Chemistry HSC Full Notes BEST NOTES

Chemistry HSC Note: This is not the final revision of my notes (I’m constantly revising them as I do papers), and there may be a few areas of error or unclear explanations. However, I’ve gone through it a number of times, and it should be mostly very accurate and comprehensive. If you find anything wrong, it would be nice if you could tell me on [email protected] so I can either discuss it or change it. Good luck for the HSC guys Organic Chemistry – the study of compounds containing carbon THIS IS BACKGROUND INFO Organic chemistry is separate because we can look at all of the included chemical groups in a unifying way, through the bonding properties of carbon. Study of major groups: o Oxygen-containing compounds e.g. alcohols o Hydrocarbons e.g. petroleum o Carbohydrates e.g. sugars o Nitrogen-containing compounds e.g. amino acids  proteins Etc. Hydrocarbons When all bonds are single, they are called alkanes. This is a family of compounds, represented by a general formula CnH2n+2, aka a homologous series. They have similar properties and reactions. There are ‘straight’ chain alkanes. e.g. C

C C

C C

C

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

Branched chain (one is attached to at least 3) C C

C C

C

C

Methane, CH4, ethane, C2H6, Propane, C3H8, and Butane, C4H10 are all alkanes. Physical Properties C1 to C4 are gases at room temp, C5 to C18 are colourless liquids, others are solids. The density of alkanes are significantly less than water (1.00g/mL), are non-conductors of electricity and are insoluble in water. The reason for their insolubility is that C-C bonds are nonpolar, and C-H bonds are only slightly. This slight polarity is cancelled by symmetry in structure. Weak dispersion forces, relatively low boiling/melting points. Boiling/melting points increase as molecular weight increases, due to stronger dispersion forces (more electrons). Volatility decreases as molecular weight increases. Alkenes Chem notes

David Lee BHHS 2007

1

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

(2)Full Structural Formula – shows planar geometry around double bond, and tetrahedral around other carbon atom (3)Intermediate type – infers tetrahedral shape (4)Condensed structural formula – no attempt to show structure, but enough information is provided Isomers are different compounds with the same molecular formula but different structural formula. The double bond can be at different positions in the compound. e.g.

Physical Properties Straight-chain alkenes similar to alkanes. Densities similar to corresponding alkanes, insoluble in water. Alkynes Contain a triple bond between carbons. CnHn-2. As with alkenes, isomers are possible. They are nonpolar, low boiling points and insoluble in water Naming Alkanes, Alkenes and Alkynes o Stem telling length of carbon chain C1 meth- C4 butC7 heptC2 ethC5 pentC8 octC3 prop- C6 hexo Look at the longest possible chain, then pick a prefix o Look for branches, and use a number to denote their position, starting from the closest end e.g. 2,3 – dimethylpentane or 2 - dimethylpentane o If double or triple bonds present, set this as priority (start counting closest to the bond) and first state branches then double/triple bond e.g. 2– methyl – 1 – propene (methyl on second branch and double bond on first) o If compound is cyclic, add a cyclo- before the name of the main branch e.g. 1,2,3 trimethylcyclohexane Saturated and unsaturated compounds Alkenes and Alkynes – unsaturated, possible to attach more hydrogen Alkanes – saturated, max no. of H atoms that skeleton can hold Functional Group The functional group of carbon compounds is the most reactive area of the compound. In alkenes and alkynes, the double/triple bonds are the functional groups. When a hydrogen atom is replaced with a halogen atom, e.g. OH, the halogen becomes the functional group. Molecules with a particular functional group react similarly, regardless of the attached chains. Chem notes

David Lee BHHS 2007

2

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

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

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

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

Ethylene is produced from natural gas or crude oil (mixtures of hydrocarbons, containing mainly alkanes and cycloalkanes and smaller amounts of unsaturated including alkenes), which is called feedstock. The feedstock is refined by fractional distillation to obtain alkenes since alkanes are susceptible to combustion and unreactive (not useful as starting material). Ethylene is the most versatile, but not found in large quantities in feedstock. Produced from other hydrocarbons in ‘cracking’ (a process where hydrocarbons of higher mol mass are converted to lower mol mass via breaking of chemical bonds). There is greater demand for some fractions than others (e.g. gasoline > heavier hydrocarbons), and fractions from crude oil are not in optimum ratios, hence cracking. Note that air needs to be excluded to prevent combustion. Ethylene is simple and can be synthesised from many different hydrocarbons. Three ways: 1. Thermal cracking – requires very high temps and generally not used. End products hard to control since many places where bonds could break, early method. Accelerates reaction and drives equilibrium to reactants. 2. Catalytic cracking of fractions separated from petroleum. – material is passed over a catalyst at a temperature of about 500oC, and the particles adsorb onto the catalyst and have their bonds weakened, resulting in decomposition. E.g. C10H22(g) -> C8H18(g) + C2H4(g). Alkane splits further into smaller alkenes until propene/ethylene formed. Catalysts allow it to be carried out at lower temperatures. Zeolite (by mid 1970’s) is the main catalyst, and is a crystalline substance of Al, Si and O. Usually fine powder (higher surface area for action of catalyst) circulated through feedstock. Zeolite gives greater control over products under different conditions of temperature and pressure (thus increasing yields of desired products) i.e. C18H38(g) ----(zeolite catalyst)---> 4 C2H4(g) + C10H22(g) 3. Steam cracking of ethane and propane – ethane from natural gas deposits fed into furnaces with steam, heated between 750 – 900oC causing much ethane to be converted to ethylene i.e. C2H6(g) -> C2H4(g) + H2(g) Propane can also be used: C3H8(g) -> C2H4(g) + CH4(g) Chem notes

David Lee BHHS 2007

3

•

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

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

Dilute H2So4 cat w/ water Br2, non aqueous solvent

KMnO4, H+ Acidified

HCl, non aqueous solvent

Reaction of alkenes Characteristic reaction of alkenes is addition reaction. Two new atoms or groups of atoms are added across double bond, one to each carbon. The C=C is converted to a single bond and a saturated hydrocarbon is produced. General eqn: H2C=CH2 + X-Y => XH2C-CH2Y 1) Addition

of hydrogen to ethylene (hydrogenation) - ethylene to ethane by heating with hydrogen in presence of nickel, platinum or palladium 2) Dibromoethane - Used as a petrol additive - halogen reactions are useful for distinguishing between saturated and unsaturated hydrocarbons. E.g. A non aqueous solution of bromine (e.g. solvent carbon tetrachloride) when added to an alkene causes the solution to lose its colour as bromine becomes incorporated into the alkene: [CH2=CH2(g) + Br2(l)  CH2Br-CH2Br(l)] (petrol additive)

Chem notes

David Lee BHHS 2007

4

Alkanes do not react with NA bromine unless exposed to UV. In aqueous solutions, the reaction may be the same as above, but due to the presence of water products can include: [CH2=CH2(g) + HOBr(aq)  CH2OH-CH2Br] Hydrogen bromide reaction: [CH2=CH2(g) + HBr(g) -> CH3 – CH2Br] (What states?) 3) Chloroethene – Monomer for PVC CuCl +150o C

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

4) Styrene –

produces from benzene and ethylene via the intermediate ethylbenzene

5) Ethanol – Used as a fuel in automobiles and as an industrial solvent ( dilute ) H SO

2 4 → CH3-CH2OH(l) ] [CH2=CH2(g) + H2O(l)    

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

Ag + 250 C [C2H4(g) + ½ O2(g)   → C2H4O(g)] +

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

•

Identify the following as commercially significant monomers o Vinyl chloride o Styrene By both their systematic and common names

Vinyl chloride – chloroethene CH2 = CHCl Monomer for the production of PVC plastics which are widely used in applications such as electrical insulation, plumbing and garden hoses with various additives to change physical properties Styrene – ethylbenzene C6H5CH=CH2 (also known as phenylethene) Production of polystyrene, most stiffened of common plastics due to large phenyl side group. Stable due to presence of C-C and C-H bonds only, minimal chain branching means it can be formed into clear objects. Tool handles, car battery cases, CD cases. Gas can be bubbled through to create foam (foam drink cups), making it soft and light. •

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

Prac – Reactions of hydrocarbons with bromine water Risk analysis: Hazard Risk Chem notes

David Lee BHHS 2007

Control 5

Bromine water

Corrosive, and toxic, can cause skin burns

Cyclohexane

Highly flammable Eye and skin irritant with severe redness and pain Highly flammable, fire hazard Eye and skin irritant with severe redness and pain

Toluene

Wear safety goggles Use small amounts to minimise vapour Wear safety glasses Keep away from hot surfaces, flames or sparks Keep away from hot surfaces, flames or sparks Polyvinyl gloves

Aim: To compare reactivities of an alkene (cyclohexene), alkane (cyclohexane), and an aromatic hydrocarbon (toluene) in bromine water Method: 1). Four semi-micro test tubes were half-filled with bromine water, cyclohexane, cyclohexene and toluene respectively, using eye droppers 2). Bromine water was mixed with the other substances by placing a few drops of bromine water in each micro-test tube with a dropper 3). The test tubes were tapped, and observations recorded Results: Cyclohexane – none Cyclohexene – forms clear solution Toluene – none The functional group reacting with bromine is the double bond present in alkenes, this decolourises bromine water. Addition reactions. These reactions are addition reactions: Bromine w/ water

Br2( aq ) + H 2 O(l ) ←→ HOBr − ( aq ) + H + ( aq ) + Br − ( aq ) Bromine w/ cyclohexene (top)

Br2( aq ) + C 6 H 10 (l ) ←→ C 6 H 10 Br2 ( aq ) Bromine water w/ cyclohexene (bot)

HOBr( aq ) + C 6 H 10 (l )  → C 6 H 10 BrOH ( aq )

