ions Total dissolved solids Hardness Turbidity

Acidity Dissolved oxygen and biochemical

oxygen demand

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

Water Property

Testing Procedure Water Quality Information

Concentrations of common ions

Quantitative test – Na+, Ca2+, Mg2+ are measured by AAS

Qualitative test – PO43- is present if

adding ammonium molybdate [(NH4)2MoO4] to an acidified water sample produces a yellow precipitate

High concentrations of Na+ and Cl- indicate salinity → important to monitor as aquatic organisms only survive in a narrow range of salt concentration

Mg2+ and Ca2+ indicate water hardness PO4

3- is tested to monitor algal blooms and eutrophication

Total dissolved solids (TDS)

Measured by using a conductivity meter:

1. Calibrate the conductivity meter2. Insert the electrodes into water

sample and swirl once.3. Read measurement.

TDS increased by irrigation and industrial and sewage effluent discharge into waterways

TDS in drinking water needs to be < 500 ppm

If TDS > 1000 ppm → freshwater marine bugs/ plankton shrivel and die

Hardness Qualitative test – add a small soap scraping into hard and soft water samples, shake both vigorously and observe for lathering/scum; hard water does not lather

Quantitative test – volumetric titration to measure the concentration of Mg2+ and Ca2+ ions; the result is expressed as mg/L of CaCO3 or MgCO3

High concentrations of Mg2+ and Ca2+ indicate hard water

Hard water produces CaCO3 and MgCO3 a grey scale which sticks to baths and sinks

Turbidity 1. Pour water sample into a turbidity tube in small volumes

2. Repeat until the black markings can not be distinguished.

3. Read the volume below the water level.

Turbidity is a measure of water cloudiness – a greater amount of suspended solids → higher turbidity

Caused by clay/silt from soil drainage, algae and fine organic matter

Prevents light penetration → reduced photosynthesis → lesser quantity of oxygen → decreased plant growth →

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limits food supply to marine animals → decreases biodiversity. Also, suspended particles clog gills of fish → suffocation → improper reproductive development

Acidity Determine pH of water using a pH meter

Determine pH using test strips and comparing colours to a standard colour chart

pH levels outside 6.5 – 8.2 indicates pollution by limestone, leaf litter, detergents and fertilisers

Acidic → corrosion of water pipes and tanks + fish suffering from skin irritations, impaired gill functioning and/or ulcers

Dissolved Oxygen (DO)

1. DO sample bottle filled and recapped while immersed in test water; ensure no air bubbles are present.

2. Add 2 drops of MnSO4 and 2 drops of KI. Recap and mix.

3. Brown precipitate appears – light brown → low DO; dark brown → high DO.O2(g) + 2Mn2+

(aq) + 4OH-(aq) →

MnO(OH)2(s)

4. Add 8 drops of sulfuric acid → forms a yellow iodine solution.

5. Titrate solution with sodium thiosulfate using a starch indicator.

Dissolved O2 is essential for respiration in aquatic life

DO increases through photosynthesis of aquatic plants and algae; and wave action near the water’s surface

DO decreases in high temperatures; high salinity; as plants and animals respire or decompose

DO in clean water = 7-9 ppm; DO in polluted water < 5 ppm

Biochemical Oxygen Demand (BOD)

1. Collect 2 samples of test water.2. Determine DO of one sample

immediately using Winkler method [above].

3. Store the second sample in the dark for 5 days.

4. Carry out Winkler method on second sample – the difference in measurement between the two samples is the BOD

BOD assesses organic pollution as it measures the concentration of dissolved oxygen that is needed for the complete breakdown of the organic matter in the water by aerobic bacteria.