Toluene did not react as aromatic molecules have delocalised electrons which do something??? ************ Bromine with cyclohexane, this is substitution: UV

Br2 ( aq ) + C6 H 12 (l ) →  C 6 H 11 Br( aq ) + HBr(l ) Requires UV light to break off hydrogen atom and allow reaction Chem notes

David Lee BHHS 2007

6

Prac – Reaction of lycopene with bromine water Aim: To determine the effect of bromine water in varying amounts on the spectrum of colours reflected by lycopene Method: 1). Five semi-micro test tubes were filled halfway with tomato juice 2). An eye dropper was used to place 1, 2, 3, 4 and 5 drops of bromine in each of the test tubes respectively 3). The solutions were stirred with the stirring rod until colour appeared 4). The colours and the corresponding amounts of bromine water in each test tube were recorded Results: Test tube no. 1 2 3 4 5

No. of drops of Bromine 10 8 6 4 2

Colour Blue Turquoise Green Khaki Orange

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

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

Polymerisation is the process of bonding monomers together to form long chains. Polymers are macromolecules consisting of small repeating units called monomers joined by covalent chemical bonds. Polymers can be divided into two categories: 1. Natural polymers – naturally occurring polymers used by humans since ancient times (E.g. cellulose, silk, rubber) 2. Synthetic – more recent man-made polymers. Replacing natural since they do not corrode, are lightweight and relatively cheap. Celluloid was first commercially manufactured plastic, but highly flammable nature meant it was replaced Ethylene serves as a monomer due to the reactivity of its double bond. It has a structure that can change to accommodate the additional bond needed to join repeating units together. •

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

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

Chem notes

David Lee BHHS 2007

7

•

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

General outline: Ethylene can be changed from gas to liquid under high pressure. This liquid ethylene can be heated in the presence of a catalyst to form polyethylene. Two forms of polyethylene can be produced, each with differing methods and varying properties: o LDPE (reaction conditions 100 – 300oC, 1500 – 3000 atm) – Polymerisation consists of three stages  Initiation – organic peroxide catalyst. They produce free radicals (molecules with unpaired electron), such as H-O. which is a hydroxy radical. This causes the double bond in ethylene to break and form a bond with the radical. CH2=CH2 +R●  RCH2-CH2●  Propagation - The resulting molecule contains an unpaired electron. Bonds to another ethylene molecule through the same process etc. (Chain propagation reactions). Backbiting, where the chain curls onto itself and the free electrons takes a hydrogen atom from an existing CH2 group, causes branching. Termination – at various times, it is possible for two free radical polymers to react to form a covalent bond, ending propagation (chain terminating reaction). o HDPE (50 – 75oC, 350 in industry): H 2 SO4 / H 3 PO4 CH 3 − CH 2 − O − H    → CH 2 = CH 2 + H 2 O

Reverse reaction is hydration, requires dilute aqueous sulfuric acid: H SO dilute

CH 2 = CH 2 + H 2 O 2 4 → CH 3 − CH 2 − O − H They are general, apply to any alkanol or alkene e.g. 1-pentanol to 1-pentene •

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

Risk Analysis Hazard Iodine

Oxalic acid

Risk Toxic – fatal if swallowed. Corrosive, causes burns and damaging to lungs if inhaled Poisonous if swallowed, inhaled or absorbed through skin

Control Safety glasses, effective ventilation Gloves, avoid generation of dust

Prac – Ethanol as a solvent Aim: To test the solubility of various materials in ethanol Method: 1). 20 mL of ethanol was poured into each of 10 test tubes using measuring cylinders 2). A rice-grain amount of each solid was placed into successive test tubes, and a few millilitres of each liquid placed into the remaining test tubes using an eye dropper. 3). Each test tube was agitated by tapping and gently shaking 4). Observations were recorded Results: Solute Sodium chloride Napthalene Cyclohexanol Glycerol Iodine Oxalic acid Boric acid Glucose Wax Urea

Chem notes

Solubility No Slightly Yes Yes Yes, dark red Yes, purple Yes Yes No Yes

David Lee BHHS 2007

13

Ethanol has a single hydroxyl group which is attached to an aliphatic Allows it to dissolve substances with polar covalent bonds, hydrogen bonds form Able to dissolve hydrocarbons and non-polar due to the formation of dispersion forces with the aliphatic group. Widely used as alternative solvent in dissolving medicines, cosmetics, food flavourings, alcoholic beverages, low toxicity so relatively safe

o o o

o

Ionic (sodium chloride): Unable to dissolve since strong ionic bonds holding atoms together, and intermolecular formed inadequately strong to break apart lattice Polar covalent bonds (cyclohexanol C6H11OH, glycerol C3H5(OH)3, oxalic acid C2O2(OH)2, Boric acid B(OH)3, Glucose C6H12O6, Urea CO(NH)2) : Polar covalent bonds such as those in hydroxyl groups allowed ethanol to bond strongly and dissolve them Macromolecules (Wax C24H50) : Though only held together with weak dispersion forces, its large size means a larger surface area of contact between molecules and thus more total dispersion forces. Ethanol was unable to dissolve Non-polar molecules (Iodine, Napthalene, heptane, pentane) – iodine is diatomic and has very weak dispersion forces holding together, but ethanol can form dispersion forces with iodine molecules and pull away. Napthalene is an aromatic hydrocarbon and does not attract strongly, but dissolves in a similar fashion, same for heptane and pentane which are short-chain hydrocarbons. With 1,2,3 – propanetriol: • Outline the use of ethanol as a fuel and explain why it can be called a renewable resource Ethanol is a flammable liquid, burning with the reaction:

C 2 H 5 OH (l ) + 3O2 ( g ) → 2CO2 ( g ) + 3H 2 O( g ) It is also easily transportable, and was used by hikers and campers. It has thus been proposed as an alternative fuel source, having already been used as an ‘extender’ in world war 2. The purpose of ethanol is to: 1. Reduce greenhouse gas emissions 2. Reduce reliability on non-renewable fossil fuels Engines would not need any modification to run 10-20% ethanol fuel, and is renewable since synthesised in sugar cane from carbon dioxide, water and sunlight. Burning produces carbon dioxide and water which can then be re-used to produce ethanol, so it follows an almost indefinite material cycle. It has thus been promoted for motor cars to supplement and replace petrol. • • •

Describe conditions under which the fermentation of sugars is promoted Summarise the chemistry of the fermentation process Present information from secondary sources by writing a balanced equation for the fermentation of glucose to ethanol o Fermentation requires a carbohydrate as starting material, such as glucose, sucrose or starch

Chem notes

David Lee BHHS 2007

14

o Disaccharides such as sucrose and polysaccharides such as starch need to first be broken down into monosaccharides such as glucose/fructose by enzymes in the mixture o This carbohydrate is placed in the presence of yeasts, which produce enzymes that break it down to ethanol and carbon dioxide 

yeast C6 H 12 O 6( aq )  → 2CH 3 − CH 2 − OH ( aq ) + 2CO2( g )

The optimum conditions are 37oC and anaerobic conditions o Reaction is exothermic, so temperature needs to carefully controlled o Fermentation can occur until 15% ethanol, then the yeasts cannot survive and fermentation stops – extra ethanol can be added or distillation used to increase concentration o

•

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

Prac – Fermentation of glucose Aim: To investigate the fermentation of glucose Method: 1). One gram of beef extract and 25 grams of glucose and 7 grams of yeast were measured out using an electronic beam balance 2). A test tube with barium hydroxide was weighed on the triple beam balance and its weight recorded 3). The beaker was filled with 300mL of tap water and heated over a bunsen burner until 40oC 4). The beef extract, warm water, yeast and glucose were quickly poured into the conical flask and weighed on the electronic beam balance 5). The apparatus was set up as shown below: 6). After a week, the tubes and stopper were removed, then the fermented solution and test tube with barium hydroxide were weighed on the triple beam balance and analysed 7). Experiment was repeated 2 times 7). Steps 1-6 were repeated without the yeast, to act as a control Result 1 2 3 Average

Results: Loss in mass of conical flask and contents (g) 8.0 9.8 9.1 9.0

Gain in mass of test tube contents (g) 5.9 4.8 5.0 5.2

Loss in mass of conical flask and contents (g) 0 0 0 0

Gain in mass of test tube contents (g) 0 0 0 0

Control: Result 1 2 3 Average Chem notes

David Lee BHHS 2007

15

The formation of carbon dioxide was evidenced by the formation of barium carbonate in the test tube. The control showed that the loss in mass, and thus creation of carbon dioxide, was caused by the yeast fermenting glucose. Equation: yeast C6 H 12 O 6( aq )  → 2CH 3 − CH 2 − OH ( aq ) + 2CO2( g )

Reaction of Barium hydroxide with carbon dioxide:

Ba (OH ) 2 ( s ) +CO2( g ) → BaCO3( s ) + H 2 O(l ) Nine grams of carbon dioxide was given off in this experiment (loss in mass of conical flask) the gain in the test tube was not used since some carbon dioxide escaped through holes around the stopper.

moles of CO 2 9.0 g (14.01) + (16.00) * 2 = 0.20moles =

Through stoichiometry, 1 mole of glucose produced 2 moles of CO2 and ethanol each. Therefore 0.20 moles of ethanol produced Mass of ethanol = molar mass ethanol x moles produced = 10g (3.3 % w/v) This reaction did not go to completion, 6 grams of glucose left. This is most likely because the yeast were saturated in ethanol and could no undergo further fermentation, or were no left for sufficient time. Some discrepancy could have been caused by measurement error. •

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

Industrial production of ethanol from sugar cane 1. Feed preparation o Crushing – sugar cane is crushed to remove high-quality sugars and molasses, which are used in fermentation o Saccharification – bagasse, the other constituents of sugar cane (50% cellulose) undergo a multi-step hydrolysis process, using an enzyme and sulfuric acid as catalyst to produce glucose compounds 3. Fermentation – yeast and anaerobic conditions ferment glucose compounds to a certain concentration of ethanol (max 15%). (Note: Fermentation less efficient than hydration) 4. Purification of mixture – waste products are removed and distillation is used to concentrate ethanol 5. Addition of gasoline – varying amounts of gasoline are added to produce a commercial product, ‘gasohol’ or E10 is 10% ethanol 90% gasoline •

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

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

David Lee BHHS 2007

16

Ethanol is flammable liquid that is suitable as a fuel. It can be fermented from biomass such as sugarcane and corn, requires land for agriculture and infrastructure for a fermentation industry to be set up. If these requirements are met adequately, the advantages of ethanol as an alternative fuel are: o Renewable – products of ethanol combustion can theoretically be completely recycled to produce more ethanol o Intrinsic anti-knock properties – circumvents the need for toxic anti-knock agents due to presence of oxygen, increases octane of fuel o Burns more cleanly due to presence of oxygen, , reducing toxic emissions (such as hydrocarbons) o Reduction of net emission of greenhouse gases due to reuse of carbon dioxide by biomass used to produce ethanol o Lower blends ( Cu(s) Overall: Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s) The salt bridge allows migration of ions into each beaker (cations into cathode solution and vice versa) to neutralise charge buildup and maintain cell voltage. A larger volume of electrolyte solution means a large surface area of electrolyte in contact with electrodes. This increases the rate at which charged particles are removed from the electrodes, decreasing the internal resistance and thus the current. However, the voltage within the cell is caused by the intrinsic properties of the electrodes (difference in electronegativity), and thus is not Chem notes

David Lee BHHS 2007

22

altered significantly. It is affected slightly since a charge buildup generates a slight back emf, which reduces the voltage, and more electrolyte action reduces this. Note: Concentration of electrolyte does affect voltage. The current output needs to be measured quickly since the ions in salt bridge are being used up, and charge begins to build in half-cells opposing current flow. Sources of error: Electrodes not completely polished – reduces effective surface area of action for electrolyte and thus current Method B: 1). The copper half-cell in A was set up 2). Another half-cell consisting of a magnesium strip in magnesium sulfate solution was set up 3). The electrodes were connected to the terminals of a voltmeter, and voltage reading recorded 4). Steps 2-3 were repeated with: a. Aluminium in 1M aluminium nitrate solution b. Tin strip in tin(II) nitrate solution Results: Test Half-cell Polarity (relative to Total cell voltage Cu/Cu2+) Zn/Zn2+ -ve 0.50 Mg/Mg2+ -ve 1.1 2+ Sn/Sn -ve 0.35 Al/Al3+ -ve 0.90 Standard potential Cu – +0.34 V (Oxidation potential -0.34V) Half Reaction (write in Predicted voltage E0 (V) Experimental E0 (V) during exam) Zn -0.76 -0.16 Mg -2.36 -0.76 Al -1.68 -0.56 Sn -0.14 -0.01 The more active the metal, the greater the potential difference. Note: Oxidation potential is the ability of a substance to oxidise in relation to hydrogen, reductional potential is ability to reduce in relation to hydrogen. E.g. Copper has reduction 0.34 meaning it has higher ability to reduce, but its oxidation is -0.34, it has less ability to oxidise. • Gather and present information on the structure and chemistry of a dry cell or lead-acid cell and evaluate it in comparison to one of the following: o Button cell (Silver Oxide cell) In terms of: o Chemistry o Cost and practicality o Impact on society o Environmental impact Criteria

Chem notes

Dry Cell (Leclanche Cell)

David Lee BHHS 2007

Button Cell (Silver Oxide Cell)

23

Structure

Chemistry

Cost and Practicality

Impact on Society

Environmental impact

•

Oxidation Zn(s) -> Zn2+(aq) + 2eReduction 2MnO2(s) + 2H+(aq) + 2e- -> Mn2O3(s) + H2O(l) Hydrogen ions provided by ammonium NH4+ NH3(aq) + H+(aq) Overall Zn(s) + 2H+(aq) + 2MnO2(s) -> Zn2+(aq) + Mn2O3(s) + H2O(l) Adv: o Inexpensive o Robust o Easy to store and use Dis: o Short life o Voltage not as constant as silver button (comparison) o Cannot deliver very high currents o Cannot be recharged o First commercial battery, made portable electric devices possible o Used widely in toys, torches, radios etc. o Manganese readily oxidised to stable manganese (IV) dioxide, becomes immobilised and not dangerous o Small quantities of ammonium salts and zinc not harmful o Not rechargeable, large volume in landfills so space is an issue

Oxidation Amalgamated zinc Zn(s) + 2OH-(aq) -> ZnO(s) +H2O(l) + 2eReduction Silver oxide Ag2O + H2O +2e- -> 2 Ag(s) + 2OHElectrolyte KOH Overall Zn(s) + Ag2O -> ZnO(s) + 2Ag(s) Adv: o Compact o Provides constant voltage over long period of time, since solid reactants and products have fixed concentration o Overall long operating life due to solid components Dis: o Silver is expensive o Not rechargeable o

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

o

Contains traces of mercury, causes problems with unsafe disposal Non-rechargeable so takes up space

o

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

Eocell = Eocathode - Eoanode Chem notes

David Lee BHHS 2007

24

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

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

o

o o

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

Stable isotopes No radiation emission Z ≤ 83 Ratio of neutrons to protons within zone of stability

Unstable isotopes Emission of radiation Z > 83 Ratio of neutrons to protons out of zone of stability

Zone of stability: A nucleus is unstable when its ratio of neutrons to protons is outside zone of stability. For light elements (Z < 20), 1.0. Stability ratio steadily increases as atomic number increases, up to 1.5 for Z = 83. Past this, all are unstable due to large size of nucleus. Mode α decay β decay Positron emission Electron capture γ emission

•

Emission

Atomic mass

Atomic number

4 2 He 0 −1 e 0 1e

-4

-2

0

+1

0

-1

0 −1 e (absorption)

0

-1

0 0γ

0

0

Describe how transuranic elements are produced

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

Some isotopes undergo fission when bombarded, others undergo nuclear reactions to form new elements When non-fissionable atoms such as Uranium 238 are bombarded with high speed particles, it absorbs the particle to become an unstable atom It then rapidly decays to form a new element

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

David Lee BHHS 2007

25

Linear accelerator – accelerates positive particles in straight line along axes of series of positive and negative cylinders, accelerating it. Often more than a kilometre in length Cyclotrons – accelerates positive particles by passing them through alternating positive and negative electric fields. A strong magnetic field is used to constrain particles to spiral path, reducing size of machine.

o o

They can also be produced in nuclear reactors, a source of neutrons. Neutrons do not experience electric repulsion like positive nuclei, and thus speeds in nuclear reactors are adequate. These create transuranic elements with a proton deficiency. e.g. 238 1 → 239 92 U + 0 n  92 U

 →−10 e+ 239 93 Np

Neptunium, first discovered transuranic element obtained in by chemical separation of nuclear fission reactor products. This can be further bombarded to create plutonium: 239 2 0 → 241 → 241 93 Np + 0 n  93 Np  94 Pu + −1 e

Unstable 238 12 →250 92 U + 6 C  98 Cf

( )

1  → 246 98 Cf + 4 0 n

23 transuranic elements have been created thus far. •

Process information from secondary sources to describe recent discoveries of elements

Recently discovered elements include: o Ununoctium (118, October 10 2006) – heaviest element discovered to date. It was indirectly detected by a team of researchers working in Russia at Dubna University’s Joint Institute for Nuclear Research when they detected its decay products after bombarding californium-249 atoms with calcium-48 ions. Very unstable, half-life 0.89ms. Reaction:

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

o

243 48 288 →115 Uup +301 n 95 Am + 20 Ca 

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

Identify instruments and processes that can be used to detect radiation

o

o

Chem notes

Photographic film – photographic film darkens in the presence of radiation. Used in radiation badges worn by laboratory workers handling radioactive substances to determine radiation dosage Geiger-Muller tube – Radiation enters window and ionises gas particles inside Geiger tube (inert gas such as argon), and the resulting charged particles are accelerated to the two plates with a potential difference. They further ionise other argon atoms through David Lee BHHS 2007

26

collision creating a cascade effect. This creates a signal which is amplified and converted into an audio signal.

o

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

o

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

•

Describe how commercial radioisotopes produced

Radioisotopes can be produced by bombardment of high-speed particles. Radioisotopes are commercially produced in: o Nuclear reactors (proton deficient, neutron enriched) – convenient source of electrons. Target nuclei are placed in reactor core and are then bombarded by neutrons to produce desired isotope. These are then separated chemically or physically from other substances within reactor. Currently operating in Australia for this purpose is HIFAR reactor, managed by ANSTO e.g. Creation of technetium: 98 1 99 m → 42 Mo  →9943 Tc + −10β 42 Mo + 0 n 

Technetium decays, releasing gamma ray inside body o

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

National Medical cyclotron e.g. Creation of gallium-67 Chem notes

David Lee BHHS 2007

27

68 1 67 →31 Mo + 201n 30 Zn + +1 p 

• •

Identify one use of a named radioisotope: o In industry o In medicine Describe the way in which the above radioisotopes are used and explain their use in terms of their chemical properties