Identify factors that affect the concentrations of a range of ions in solution in natural bodies of water such as rivers and oceansPathway from rain to water body

If rain runs off bushland into streams, it only picks up small quantities of Ca2+, Mg2+, NO3- and

PO43- from surface nutrients and decomposing minerals

If water penetrates down to underground aquifers, it will contain increased concentrations of these ions, and in addition also have SO4

2-, Cl- and CO32-

If water percolates even further into deep underground aquifers, the water would contain even higher concentrations of these ions, and in addition also have Fe3+, Mn2+, Cu2+ and Zn2+

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The Nature and Amount of Human Activity in the Catchment Land clearing → destabilises soils → increased sediments carried with the water → increases

turbidity Increased fertiliser use on agricultural land, the water run-off will have increased

concentrations of PO43- and NO3

- and organic wastes from animals

Effluents Discharged into Water Bodies Industrial effluents discharge heavy metal ions, such as Pb2+, Hg2+, Cu2+, Zn2+

Raw or treated sewage discharge introduces PO43- and NO3

-

Describe and assess the effectiveness of methods used to purify and sanitise mass water supplies1. Flocculation – Lime [Ca(OH)2] is added to increase pH → encourage the formation of a

precipitate. Iron (III) chloride is added, it absorbs suspended particles and bacteria to produce a gelatinous precipitate of iron hydroxide. This step is cost effective and relatively fast.

2. Sedimentation – the floc formed settles at the bottom of the tank for removal. This step reduces running costs as it is based on gravity but it is slow ∴not efficient.

3. Filtration – the water is filtered through a sand and gravel bed to remove remaining suspended particles. This step does not remove dissolved particles and microscopic pathogens.

4. Chlorination – chlorine reacts with water to produce hypochlorous acid (HOCl) which kills residual bacteria and prevents algal growth as water travels down pipes to households. Chlorination is cost effective and kills most bacteria but in 1998 Sydney residents had to boil their water due to a Cryptosporidium and Giardia outbreak.

Describe the design and composition of microscopic membrane filters and explain how they purify contaminated waterA membrane filter is a thin film of a synthetic polymer, like polypropylene, through which there are 0.2 μm pores. Sheets of this polymer are pleated around a porous core and held with mesh. As contaminated water enters through an opening cavity, it is forced by gravity, vacuum pumps and centrifugal pressure through the microscopic pores of the polymer, leaving behind harmful microorganisms. Clean water passes out a central cavity.

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

Heavy Metal Pollution Heavy metals are those of relatively high atomic mass such as the transition metals and lead

and arsenic. Mercury, lead, cadmium, chromium and arsenic are the heavy metals that need monitoring in

water as they are poisonous to humans. AAS is used to monitor heavy metal concentrations due to its advantage of being able to assess

the concentration of several metals in one sample, provided there is an adjustable lamp. [Refer to the principle of chemistry used in AAS on page 33]

Gather, process and present information on the features of the local supply in terms of: Catchment area Possible sources of contamination in this catchment Chemical tests available to determine levels and types of contaminants

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Physical and chemical processes used to purify water Chemical additives in the water and the reasons for the presence of these additives

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9.6.1 The chemical composition of the ocean implies its potential role as an electrolyte

Identify the origins of the minerals in oceans as: leaching by rainwater from terrestrial environments hydrothermal vents in mid-ocean ridges Leaching involves rainwater, streams and rivers dissolving soluble minerals from terrestrial

rocks and soils and carrying them to oceans Seawater seeps through mid-oceanic ridge cracks → gets heated by upwelling magma → hot

water dissolves ionic substances → hot, mineral-rich water is vented back into the ocean → water is cooled and minerals crystallise to deposit on the ocean floor.

Outline the role of electron transfer in oxidation-reduction reactions Oxidation-reduction reactions involve the transfer of electrons from the oxidation site (anode) to the reduction site (cathode). Ions in solution act as an electrolyte/salt bridge → allows for migration of ions → ability to conduct electricity.

Identify that oxidation-reduction reactions can occur when ions are free to move in solid and liquid electrolytes Ionic substances in solid form cannot conduct electricity (i.e. be an electrolyte) as they are in fixed positions in a crystal lattice. The ions need to be free to move to conduct electricity and be apart of oxidation-reduction reactions.

Describe the work of Galvani, Volta, Davy and Faraday in increasing understanding of electron transfer reactions Process information from secondary sources to outline and analyse the impact of the work of Galvani, Volta, Davy and Faraday in understanding electron transfer reactions

Scientist/ Date Experiments ContributionGalvani (1791) Connected a wire from frog’s spinal

cord to a brass hook which was in contact with an iron railing to observe continual muscle contraction.

Credited with the generation of electric current.Electric fluid in animals caused muscle contraction – animal electricity.

Volta (1800) Repeated Galvani’s experiment using different metals.Constructed the first voltaic pile (galvanic cell) by sandwiching electrolytic cardboard discs soaked in brine between two different metal plates (Zn and Cu), then attaching a wire to each end to produce an electric current.

Demonstrated that two different wires in solution generated an electric current to cause muscular contraction.

Davy (1807) Constructed the largest galvanic cell and passed a strong electric current

Recognised that chemical reactions were the source of electric current.

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through molten salts of various compounds suspected of containing undiscovered elements.

Identified new elements: Group I metals – Na, K; Group II metals – Ba, Ca, Sr, and Mg.

Faraday (1830) Extended on Davy’s electrolytic decomposition studies.Investigated quantitative relationships in electrolytic processes.

Established electrolysis laws, e.g. the quantity of a substance produced during electrolysis is directly proportional to the quantity of electric charge passed through the circuit.Developed the terminology used today.

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9.6.2 Ships have been made of metals or alloys of metals

Account for the differences in corrosion of active and passivating metals Corrosion is the degradation of metals as it reacts with environmental substances, like oxygen, so that it loses its strength and becomes unable to fulfil its intended purpose.Passivating metals, e.g. aluminium, form an impervious oxide layer as the metal reacts with oxygen and water when first exposed. This passivating layer prevents more oxygen and water penetrating through to the unaffected metal underneath → provides protection from further corrosion.Active metals, e.g. impure iron, form a porous, flaky oxide layer which water and oxygen can easily penetrate to continue corrosion until the metal has corroded completely.

Identify iron and steel as the main metals used in ships Iron and steel (an iron alloy) are the main metals used in ships.

Identify the composition of steel and explain how the percentage composition of steel can determine its properties Gather and process information from secondary sources to compare the composition, properties and uses of a range of steel

Substance Composition Properties in relation to composition Uses in relation to properties

Pure Iron 100% Fe Soft; malleable; corrodes very slowly as impurities which act as cathodes for corrosion are absent

Not used commercially

Mild Steel <0.2% C Added carbon makes it relatively soft but harder than pure iron; more brittle but relatively malleable; corrodes quickly

Carbodies – malleability allows for easy moulding

Stainless Steel

4-30% Cr1-22% Ni1-10% Mn

Chromium – corrosion and heat resistanceNickel and Manganese – increased strength and hardness

Kitchen sinks and cutlery – does not corrode in the presence of acids/ bases from foods and detergents

Tool Steel 5% Co4% Cr14% W0.5% Mn

Chromium and manganese – corrosion resistanceCobalt and Tungsten – increased strength and hardness

High speed cutting tools – does not go soft when heated due to strength and hardness

Describe the conditions under which rusting of iron occurs and explain the process of rustingUse available evidence to analyse and explain the conditions under which rusting occurs Rusting EquationsOxidation: Fe(s) → Fe2+

(aq) + 2e-(aq)

Reduction: O2(g) + 2H2O(l) + 4e-(aq) → 4OH-

(aq)

Redox equation: 2Fe(s) + O2(g) + 2H2O(l) → 4OH-(aq) + 2Fe2+

(aq) → Fe(OH)2(s)

Fe(OH)2 (s) further oxidises to form rust, Fe2O3. H2O, i.e. Fe(OH)2(s) + O2(g) → Fe2O3. H2O + 2H2O