Technicium-99m (nuclear reactor) Most widely used in medicine for diagnosis, such as locating brain tumours or studying other parts of the body by being attached to red blood cells. o Short half-life of 6 hours means patient exposure is minimised o Versatile chemistry and can be incorporated into range of biomolecules targeting different organs Iodine-131 – Produced in cyclotron or nuclear reactor, testing of thyroid function and treatment of thyroid ailments such as overactive thyroid or thyroid cancer (beta decay destroys some thyroid cells) o Iodine-131 is naturally absorbed by cells in the thyroid gland o Relatively short half life to minimise exposure (8 days) Specific problems: o Ionising radiation of iodine-131 deals collateral damage to other cells o Radiation penetrates the body and can damage organic tissue near to the patient o Transport and production in nuclear reactors requires stringent safeguards, which is problematic Cobalt-60 – used to measure thickness of materials. With fixed geometry for source and detector, penetration of radiation emitted from radioisotope (beta particles in this case) determines thickness of material. Gamma ray producer o Long half-life so does not need to be replaced frequently (5.3 years) o Low energy emissions, so absorption is significant and can be detected o Low energy emissions, minimises safety procedures required Sodium-24 Used to detect leaks in water pipes or underground oil pipes. Dissolved into water source, can be subsequently detected in soil around areas of leakage. o Dissolves easily into water o Relatively short half-life to minimise environmental damage (15 hours) •

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

Benefits in medicine: o Created wide range of non-invasive diagnostic procedures otherwise impossible, such as technetium-99m used to identify brain tumours, gallium-67 for cancers, an area very dangerous for surgery o Allowed radiation therapy to treat many forms of cancer, e.g. iodine-131 for treatment of thyroid cancer Benefits in industry: Chem notes

David Lee BHHS 2007

28

o o

Ability to make more sensitive, precise and reliable monitoring equipment e.g. cobalt-60 for measuring thickness of materials Allows otherwise difficult activities such as detecting leaks in extensive water distribution systems, sodium-24 can be dissolved and radiation detected near leaks

Problems: o Exposure of radiation doses to workers in medicine, industry and research can damage tissues e.g. ionizing radiation of sodium-24 causes cancer by removing electrons from the biological molecule DNA. o Extra safety precautions are required for sites with radioactive materials, such as proper storage facilities and protective clothing e.g. industries dealing with cobalt-60 and technetium-99m need to filter out the fine radioactive dust produced, which can pose a lung cancer risk o Disposal of radioactive waste requires space, and can be problematic since isotopes such as Cobalt-60 remain radioactive a long time after they are no longer useful, may leak into environment without strict procedures

The Acidic Environment Definitions and properties of acids/bases An acid is a substance that produces hydronium ions (H3O+) in solution A base is a substance containing oxide or hydroxide ions (O2-, OH-) or which in solution produces the hydroxide ions. A soluble base is an alkali, i.e. one which dissolves or reacts in aqueous solution to produce hydroxide ions. Note that oxygen ion-containers are insoluble or only react. Common properties of acids: 1. sour taste 2. sting or burn the skin 3. conduct electricity in solution 4. turns blue litmus red 5. React with reactive metals to form salt + hydrogen gas 6. Reacts with carbonates to form CO2, salt and water 7. Reacts with metal oxides/hydroxides to form salt and water Common properties of alkalis: 1. have a soapy feel 2. have a bitter taste 3. conduct electricity in solution 4. turns red litmus blue 5. React with amphoteric metals to produce hydrogen gas Acids and bases react to form salt and water (there are exceptions). Note that the salt is in aqueous solution, separated as ions and not precipitated. E.g. reaction of sodium hydroxide with hydrochloric acid: HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l) Can also be written as: H+(aq) + NaOH(aq) - > Na+(aq) + H2O(l) Or H+ + Cl- + Na+ + OH- -> H2O(l) + Na+ + ClThe chloride and sodium ions are spectator ions. The net ionic equation is simply H+ + OH- -> H2O(l) A salt is an ionic compound formed when a base (alkali) reacts with an acid. Chem notes

David Lee BHHS 2007

29

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

Classify some common substances as acidic, basic or neutral

Acidic: Vinegar (acetic acid) – used in cooking (~3) Lemon juice (citric acid) (~2.5) Vitamin C (ascorbic acid) – vitamin supplement Hydrochloric acid – pH maintenance in swimming pools, clean bricks cement and tiles (~1) Neutral: Water Salts (e.g. sodium chloride, copper sulfate) Milk Basic: Baking soda (sodium bicarbonate) (NaHCO3) (~8.5) Oven and drain cleaners (sodium hydroxide) – sodium hydroxide also used in soap, and alumina (~ 13) Lime (calcium hydroxide) – making mortar (~ 11) Ammonia – used to make fertilisers (~ 12) • •

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

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

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

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

Chem notes

David Lee BHHS 2007

30

pH in soil testing – steps: • small amount of moist soil placed in a well on a ceramic test plate • White barium sulfate (neutral pH, insoluble) sprinkled on so colour change more easily discerned • 2/3 drops universal indicator added to barium sulfate • Colour change of indicator compared to colour chart Lime (CaO) or dolomite (CaCO3/MgCO3) added if too acidic CaO + H2O  Ca(OH)2  Ca2+ + 2OHCO32- + H2O  HCO3- + OHSulfur is added if too alkaline S + O2  SO2 SO2 + H2O  H2SO3  H+ + HSO3o

Testing home swimming pools which need to be neutral. Acidic water burns eyes, alkaline water causes skin rashes. Operation of electrochemical cell to produce chlorine makes pool more alkaline. A pool sample is collected into a vial and an indicator, usually phenol red (yellow -> red) (6.6 – 8.0) is added and compared to a colour chart. HCl added if too basic. Monitoring wastes from laboratories that process photographic film, as photographic solutions are highly alkaline and discharges need to be neutral in order to not adversely affect the environment pH in aquariums – fish excrete ammonia which reacts with water, making it more basic. Universal indicator is used

o

o o •

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

Prac – Red Cabbage indicator Aim: to investigate the colour changes of an indicator extracted from red cabbage Method: 1). A handful of shredded red cabbage was boiled in water over a bunsen burner for about 5 minutes 2). The resulting liquid solution was poured into a filter paper/filter funnel apparatus and collected in a beaker 3). 20 mL of 1 M HCl solution was poured into a measuring cylinder 4). 10 mL was poured into a small test tube which was labelled ‘0’ 5). The remaining 10 mL was poured into a beaker and diluted to 100 mL 6). 20 mL of the resulting solution was poured back into the measuring cylinder, and steps 4 – 5 were repeated 5 times, each test tube being labelled a successive integer higher 7). Steps 3-6 were repeated starting with 1 M NaOH solution and labelling from 14 down 8). A few drops of red cabbage indicator were added to each test tube and observations recorded Results: Red (0-1) Pink (2)  Purple (3-4)  Clear (5-6)  Purple (7-9) Blue (10-11)  Green (1213)  Yellow (14) There are 6 discrete colour stages for this indicator. This suggests multiple molecules within the red cabbage indicator solution acting to produce colour changes. Each molecule has a specific colour when a proton is added or taken away. The molecules in this case are anthocyanins, and there are about 15 different ones in red cabbage indicator. •

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

Chem notes

David Lee BHHS 2007

31

See book •

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

An acidic oxide is one which either reacts with water to form an acid, or reacts with bases to form salts (or both). E.g. carbon dioxide and diphosphorous pentoxide P2O5 CO2(g) + H2O(l)  H2CO3(aq)  2H+ (aq) + CO32-(aq) (carbonic acid) CO2(aq) + 2NaOH(aq)  H2O(l) + 2Na+(aq) + CO32-(aq) (sodium carbonate) Or alternatively: H2CO3(aq) + 2NaOH(aq)  2H2O(l) + 2Na+(aq) + CO32-(aq) The latter is more correct, as the acidic oxide would react with water to form the acid first. It would depend on the relative concentrations of the oxide and the acid in solution, as it is an equilibrium reaction. However, since both create the same products, this is negligible. And similarly for P2O5 A basic oxide show basic character, and react with acids to form salts, but not with alkali solutions e.g. CuO + H2SO4(aq)  CuSO4(aq) + H2O(l) CuO(s) + H2O(l)  Cu2+(aq) + 2OH-(aq) Amphoteric oxides are those showing both acidic and basic character, and those that react with neither acids or bases are neutral oxides e.g. NO, CO, N2O •

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

Acidic nature of oxides increases from left to right. •

Define Le Chatelier’s principle

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

Identify factors which can affect the equilibrium in a reversible reaction

Reversible reactions occur when products can react to generate reactants. When a reaction starts, forward reaction generates products from reactants. Backward reaction then generates products, Chem notes

David Lee BHHS 2007

32

which form at an increasing rate as product concentration increases. The equilibrium occurs at the point where formation of products is equal to the rate of reactant formation, no net change in concentration. Factors: Concentration species – increasing/decreasing concentration of a species will cause reaction equilibrium to shift so that it decreases/increases the species concentration. This because it naturally results in more/less collisions or more/less decomposition to form more/less of that chemical. Note that reactions involving solids and liquids experience little effect, as concentrations remain almost unchanged (note: this does not include dissolved substances). Pressure in a gaseous reaction – an increase/decrease will cause a increase/decrease in concentration (and vice versa for volume). Depending on which side of the reaction has more particles, the equilibrium will shift in that direction in order to reduce number of particles and thus pressure (or vice versa). Note that increasing reaction by increasing concentration of gas not involved in reaction e.g. argon has no effect Temperature – If the temperature is lowered, the amount of energy in the system decreases and the exothermic reaction is favoured since less particles have sufficient energy to form products with a higher potential energy. And vice versa Catalysts – increases speed at which equilibrium is reached, does not alter equilibrium position as activation energy of both product and reactants formation is decreased

o

o

o

o

Notable exceptions: o When solid or liquid is involved in reaction – the concentration of these substances stays constant o The addition of water to an aqueous reaction involving water – concentration of water does not change significantly, but other substances more dilute •