Conditions for Rusting

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Oxygen – undergoes reduction, forms the cathode Water – acts as a salt bridge as it is an electrolyte → allows for the migration of ions, i.e. Fe2+

moves to cathode, OH- moves to anode

Factors that Accelerate Rusting Salt water – contains Na+ and Cl- ions → a better electrolyte → faster migration of ions Points under stress – e.g. sharpened points, bends, heads; forms a site for oxidation as the

crystal lattice is distorted to expose more iron atoms Iron in contact with a less active metal – e.g. tin, copper; less active metal becomes the

cathode, forcing iron to be the anode

Identify data, select equipment, plan and perform a first-hand investigation to compare the rate of corrosion of pure iron and an identified form of steel Equipment:

Stainless steel Mild steel 0.1 mol L-1 NaCl 2 test tubes

Method:1. Place pieces of stainless steel and mild steel (of a relatively similar size and shape) into separate

but identical test tubes.2. Fill both test tubes with 25 mL of 0.1 mol L-1 NaCl.3. Leave for 7 days, then record observations and degree of rusting, i.e. look for the amount of

red-brown deposit.The mild steel piece corrodes more rapidly than stainless steel as there is more red brown deposit. The carbon content in mild steel provides sites for reduction → galvanic cell is produced → corrosion occurs. The chromium content in stainless steel makes it corrosion resistant.

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9.6.3 Electrolytic cells involve oxidation-reduction reactions

Describe, using half equations, what happens at the anode and cathode during electrolysis of selected aqueous solutionsAnode (+)Oxygen gas produced at the inert anode if the anions present are more difficult to oxidise than water or hydroxide ions. Sulfate, nitrate, carbonate or phosphate ions are never oxidised at the anode.Cathode (-)Hydrogen gas produced at the inert cathode if the cations present are more difficult to reduce than water or hydrogen ions. Potassium, sodium, magnesium or aluminium ions are never reduced at the cathode.

Describe factors that affect an electrolysis reaction, including: Effect of concentration Nature of electrolyte Nature of electrodes

Factor Effect on Electrolysis ReactionConcentration of Electrolyte

High concentration of ions improves conductivity → increased rate of electrolysis.Changing concentrations of species which have similar standard potentials → different products. E.g. In molten or concentrated aqueous NaCl solutions (>2 mol L-1), where Na+ and Cl- are present but there is no/ very little H2O, Na+ reduces and Cl- oxidises. In dilute NaCl solutions, Na+ reduces but H2O oxidises.

Nature of Electrolyte The electrolyte itself affects the products of the reaction, e.g. CuCl2, CuSO4 and Na2SO4 each produce different products. Both, one or neither of the ions in the electrolyte may be reacted, depending upon their ease of oxidation or reduction relative to water.

Nature of Electrodes The greater the surface area of the electrodes, the greater the conductivity.The smaller the distance between the electrodes, the greater the conductivity.

Plan and perform a first-hand investigation and gather first-hand data to identify the factors that affect the rate of an electrolytic reactionTesting the Effect of Different ElectrolytesMethod:

1. Half fill a 150 mL beaker with 1 mol L-1 solution of potassium bromide.2. Insert two graphite electrodes and connect these to a voltameter which is set at 4-6 volts.3. Pass a current through the solution for about one minute and observe colour changes around

each electrode.4. Repeat steps 1-3 with solutions of potassium iodide and copper sulfate.

Testing the Effect of the Nature of ElectrodesMethod:

1. Half fill a 150 mL beaker with 1 mol L-1 solution of copper sulfate.

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2. Insert a copper anode and graphite cathode and connect these to a voltameter which is set at 4-6 volts.

3. Record observations to the electrodes and the solution.4. Repeat steps 1-3 using two graphite electrodes.5. Compare results of both experiments.

Risk Assessment: Safety goggles were worn to protect the eyes from any splashes of chemicals which would cause burns and irritations.

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9.6.4 Iron and steel corrode quickly in a marine environment and must be protected

Identify the ways in which a metal hull may be protected, including: Corrosion resistant metals Development of surface alloys New paints

Method How Method Prevents CorrosionUsing corrosion resistant metals

Aluminium and chromium form passivating oxide layers which are impervious and tightly bound to the steel → prevents exposure to oxygen and water. However, chromium oxide layer reacts with chloride ions in sea water so is not suitable for ships’ hulls.

Surface alloys

Formed by the bombardment of metal ions (e.g. chromium or nickel) onto the surface of steel where they become embedded as atoms. This alloyed surface is corrosion resistant and more economical than using solid pieces of steel alloys.

New paints

Forms a physical barrier that stops oxygen and water coming into contact with the metal. New polymer based paints form a pyroaurite interlayer between the paint and metal. This is a highly insoluble, ionic layer that is tightly bound to the metal surface → prevents the migration of ions from one place on the steel to another → prevents corrosion.Zinc rich paints prevent corrosion by providing cathodic protection as zinc is a sacrificial anode.

Predict the metal which corrodes when two metals form an electrochemical cell using a list of standard potentialsWhen comparing two metals, the one with the algebraically smaller (higher up on the SPT) standard potential will corrode (oxidise) more easily than the other.For example, iron has Eѳ = -0.45V and tin has Eѳ = -0.14V ∴ iron is more readily oxidised/corroded (as in galvanic cells).

Outline the process of cathodic protection, describing examples of its use in both marine and wet terrestrial environmentsGather and process information to identify applications of cathodic protection, and use available evidence to identify the reasons for their use and the chemistry involved [refer to next syllabus dot point aswell]Cathodic protection is a method of protecting a metal from corrosion by making it the cathode of a galvanic cell. There are two types – sacrificial anode or applied voltage.

Sacrificial AnodeA sacrificial anode is a block of a more active metal than iron, e.g. zinc or magnesium, which is in contact with the ship’s hull via an electrolytic medium, e.g. salt water or moist earth, which acts as a salt bridge. The active metal corrodes preferentially to the iron; iron is made the cathode ∴ it does not oxidise and corrode. This method works well in a marine environment as the ship’s hull is immersed in a sea water electrolyte. It does not work well in terrestrial environments as there is no electrolyte/salt bridge → no galvanic cell.

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Applied VoltageThe hull is attached to the negative terminal of a voltage source (4-5V) → supplied with electrons → becomes the cathode → prevents oxidation/corrosion. An inert electrode, e.g. platinum, is mounted on the ship’s hull and connected to the positive terminal → forms the anode. The inert electrode is insulated from the ship’s hull, preventing the migration of ions ∴ a galvanic cell cannot form. This method is used in terrestrial environments to prevent corrosion in underground storage tanks and pipelines.

Describe the process of cathodic protection in selected examples in terms of the oxidation- reduction chemistry involvedIron: Fe2+ + 2e- → Fe Eѳ = -0.45V Zinc: Zn2+ + 2e- → Zn Eѳ = -0.76VZinc has a lower Eѳ than iron → zinc oxidises/corrodes more easily than iron → zinc is used as sacrificial anode by galvanising (coating) steel with zinc → provides cathodic protection, i.e. zinc corrodes preferentially, protecting the iron (steel).Even if Fe2+ ions are formed, Zn reacts with it to convert it back to Fe, i.e. Fe2+ is reduced:Zn + Fe2+ → Zn2+ + FeWater is reduced at the iron cathode (ship’s hull): 2H2O + O2 +4e- → 4OH-

Coating steel with tin is less effective as tin is less reactive than iron → if tin is scratched, iron becomes the anode and tin becomes the cathode → accelerates the rate of corrosion.

Identify data, gather and process information from first-hand or secondary sources to trace historical developments in the choice of materials used in the construction of ocean-going vessels with a focus on the metals used

Date Metal Used Use of Metal

Advantages Disadvantages

1760s Copper Sheathing on wooden hulls

Poisonous → deters marine organisms

Sheathing attached by iron nails forms galvanic cell → corrosion

1830s Iron Body of ships

Malleable – moulded into sheetsOvercame wood rot problemsStronger/Safer/Faster/Durable

Fouling due to marine growths

1880s-1890s

Steel Body of ships

Lighter than iron → more cargoLater, increased manganese + decreased carbon → increased usage