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

Prac – Degassing soft drinks Method: 1). A 300mL bottle of soft drink was taken and weighed on an electronic beam balance 2). The drink was shaken vigorously, and the top slowly and carefully unscrewed to prevent spillage as the gas escaped 3). The bottle was reweighed 4). 6 grams of table salt was taken and added slowly to the contents of the bottle until no more fizzing was observed 5). Bottle was reweighed 6). A new bottle was heated over a bunsen burner with cap on, then the cap unscrewed slowly to prevent spillage before reweighing it 7).Hydrochloric acid was added until the water stopped fizzing, then it was reweighed Discussion: The dissolution and reaction of CO2 in water is multistep equilibria – CO2(g)  CO2(aq) … (1) CO2(aq) + H2O(l)  H2CO3(aq) … (2) H2CO3(aq) + H2O(l)  HCO3-(aq) + H3O+(aq) … (3) HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq) … (4) Factors: Chem notes

David Lee BHHS 2007

33

Temperature – solubility decreases as temperature increases, opposite to liquids and solids. Increased average kinetic energy of CO2 molecules means they have greater overall tendency to escape from solution. Equilibrium shifts to left for all equations until new equilibrium is reached Pressure – solubility increases with increased pressure, more carbon dioxide dissolves to decrease pressure and act against change, equilibrium shifts to right. Dissolution of ions – dissolution of ions displaces carbonic acid ions and CO2 molecules from hydration shells and causes equilibrium to shift to left and increase CO2 gas concentration pH of water – increased pH means more hydroxide ions, which react with carbonic acid to neutralise it and produce water, resulting in more CO2 dissolved to produce acid to counteract change. If pH lowered, increased acidity means increased concentration of H3O+ ions, shifting equilibrium of (3) and (4) to left to decrease its concentration. This means increased concentration of the reactants on left, which has a cascade effect shifting all equilibrium to left and increasing CO2 gas concentration.

o

o o

o

• •

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen

Natural sources of sulfur dioxide: o Geothermal hot springs and volcanoes, release is unpredictable and changes with volcanic activity. Natural levels of sulfur dioxide vary widely, but account for about ¼ of worldwide emissions o Bacterial decomposition of organic matter, produces H2S which oxidises to form sulfur dioxide 2H2S(g) + 3O2(g)  2SO2(g) + 2H2O(g) Industrial sources are mainly: o Burning of fossil fuels – coal generally contains 0.5% to 6% sulfur as metallic sulfides in sulfur in carbon-containing compounds, released as sulfur dioxide during combustion in power stations. Some sulfur remains in refined petrol which is released as sulfur dioxide in automobiles. e.g.iron sulfide in coal 4FeS(s) + 7O2(g)  2Fe2O3(s) + 4SO2(g) o Processing of fossil fuels - removal of sulfur from crude oil and natural gas releases some sulfur dioxide. o

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

Natural sources of nitrogen oxide/dioxide: o Lightning – high temperatures causes atmospheric nitrogen and oxygen to combine forming nitrous oxide: O2(g) + N2(g)  2NO(g) This slowly reacts with oxygen to form: 2NO(g) + O2(g)  2NO2 o Denitrifying bacteria – converts nitrates in soil into nitrous oxide (N2O), increased use of fertiliser has increased emissions Industrial sources: Chem notes

David Lee BHHS 2007

34

Combustion – includes power stations and automobiles etc. High temperatures involved causes atmospheric oxygen and nitrogen to react, and is released into atmosphere(equations same as above).

o

•

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

Nitrogen dioxide and sulfur dioxide are washed out by rain, so there is no significant buildup in atmosphere. Nitrous oxide however, has steadily increased by about 15%, from measurements made over the last century. There are problems associated with collecting evidence for sulfur and nitrogen oxides, namely: o Concentrations of both are very low, below 0.1ppm, and only recently (since about the 1950’s) are instruments accurate enough to reliably measure the levels, so trends before this period could be invalid o Sulfur dioxide and nitrogen dioxide form sulfate and nitrate ions which are changed chemically as they move around the hydrosphere, so measuring traces of these compounds is difficult Most evidence comes from observed occurrences such as acid rain. There appears to be an increase from data but it is inconclusive due to lack of long-term trends and inaccuracies of earlier measurements. •

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

Concern about release into environment because it has detrimental effect on environment and can cause harm to people since: Sulfur dioxide: o Sulfur dioxide irritates the respiratory tract, and can cause symptoms in people with asthma or emphysema in concentrations as low as 1ppm o Forms acid rain o Dry deposition causing environmental damage Nitrogen oxides: o Nitrogen dioxide irritates the respiratory tract, and can cause extensive tissue damage in concentrations 3-5 ppm o The action of sunlight on nitrogen dioxide, hydrocarbons and oxygen increases photochemical smog, which includes ozone – poisonous substance o NO and NO2 participate in ozone layer depletion (NO + O3  NO2 + O2) o Forms acid rain o Contributes to global warming •

Explain the formation and effects of acid rain

o

Sulfur dioxide and nitrogen dioxide gases released dissolve in water to form sulfuric acid and nitric acid which is washed out of the atmosphere by rain, forming wet deposition acid rain. Reaction with hydroxyl radicals: SO2(g) + 2OH  H2SO4(aq) (OR) 2SO2(g) + O2(g)  2 SO3(g) SO3(g) + H2O(l)  H2SO4(aq)

Chem notes

David Lee BHHS 2007

35

-2NO2(g) + H2O(l) -> HNO2(aq) + HNO3(aq) 2HNO2(aq) + O2(g) –(catalysed by impurities) 2HNO3(aq) (OR) 4NO2(g) + 2H2O(l) + O2(g)  4HNO3(aq) Dry formation: Incorporated into dust and smoke and falls to the ground Effects due to low pH include: • Corrosion and tarnishing of metal and bridges, soiling and surface erosion of marble and stone structures CaCO3(s) + 2H+(aq) → Ca2+(aq) + CO2(g) + H2O(l)

o

• • • •

• • •

Crown dieback in trees Leeching of leaf nutrients Killing of leaf tissue Leeching of Ca2+ and Mg2+ ions from soil as they are mobilised due to decreased pH, reducing soil fertility Inhibits microbial activity Increased acidity of lakes, killing aquatic life e.g. snails can only tolerate up to pH 6.0 Mobilisation of Al3+ ions in soil due to reduced pH. This flows into lakes and precipitates out, clogging fish gills and suffocating them

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

•

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

Gather and process information to write the ionic equations to represent the ionisation of acids Define acids as proton donors and describe the ionisation of acids in water

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

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

o

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

o

Citric acid (C6H8O7) – occurs in citrus fruit, also widely used as a food additive for flavour or as a preservative Hydrochloric acid (HCl) – produced by stomach lining glands to break down food molecules, also made commercially to clean metals, brickwork, neutralising bases etc.

o

Chem notes

David Lee BHHS 2007

36

Sulfuric acid (H2SO4) – synthetic acid manufactured to make fertilisers, synthetic fibres etc. Phosphoric acid (H3PO4) weak Polyprotic acids have more than one ionisable hydrogen per formula unit. For others see book o

• •

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

pH scale is a scale of measurement for hydrogen ion concentration. pH = –log10[H3O+] This obeys the significant figure rule Water self ionises: H2O + H2O  H3O+ + OHKw = [H3O+][OH-] = 1.00 x 10-14 at 298K • •

Describe acids and their solutions with the appropriate terms weak, strong, concentrated and dilute Describe the difference between a strong and weak acid in terms of an equilibrium between intact molecules and its ions

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

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

Prac – Relative strength of acids Aim: To compare the relative strengths of different acids using a variety of methods Method: 1). 50mL of 0.1M hydrochloric, acetic, oxalic, citric and sulfuric acid were prepared in labelled 250 mL beakers 2). A pH meter was used in each beaker and the reading recorded 3). A pH strip was placed into each beaker for a short period and its colour compared with a chart 4). A few drops of universal indicator were dropped into each beaker and the colour compared to a colour chart Results: ##### The strongest was sulfuric acid. It is a strong acid and is diprotic, meaning that the concentration of H3O+ ions is twice the concentration of hydrochloric acid. HCl is strong but monoprotic, meaning concentration is identical to HCl concentration. Oxalic acid is diprotic and citric is triprotic, Chem notes

David Lee BHHS 2007

37

but both are weak and do not completely ionise. The tendency for their conjugate bases to re-bond with hydrogen ions limits the concentration of H3O+ in solution. Acetic is the weakest, being monoprotic and have a low degree of ionisation. Oxalic – C2H2O4 Citric – C6H8O7 Acetic/ethanoic – CH3COOH Safety: HCl - corrosive, vapour can burn mouth, throat and eyes Oxalic acid – corrosive to tissue, corrosive to respiratory tract if inhaled Wear safety glasses, goggles, use lower concentrations and smaller amounts •

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

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

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

Use available evidence to model the molecule nature of acids and simulate the ionisation of strong and weak acids Gather and process information to explain the use of acids as food additives

Acids are added to food to: o Improve taste – e.g. carbonic acid in soft drinks, acetic acid in vinegar o Preserve food – increases acidity to point where bacteria can no longer survive e.g. coating freshly cut fruit with citric acid o Increase nutritional value – e.g. adding ascorbic acid (Vitamin C, an antioxidant) • Identify examples of naturally occurring acids and bases and their chemical composition Acids: o Ascorbic acid C6H8O6 – occurs widely in fruit and vegetables, essential to health o Citric acid – found in citrus fruits o Lactic acid CH3CH(OH)CO2H – produced by anaerobic respiration in cells, found in muscle tissue and milk o HCl – stomach to break down food Bases: o Chem notes

Ammonia NH3 – result of decomposition of proteins, or anaerobic decay of organic matter found in fish urine David Lee BHHS 2007

38

Carbonates – e.g. calcium carbonate (limestone), magnesium carbonate Metallic oxides – e.g. Iron (III) oxide, copper oxide and titanium oxide, found in minerals. Metals are extracted from them.

o o •

Calculate pH of strong acids given appropriate hydrogen ion concentrations

•

Outline the historical development of ideas about acids including those of: o Lavoisier o Davy o Arrhenius Lavoisier (1780) • acids were substances that contained oxygen • Disproved since some oxygen-containing compounds such as metallic oxides were basic, and distinctly acidic substances such as hydrochloric acid contained no oxygen • Wrong but stimulated research Davy (1815) • suggested that acids were substances that contained replaceable hydrogen. • Bases were substances that reacted with acids to form salt and water. • These definitions worked well for most of that century, but the definition made no attempt to interpret the properties, only classify the substances Arrhenius (1884) • Interpreted acidic properties in terms of ionisation to form H+, and weak/strong in terms of degrees of ionisation • the conductivity of acid solutions and their reaction with many metals to form hydrogen gas evidenced that acidic solutions contained hydrogen ions • acids were substances that ionised in solution to produce hydrogen ions • A base is a substance that in solution produced hydroxide ions • He defined strong acids as those that ionised completely and weak as those that partially ionised. General equations: HA(aq)  H+(aq) + A-(aq) XOH(aq)  X+(aq) + OH-(aq) Weaknesses were:  Does not take into account role of solvent in ionisation of acid (ionisation results from reaction of acid with solvent)  Acid-base reactions can occur in solvents where there is no ionisation  Not all acidic/basic substances (e.g. metallic oxides) ionised to produce hydrogen/hydroxide ions

o

o

o

•

Outline the Bronsted-Lowry theory of acids and bases

An acid is a proton donor, a base is a proton acceptor. Gives the broadest definition of acid/base theory (it means that acids must have hydrogen). This definition : o Does not restrict bases to those which ionise to produce hydroxide ions, such as in the case of metal oxides and ammonia o Explains how neutralisation reactions don’t require dissolution of ions into aqueous solution e.g. NH3 + HCl in benzene (direct proton transfer) o Exchange of proton relies on relative properties of both substances involved, accounting for the role of the solvent o Shows that hydrolysis of salts to change pH were acid or base reactions o Provided basis for quantitative treatment of acid-base equilibria and pH calculations Chem notes

David Lee BHHS 2007

39

•

Trace developments in understanding and describing acid/base reactions

•

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

?

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

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

A salt is an ionic compound containing a cation not H+ and an anion not O2- or OH-. In aqueous solution, salts completely dissociate into ions. Type Acidic Basic Neutral Salt Ammonium nitrate Sodium acetate Sodium chloride (NH4NO3) (NaCH3COO) Potassium nitrate Sodium hydrogen Potassium nitrite (KNO3) sulfate (NaHSO4) (KNO2) Sodium sulfate Anything Sodium carbonate (Na2SO4) 3+ containing Al , (Na2CO3) Fe3+, HSO4- or Anything H2PO4 containing F-, S2etc. The pH of a salt solution depends on the nature of its ions, many cations/anions serve as acids or bases. Some generalisations: Neutral salts have anions which are the conjugate base of strong acids, and cations the conjugate acid of strong bases, since their reaction to accept/give protons is negligible Basic anions react with water to form hydroxide ions in solution. Reaction is equilibrium, occurs to small extent since conjugate acid is stronger than water, and conjugate base is stronger than basic anion Acid anions contain hydrogen atoms to react with water to form hydronium ions, derived from polyprotic acids. The anion resulting from hydrolysis of polyprotic acids is amphiprotic, and whether it is acidic or basic depends on the tendency for one hydrolysis reaction (proton donation or proton accept) to occur over the other

o

o o

Some examples: (Basic anions) S2–(aq) + H2O(l) ↔ HS–(aq) + OH–(aq) F–(aq) + H2O(l) ↔ HF(aq) + OH–(aq) (Acidic cations) [Fe(H2O)6]3+(aq) + H2O(l) ↔ [Fe(OH)(H2O)5]2+(aq) + H3O+(aq) •

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

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

David Lee BHHS 2007

40

Risk analysis: Hazard Ammonium chloride

•

Risk Released vapour causes coughing, shortness of breath

Identify conjugate acid/base pairs

Acid HCl H2SO4 HNO3 NH4+

Base

Conjugate base ClHSO4-,(and SO42-?) NO3NH3

OHCNCO32CH3COO•

Control Wear safety goggles Use small amounts to minimise vapour

Conjugate acid

H2O HCN HCO3CH3OOH

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

Note: Hydrolysis is when a substance reacts with water Amphiprotic substances can act as both a proton donor and proton acceptor. They react to both accept protons and donate protons. Their behaviour changes whether in aqueous solution or alkaline/acid solution. e.g. HCO3- (hydrogen carbonate) In aqueous solution: HCO3-(aq) + H2O(l)  H3O+(aq) + CO32-(aq) HCO3-(aq) + H2O(l)  H2CO3(aq) + OH-(aq) In basic/acidic solution: HCO3-(aq or s) + OH-  H2O(l) + CO32-(aq) HCO3-(aq or s) + H+(aq)  H2CO3(aq) Reactions go to completion since products cannot perform the reverse reaction (???). This also applies to HSO3- (hydrogen sulfite) and HSO4-. Water is amphiprotic. •

Identify neutralisation as a proton transfer reaction which is exothermic

Neutralisation reactions are proton transfer reactions, and involve the reaction between an acid and a base. They are exothermic and thus have a negative enthalpy change. The net ionic reaction in Arrhenius theory is: OH-(aq) + H+(aq)  H2O(l) Acids and bases not fitting in Arrhenius theory do not necessarily produce salt and water during neutralisation reactions e.g. neutralisation of ammonia. In LB theory, an acid and base react to form conjugate base and conjugate acid. The acid gives a proton to the base. Reactions between strong acids and bases form very weak conjugate acids/bases and go virtually to completion since the back reaction has almost no tendency to occur. Otherwise, reactions are equilibria. Chem notes

David Lee BHHS 2007

41

•

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

Volumetric analysis is a form of chemical analysis where the concentration of a substance is determined. Determining the composition of a solution require titration against another solution of known concentration, called the standard solution. The substance dissolved is a primary standard. Equipment: A primary standard: o Must be obtainable in very pure form and have known formula o Should not alter weight unintentionally during preparation/titration e.g. absorbing moisture from air o Have a reasonably high formula mass to minimise weighting errors o Purified by drying in oven and cooling in dessicator to eliminate moisture and prevent its absorption e.g. oxalic acid, sodium carbonate Use of equipment: o Pipette – solution to be used is first drawn in above mark, then solution let out until meniscus at mark, solution let out through gravity with tip against wall of container o Burette – first, rinse with portion of solution to be dispensed, overfilled then excess allowed to run out Preparation: o Accurately measure mass of primary standard e.g. electronic beam balance o Rinse a volumetric flask and beaker with distilled water o Pour the primary standard into beaker and dissolve with distilled water, less than intended volume of final solution o Pour into volumetric flask, and repeat a few more times o Use a pipette to add final few drops to complete solution Titration curves: Strong acid strong base: Equivalence point at pH 7, steep curve. Indicator used should have colour endpoint near equivalence point. Using indicator changing during equivalence point is inaccurate, too difficult to tell exact colour shade needed

Strong base/weak acid e.g. NaOH, acetic acid

Chem notes

David Lee BHHS 2007

42

Equivalence point in basic range since salt formed is basic, the anion is a conjugate base of weak acid and thus a weak base. Equilibrium reaction occurs, weak base reacts to reform conjugate acid, thus decreasing acidity since less H3O+

Strong acid/weak base Similar to above with equivalence point acid. Special case when CO2 formed during reaction e.g. HCl + Na2CO3, CO2 forms carbonic acid. Weak/weak Not good since gradient around equivalence point is quite shallow, big volume difference between indicator endpoints and equivalence points, needs to fit indicator very well or use one which changes during equivalence, hard to distinguish. •

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

Prac – Titration Aim: To standardise HCl and NaOH solutions using titration Method: Preparing standard: 1). A clean 250mL beaker was placed on an electronic balance, zeroed, and had 2.650g of Na2CO3 added using a spatula 2). Approx 100 mL of distilled water was added to the beaker, and solution stirred using stirring rod 3). A 250mL volumetric flask was rinsed with distilled water* 4). The Na2CO3 solution was poured into the volumetric flask, and step 2 repeated 5). The stirring rod/beaker were thoroughly washed using a wash bottle, the runoff dripping into the vol. flask 6). Using a 25mL pipette, the volumetric flask was filled to the 250mL mark 7). A stopper was placed on the flask and contents swirled to mix 8). The pipette was rinsed with the unknown HCl* 9). 10 mL of unknown concentration HCl was poured using a pipette into a 50mL beaker washed with distilled water*, and a few drops of methyl orange added 10). A burette was washed and filled with the Na2CO3 standard solution 11). The standard solution was quickly drained into the beaker to find an approximate end point 12). Steps 8-10 were repeated 3 times but more accurately * to ensure no cross-contamination See book for calculations •

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

Buffer solutions resist changes in pH. It contains comparable amounts of a weak acid/base and its conjugate base/acid. Take for example an acetic acid (CH3COOH) and sodium acetate (NaCH3COO) system (acidic buffer). CH3COOH(aq) + H2O(l)  H3O+(aq) + CH3COO- (1) Chem notes