ExpensiveLow manganese + high phosphorous and sulfur → brittle, less ductile, not strong

1950s Aluminium Hull; military vessels

Light weight → faster shipsStrongMalleable – readily castCorrosion resistant

More expensive than steel → did not meet yield requirementsPassivating aluminium oxide layer does not adhere well in marine environments

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Identify data, choose equipment, plan and perform a first-hand investigation to compare the corrosion rate, in a suitable electrolyte, of a variety of metals, including named modern alloys to identify those best suited for use in marine vesselsEquipment:

Brass Stainless steel

Mild steel Aluminium

0.1 mol L-1 NaCl 4 test tubes

Method:1. Place a piece of brass into a test tube half filled with 0.1 mol L-1 NaCl.2. Repeat step 1 with pieces of stainless steel, mild steel and aluminium that are of a relatively

similar size and shape to brass.3. Leave the 4 test tubes in a room for 7 – 14 days.4. Observe the test tubes each day for signs of corrosion, i.e. the amount of red-brown deposit.5. Label the metals/alloys from 1-4 in order of their degree of corrosion, 1 being the least

corrosive and 4 being the most corrosive.

Plan and perform a first-hand investigation to compare the effectiveness of different protections used to coat a metal such as iron and prevent corrosionMethod:

1. Collect a range of iron samples which have been treated differently for protection from corrosion: e.g. painted, lacquer with inhibitor, galvanised with zinc, coated with wax.

2. Use a control of an untreated iron piece.3. Record observations of signs and the extent of corrosion of both the treated and untreated

samples of iron, e.g. look for colour changes/ rusting – red brown deposits/ dullness/ scratching.

4. Compare results of the treated samples and the control to analyse the effectiveness of each form of protection.

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9.6.5 When a ship sinks, the rate of decay and corrosion may be dependent on the final depth of the wreck

Outline the effect of: Temperature Pressure

on the solubility of gases and salts The solubility of gases decreases with increasing temperature. The solubility of gases increases with increasing pressure. The solubility of salts increases with increasing temperature. The solubility of salts increases with increasing pressure.

Identify that gases are normally dissolved in the oceans and compare their concentrations in the oceans to their concentrations in the atmosphereThe most abundant dissolved gases are nitrogen, oxygen and carbon dioxide. The solubility of a gas in the ocean is directly proportional to the partial pressure of that gas above the ocean, i.e. in the atmosphere. An exception is carbon dioxide which has a very low concentration in the atmosphere but a higher concentration in sea water. This is because carbon dioxide reacts with water to form carbonic acid: CO2(g) + H2O(l) ↔ H2CO3 (aq) ↔ H+

(aq) + HCO3-(aq)

Compare and explain the solubility of selected gases at increasing depths in the oceansGas Concentration at

Increasing DepthReason

Oxygen Low Photosynthetic phytoplankton and wave action dissolving atmospheric oxygen increases concentration near ocean surface.At ocean depths, oxygen is consumed by respiration and decomposition but is not replenished by photosynthesis as light can not penetrate beyond 200 m.

Carbon Dioxide

High Photosynthetic plants consume carbon dioxide near the ocean surface.At ocean depths, no photosynthesis occurs but carbon dioxide continues to be produced through respiration and decomposition.

Predict the effect of low temperatures at great depths on the rate of corrosion of a metalUse available evidence to predict the rate of corrosion of a metal wreck at great depths in the oceans and give reasons for the prediction madeAt ocean depths, the concentration of oxygen is low and temperatures are as low as 4°C ∴scientists predict that corrosion of wrecks would be very slow or inhibited. This hypothesis comes from the knowledge that oxygen is essential for corrosion and that corrosion is accelerated at higher temperatures.