David Lee BHHS 2007

43

Addition of sodium acetate would increase the concentration of CH3COO – ions on the right. The equilibrium shifts to the left, but due to the unchanged concentration of H3O+ ions, it stays enough to the right for dissociation to cause a net increase in CH3COO – ions. This net increase enhances buffering capacity. When hydronium ions are added to solution, the equation will shift to the left according to Le Chatelier’s principle to reduce the concentration of H3O+ ions. When hydroxide ions are added, the CH3COOH will react to form water and CH3COO-, reducing OH- concentration. In both cases, pH change is reduced. Buffer in natural system Carbonic acid / bicarbonate ion buffer system in mammalian blood: H2CO3(aq) + H2O(l)  HCO3-(aq) + H3O+(aq) Maintains the blood at around 7.4 for optimum function, too high/low can result in death. •

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

o Many acids and bases are corrosive, can damage materials if spilt Neutralisation reactions can: o Reduce or nullify corrosive properties of spill, minimising damage o Utilise common, cheap, safely handled/stored materials and produces relatively harmless products e.g. sodium bicarbonate NaHCO3 Sodium bicarbonate is amphiprotic, so it can be used for both acidic and basic spills:

OH − ( aq ) + HCO3 − ( aq )  → CO3 2− ( aq ) + H 2 O(l ) H + ( aq ) + NaHCO3 ( aq )  → CO2( g ) + H 2 O(l ) + Na + ( aq ) Vinegar – commonly used in cooking, contains acetic acid:

OH − ( aq ) + CH 3COOH

( aq )

 → CH 3COO − ( aq ) + H 2 O(l )

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

o

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

Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures

Strong intermolecular forces – higher BP/MP compared to similar mole mass (roughly similar dispersion forces) -OH carboxylic acid group

Forms One hydrogen bond between molecules None ionised, neutral •

Two polar bonds between molecules (CO polar and C – O – H hydrogen bond), higher boil/melt point Small no. ionised, acidic

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

Chem notes

David Lee BHHS 2007

44

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

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

Esters – carboxylic acids combined with alcohols, equilibrium reaction Standard equation:

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

o o

H SO

4 → CH 3COOCH 3 (l ) + H 2 O(l ) e.g. CH 3COOH (l ) + CH 3OH (l ) ←2  (acetic acid + methanol)

H 2 SO4 HCOOH (l ) + CH 3CH 2 OH (l ) ←  → HCOOCH 2 CH 3 (l ) + H 2 O(l )

(formic acid + ethanol) •

Describe the purpose of using acid in esterification for catalysis The acid acts as a dehydration agent, removing water from the reaction (Le chatelier argument), thus increasing yield Acts as a catalyst, lowering activation energy to speed reaction Only small amounts of acid required Most common is sulfuric, others include tosic, scandium (III) triflate

o

o o o •

Explain the need for refluxing during esterification

Reflux – the backflow of reactants into the reaction vessel o Reactants and ester products can be volatile o Reaction is carried out at high temperatures to speed reaction, causing evaporation of alcohol and ester and thus loss of reactants/products o Vapour is also dangerous as it is flammable and toxic o To avoid loss and prevent diffusion, a cooled condenser is placed over the reaction vessel, covering it o The vapour condenses here and runs back into the reaction vessel, which  Increases yield and saves on resources  Allows the reaction to be carried at higher temperatures (faster)  Prevents flammable gases from escaping Chem notes

David Lee BHHS 2007

45

• •

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

Esters occur naturally, and are identified as fragrances and flavours in fruit and flowers e.g. orange (octyl acetate). Animal fats such as butter, or oils such as linseed are also esters, as are waxes. These fats and oils can be used to make soap (Fat or oil) + sodium hydroxide → Salt of carboxylic acid + glycerol The salt of carboxylic acid is the cleaning factor of soap Artificially produced esters: Aspirin: o Acetylsalicylic acid active ingredient, discovered by Hoffman in 1897 o C7 H 6 O3 + C 4 H 6 O3  → C9 H 8O4 + C 2 H 4 O2 Salycilic acid + acetic anhydride Most widely used pain-relieving drug, e.g. headache, and prevents blood clots

o

Ethyl acetate: o Solvent in industry, nail polish remover o Acetic acid and ethanol C2H5OH + CH3COOH → C4H8O2 + H2O Octyl acetate: o As a food flavouring such as in sweets, ice cream etc. o Acetic acid and octanol C8H17OH + CH3COOH → C10H20O2 + H2O •

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

Prac – Esterification 1). 15 mL of acetic acid, 15 mL of butan-2-ol and 10 drops of conc. Sulfuric acid was placed into a 100 mL round-bottom flask 2). The apparatus was set up as shown 3). The bunsen burner was lit and reflux allowed to occur for about 15 minutes until two layers clearly visible 4). A separating funnel was assembled on a retort stand and 100 mL of distilled water poured inside 5). Contents of round-bottom flask were transferred into funnel and the mixture shaken 6). 50mL of sodium carbonate solution was added and gently shaken, occasionally inverted and tap opened to release gas 7). Once the layers separated, the lower layer was discarded 8). Step 6-7 was repeated 9). The remaining liquid was poured into a conical flask and a teaspoon of CaCl2 added, shaken gently then allowed to stand for a few minutes CH3COOH(l) + C4H9OH(l) (H2SO4 conc.) CH3COOCH2CH2CH2CH3(l) + H2O(l) (diagram representation here) H2SO4 (see above) Na2CO3: neutralise remaining acid, excess forms distinct layer so easily separated Chem notes

David Lee BHHS 2007

46

CaCl2: anhydrous, absorbs water in mixture to leave ester Safety: 1). Acetic acid + butan-2-ol flammable, apparatus should be secured to ensure no tipping; tipping could pose fire risk due to contact with bunsen flame 2). Acetic acid corrosive to skin, avoid spilling 3). Condenser should have adequate water to prevent organic vapour escaping, flammable and respiratory irritant

Chemical Monitoring and Management •

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

Industry: Australian chemical manufacturing company Branch: Analytical chemistry – quantitative and qualitative analysis of substances present within materials Roles: o Occasional monitoring of ethylene quality, waste water (pH, suspended solids, hydrocarbons etc.), gaseous emissions (particulates, toxic pollutants) to ensure reliability of results by other chemists o Check proper operation of equipment/calibrate instruments o Train shift workers to use instruments o Look for ways of improving processes Principle: Adsorption in gas-solid chromatography o Gas chromatography – a liquid or gaseous mixture is vaporised into a stream of helium flowing over a stationary phase such as a solid o If stationary phase is solid, the components of injected mixture adsorb (stick onto) its surfaces to differing extents, and desorb at different rates o This causes different substances to pass through at different rates o A detector is able to quantitatively measure each substances as it passes out o Can be used to determine chemical composition of substances •

Gather, process and present information from secondary sources about the work of practising scientists identifying: o The variety of chemical occupations o A specific chemical occupation for a more detailed study

Many areas a chemist can work in, 13 divisions recognised by Royal Australian Chemical institute including: o Environmental chemistry (detailed) – determining how substances interact in the environment, monitoring concentrations of substances especially in air, water and soil  Environment monitoring, employed by Environmental Protection Authority, mining companies, local government – qualification can be BSc and postgrad qualification in fields such as scientific communication/management. Could collect data on air/water quality, then analyse and assess this information. Require strength in chemical analysis, and instrumental analysis. May work in a team, providing environmental advice to external bodies via reports. o Physical chemistry – study and measurements of physical aspects of compounds and reactions e.g. reaction rates, structure of substances, nature of chemical bonding Chem notes

David Lee BHHS 2007

47

Pharmaceutical chemistry – discovery, testing, synthesis and commercial development of chemicals for use as medicines Industrial chemistry – chemistry of industrial processes such as manufacture of ammonia, sulfuric/nitric acids and others

o o

•

Identify the need for collaboration between chemists as they collect and analyse data Chemistry is diverse, chemists specialise in particular branches as range of knowledge too large Some real-world problems require expertise from more than one branch Collaboration is required for proper tackling of problem Also, work of one chemist may have implications in another area, requires active communication skills This facilitates efficiency of scientific progress and scientific work Can be achieved by:  Publishing of papers  Collaboration between laboratories  Direct voice communication

o o o o o o

•

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

e.g. Petrol combustion in motor vehicles (mostly octane) Different situations/products: o Complete combustion – excess oxygen, products carbon dioxide and water

2C8 H 18(l ) + 25O2( g )  →16CO2( g ) + 18 H 2 O(l ) Incomplete combustion – insufficient oxygen, products carbon monoxide, carbon dioxide, soot, unburnt hydrocarbons, water. Ensure adequate oxygen supply to fuel Nitrogen oxides – reaction of oxygen with atmospheric nitrogen due to high temps forms nitric oxide and nitrogen dioxide (see acidic environment). Rhodium-platinum catalyst converts more polluting gases into less harmful ones

o o

2 NO( g ) + 2CO( g )  → N 2( g ) + 2CO2 ( g ) 2CO( g ) + O2( g )  → 2CO2( g ) Sulfur oxides – some sulfur compounds in fuels:

o

S ( 2 ) + O2 ( g )  → SO2 ( g ) 2 S ( s ) + 3O2( g )  → SO3( g ) Monitoring can ensure minimum possible toxic chemicals released, important since: o CO affects judgement/perception as levels as low as 10 ppm, can cause death by asphyxiation o Soot contributes to particulate pollution, bad for asthma sufferers o Nitric oxide affects respiratory systems and is generally toxic, excessive production by motor vehicles can affect health of population o Sulfur oxides and nitric oxides contribute to acid rain o Motor industry can use information to build more efficient engines (more complete combustion) •