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Gas % dissolved in the atmosphere

% dissolved in sea water

Nitrogen 78 48Oxygen 21 36Carbon Dioxide 0.036 15

Perform a first-hand investigation to compare and describe the rate of corrosion of materials in different:

Oxygen concentrations Temperatures Salt concentrations

Oxygen ConcentrationsMethod:

1. Add a mild steel nail into a test tube containing tap water. Secure with a rubber stopper.2. Add an identical mild steel nail to a second test tube containing tap water which has been

boiled and allowed to cool back to room temperature. Secure with a rubber stopper.3. Leave both test tubes for 7-14 days.4. Each day, record observations for the amount of red-brown deposit formed on both nails.

TemperaturesMethod:

1. Set up 3 identical test tubes, each half filled with tap water and containing a mild steel nail.2. Place one test tube in the fridge, one in a room, and one in an incubator.3. Leave test tubes for 7-14 days.4. Each day, record observations for the amount of red-brown deposit formed on each nail.

Salt ConcentrationsMethod:

1. Add a mild steel nail into 4 test tubes.2. Half fill each test tube with the following solutions: T1 – deionised water (control); T2 – 0.01

mol L-1 NaCl; T3 – 0.1 mol L-1 NaCl; T4 – 1.0 mol L-1 NaCl.3. Leave test tubes in a room for 7-14 days.4. Each day, record observations for the amount of red-brown deposit formed on each nail.

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9.6.6 Predictions of slow corrosion at great depths were apparently incorrect

Explain that ship wrecks at great depths are corroded by electrochemical reactions and by anaerobic bacteriaThe formation of rusticles and black metal sulfides around ship wrecks is due to anaerobic bacteria, such as sulfate reducing bacteria, which rely on energy obtained from the reduction of sulfate ions to hydrogen sulfide. Cathode: SO4

2-(aq) + 10H+

(aq) + 8e- → H2S(aq) + 4H2O The electrons required for the reduction are supplied by the oxidation of iron.Anode: Fe(s) → Fe2+

(aq) + 2e-

Describe the action of sulfate reducing bacteria around deep wrecksSulfate reducing bacteria obtain energy by reducing sulfate to sulfide: SO4

2-(aq) + 5H2O(l) + 8e- → HS-

(aq) + 9OH-(aq)

These bacteria multiply by feeding on the abundant organic food sources, such as textiles and wood, supplied from ship wrecks.Corrosion by bacterial reduction accounts for the rusticle formations that hang from ship wrecks.

Explain that acidic environments accelerate corrosion in non-passivating metalsThe rate of corrosion of non-passivating metals, such as iron, in deep ocean ship wrecks is accelerated by the creation of an acidic environment.

Fe2+ ions, produced from the oxidation of iron, react with H2S, produced from the reduction of SO4

2-, to form H+ ions → makes sea water acidic: Fe2+(aq) + H2S(aq) → FeS(s) + 2H+

(aq)

The hydrolysis of iron also from H+ ions: Fe2+(aq) + 2H2O(l) → Fe(OH)2(s) + H+

(aq)

The presence of H+ ions increase the rate at which oxygen gas is reduced: O2(g) + 4H+(aq) + 4e- → 2H2O(l)

Perform a first-hand investigation to compare and describe the rate of corrosion of metals in different acidic and neutral solutionsMethod:

1. Collect 4 pieces of iron each of a similar size and shape.2. Collect 4 test tubes each containing a solution of HCl with either pH = 2, 4 and 6 or distilled

water with pH = 7.3. Immerse a piece of iron in each test tube.4. Record immediate observations, after 2 hours, after 1 day and after 5 days. Look for signs of

corrosion, such as metal flaking into solution, bubbles on metal surface, coloured deposits on metal.

5. Repeat steps 1-4 with other metals, such as zinc, copper and brass.Results:

Iron – bubbles in the first few minutes at pH = 2; red-brown deposits in all solutions after 5 days Zinc – corroded fastest at pH = 2; no signs of corrosion in pH = 7 Copper – no signs of corrosion in any solution Brass – few bubbles but otherwise no signs of corrosion

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9.6.7 Salvage, conservation and restoration of objects from wrecks requires careful planning and understanding of the behaviour of chemicals

Explain that artefacts from long submerged wrecks will be saturated with dissolved chlorides and sulfates

Metals have been seriously corroded because of the presence of electrolytes Artefacts are severely encrusted with calcium carbonate deposits or coral due to marine

organisms causing differential aeration, i.e. they consume oxygen → limited supply near artefact → artefact becomes anodic → accelerates corrosion

Porous objects, such as leather and wood, are impregnated with sea water containing dissolved chlorides and sulphates

Describe the processes that occur when a saturated solution evaporates and relate this to the potential damage to drying artefactsRecovered wrecks should not be left to dry in the air or in an oven as water evaporates and leaves solid salts. These salts crystallise in the pores of the objects → the destruction of the cellular structure → distortion, cracking or chemically react with the artefact.

Identify the use of electrolysis as a means of removing saltChlorine accelerates corrosion as it provides an acidic environment as it reacts with water to form HCl.Chlorine is difficult to remove from metal artefacts by leaching in clean water so electrolysis is used to free and remove chlorine ions from insoluble compounds.

Identify the use of electrolysis as a means of cleaning and stabilising iron, copper and lead artefactsElectrolysis with a dilute solution of NaOH is used to remove chlorine from recovered metal objects. NaOH speeds the removal of chlorine as OH- replaces Cl- in insoluble compounds, like Fe(OH)Cl, and also retard further corrosion. The recovered object becomes the cathode and a stainless steel mesh surrounding the object forms the anode.Oxidation: 4OH- (aq) → O2(g) + 2H2O(l) + 4e-

(aq)

Reduction: 2Fe(OH)Cl(s) + 4e-(aq) → 2Fe(s) + 2OH-

(aq) + 2Cl-(aq)

Discuss the range of chemical procedures which can be used to clean, preserve and stabilise artefacts from wrecks and where possible, provide an example of the use of each procedureWhen artefacts are first recovered, they must be kept wet in the sea water they were found in to maintain stability and prevent evaporation and crystallisation of salts.DesalinationArtefacts, like leather shoes, undergo desalination to remove impregnated chlorine and sulfate salts. The object is immersed in clean water for periods ranging from hours to weeks. Salts diffuse out into the water; the water is replaced as the salt concentration reaches equilibrium. The treatment is complete when the diffusion rate is less than 50 ppm.Removing ConcretionsBacteria feed on the wreck/artefact → increases the concentration of carbon dioxide → precipitation of calcium carbonate → crusty deposits/ concretions: Ca2+

(aq) + CO32-

(aq) → CaCO3(s)

There are two ways to remove concretions: Treat the artefact with dilute acid, such as HCl or CH3COOH: CaCO3(s) + 2H+

(aq) → Ca2+(aq) + H2O(l) + CO2(g)

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X-ray the artefact to observe the thickness of the concretion, then break away slabs of concretions using a hammer or a dentist drill for delicate areas.

ProtectionRecovered wooden objects are washed with cold water then conserved using polyethylene glycol. Wooden objects are regularly sprayed with PEG of increasing concentrations over long periods of time until PEG completes replaces sea water. PEG fills in wood cavities → restores wood strength + stops further degradation.

Perform investigations and gather information from secondary sources to compare conservation and restoration techniques applied in two Australian maritime archaeological projects

Vernon Anchors Endeavour CannonConservation

ProcessIron surface blasted with copper slag →garnet polished → coated with zinc epoxy paint to prevent rusting.Timber stocks saturated with zinc napthenate solution to prevent deterioration.

Disinfected using 10% formalin in sea water → coral concretions removed using hammers → electrolytic treatment using 2% NaOH to remove chlorine → washed with distilled water with chromate ions (chromic oxide protective layer forms) → dried for 48 hours → treated with protective wax.

Display Outside; exposed to the elements – rain, sea spray, sun, humidity.

Inside; conditioned environment – controlled temperature, restricted light, fenced (no contact)

Maintenance Inspected for deteriorationRegularly hosed down with fresh water to reduce salt build up as anchors are situated near the sea

Minimal maintenance as display environment is controlled

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