Identify and describe the industrial uses of ammonia

Chem notes

David Lee BHHS 2007

48

Fertilisers – reacted with sulfuric acid to form ammonium sulfate or nitric acid to form ammonium nitrate fertiliser. Application to soil provides good source of nitrates essential for crop growth, improving yields. 2NH3 + H2SO4  (NH4)2SO4(s) Conversion to nitric acid – nitric acid is used in making explosives, dyes, fibres and plastics Neutralisation of acid – petroleum industry uses ammonia to neutralise acid components of crude oil and protect equipment from corrosion Water treatment – addition of ammonia and chlorine to water produces more stable disinfecting residual than chlorine alone

o

o o o

• • • • • • •

Identify that ammonia can be synthesised from its component gases, nitrogen and hydrogen Describe that synthesis of ammonia occurs as a reversible reaction that will reach equilibrium Identify the reaction of hydrogen with nitrogen as exothermic Explain why the rate of reaction is increased by higher temperatures Explain why the yield of product in the Haber process is reduced at higher temperatures using Le Chatelier’s principle Analyse the impact of increased pressure on the system involved in the Haber process Explain that the use of a catalyst will lower the reaction temperature required and identify the catalyst(s) used in the Haber process

N 2( g ) + 3H 2 ( g ) ←→ 2 NH 3( g )

∆H = −92 Kj / mol

Nitrogen and hydrogen combine to form ammonia, which in turn decomposes to reform reactants (reversible reaction). Equilibrium reached when rate of forward/reverse reactions the same. At higher temperatures, the average kinetic energy of the reactants is higher. Thus: 1. Larger fraction of molecules have adequate energy to overcome activation energy and react upon collision 2. Molecules move faster, more collisions between molecules These increase the rate of reaction, and apply to both forward and backward reactions. However, it affects forward reaction more. Reaction is exothermic, higher temperatures shifts equilibrium to left to reduce temperature change. Number of moles of gas on each side of reaction is different. According to Le Chatelier’s principle, increasing pressure shifts equilibrium to right since there are less moles of gas, decreasing pressure and thus minimising pressure change. Vice versa Catalysts: iron-iron catalyst, small amounts of K2O, Al2O3 o Hydrogen and nitrogen molecules adsorbed onto surface, increasing collisions o Allows reaction via new chemical pathways with lower Ea o Reduces activation energy, allowing molecules with lower kinetic energy to react, thus lowering the required temperature o Reaction is faster, equilibrium unaffected •

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

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

David Lee BHHS 2007

49

Higher temperatures means faster reaction rate but lower yield (equilibrium more to reactants), and higher energy costs associated with temperature maintenance The reverse is true for lower temperatures, their advantages conflict, optimal temperature currently used is 700K Higher pressures increases yield but places more stress on reaction vessel, current optimum 2.5 x 104 kPa Catalysts can speed up reaction rate and lower reaction energy, lowering temperature required for same production rate

o o o

o •

Explain why monitoring of the reaction vessel used in the Haber process is crucial and discuss the monitoring required Temperature and pressure of reaction vessel – keep within range for optimum production rate, excess temperatures can damage catalyst and lower yield Ratio of incoming reactants – maintain stoichiometric ratio and prevent buildup of one reactant, slowing reaction Impurities in incoming gases – O2 can cause explosion, CO/CO2 can poison catalyst and reduce its lifespan Rate of ammonia removal – inadequate rate of removal shifts equilibrium to reactants, reducing yield

o o o o

•

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

o o o o o

Haber developed process in 1908, before WW1 Growing population in early 20th century required large amounts of fertiliser to feed population Growing militancy of Germany required more ammonia for explosives Haber process able to meet these demands Germany originally obtained nitrates as saltpere/guano from Chile, but advent of WW I caused allies to set up a naval blockade, preventing imports, but Haber process allowed Germany to be self-reliant This prolonged their resistance against the Allies, increasing the length of World War 1 and resulting in loss of many more lives Significant impact in world history

•

Deduce ions present in a sample from a range of tests (Maybe need mixtures – go to conquering chem for good summary) Cations: Cations: Confirmation tests: Ba2+ - apple green flame Ca2+ - brick red flame Cu2+ - blue-green flame, dissolves in ammonia to form deep blue solution Fe2+ - decolourises acidified dilute KMnO4 solution Fe3+ - deep red solution with SCNChem notes

David Lee BHHS 2007

50

* Fe2+ can form white or green which may tun brown due to oxidisation over time (older samples) * Copper forms CuI with iodine, not used in initial procedure Note: Silver chloride is also a white precipitate, but silver sulfate is insoluble and colourless Anions: HNO3 Ba(NO3)2 Pb(NO3)2 AgNO3 Other 1 Carbonate Bubbles White White pH between 8 CO32precipitate precipitate, and 11 soluble in soluble in HNO3 HNO3 * Sulfate (acidified) (acidified) SO42thick white *white precipitate precipitate Phosphate (ammonia)* Yellow (acidified) 3PO4 white precipitate (NH4)2MoO4,2 precipitate soluble in yellow soluble in HNO3 precipitate HNO3 (may need warming) Chloride Cl(acidified)  dissolves in white ammonia, precipitate darkens in sunlight 1

CO32-(aq) + 2H+(aq)  CO2(g) + H2O(l) Sulfate weaker lewis base than phosphate SO42-(aq) + H3O+  HSO4-(aq) + H2O(l), equilibrium enough to left so adequate sulfate to produce noticeable precipitation with barium/lead (CHECK – how would this effect lead?) In basic conditions, enough phosphate for precipitation 2 12(NH4)3MoO4 + PO43- + 3H+  (NH4)3PMo12O40 *

•

Perform first-hand investigations to identify the following ions: o Phosphate o Sulfate o Carbonate o Chloride o Barium o Calcium o Lead o Copper o Iron

Some safety information: o Barium compounds are toxic (see barium chloride below) (barium sulfate is mostly harmless) o Silver chloride – eye, skin and respiratory irritant o Lead nitrate – poisonous if swallowed, causing spasms, nausea etc. Can be absorbed through skin to cause irritation, redness over short periods •

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

Chem notes

David Lee BHHS 2007

51

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

Effects on People

Corrosion of copper plumbing

Over consumption – Liver, kidney damage (liver is storage point for copper) (Chuttani et al., 1965) Over consumption Vomiting, diarrhea, nausea (USEPA, 1980) Contact with skin – eczema, edema of eyelids (Patty, 1963) Death (NRC, 1977)

Copper sulfate crystals to control algal growth

• •

Effects on aquatic ecosystems Kills plants as well as algae (bioaccumulation) Reduces survivability of aquatic invertebrates (WSDOE, 1992) Reduces survivability of fish (Holland et al, 1960)

Perform first hand investigations to measure the sulfate content of lawn fertiliser and explain the chemistry involved Analyse information to evaluate the results of the above investigation and to propose solutions to problems encountered in the procedure

Hazard Risk Hydrochloric acid (if used) Corrosive. Can cause permanent eye damage Kills tissue Barium Chloride Eye, skin and respiratory irritant Hot plate (if used) Can cause burns upon contact

Control Wear safety glasses, use lower concentrations Wear safety glasses, use lower concentrations Gloves, keep body parts away

Prac – Sulfate content of fertiliser 1). 1.0g of ammonium sulfate was placed into a measuring cylinder, and 2 ml of water was added; mixture was shaken to dissolve 2). 10 mL of 1.0M barium chloride solution added to form white precipitate 3). This precipitate was allowed to settle overnight, then the clear liquid decanted 4). 10mL of water was added to precipitate to further dissolve chloride and ammonium ions, mixture shaken 5). Steps 3-4 repeated twice 6). Two pieces of filter paper were weighed and then shaped into a cone 7). The barium sulfate slurry was evenly poured into this cone, then the cone placed on a source of heat for an hour to evaporate remaining water 8). Barium sulfate + filter paper was weighed, results recorded and analysed Results Mass of ammonium sulfate: 1.000 Mass of BaSO4 precipitate: 0.655 g Mass of SO42- =

96.06 x0.655 = 0.270 g 233.36

27% sulfate ions (w/w) in fertiliser Chem notes

David Lee BHHS 2007

52

Evaluation The sources of error are: 1). Loss of barium sulfate due to slight solubility – using increased concentrations of reactants would decrease its effect on results, but increases contamination by adsorption. Use between 0.005 and 0.05 mol/L of initial sulfate concentration 2). Adherence of barium sulfate to walls of the measuring cylinder 3). Incomplete drying of precipitate, contains water while weighed – repeated cycles of drying/cooling/weighing until constant mass obtained 4). Remaining chloride and ammonium ions not completely removed by washing – more repetitions, although this would cause greater loss of BaSO4 due to dissolution 5). Contamination by adsorption of substances in solution during precipitation – form precipitate slowly by slowly mixing reactants, and forming in hot solutions maximises particle size and reduces adsorption 6). Small size of barium sulfate crystals – some may have fallen through double filter paper, use of a sintered glass funnel to trap it would result in more accurate results OR Weighed amount of agar added as coagulant OR Form in hot solution to maximise particle size + small amount of HCl • •

Describe the use of AAS in detecting concentrations of metal ions in solutions and assess its impact on scientific understanding of the effects of trace elements Interpret secondary data from AAS measurements and evaluate the effectiveness of this in pollution control

Atomic absorption spectroscopy – used to measure low concentrations of elements in ppm range, mainly metals. Each element has unique emission spectrum, so by measuring, studying and using spectra we can determine qualitatively and quantitatively the elements present in sample (by looking at spectra, measuring intensity) Practical arrangement + workings: o Measures how much light of a specific wavelength is absorbed by sample being studied o Sample to be analysed fed into flame, vaporises and converts molecules and ions into atoms o Atomic emission lamp producing a specific emission spectrum matching element to be studied is passed through the flame o Electrons absorb energy and are excited to higher energy levels o Light is passed through prism to concentrate light of desired wavelength onto detector o Wavelengths characteristically absorbed by element in question shows obvious drops, indicating absorbance o Absorbance linearly proportional to concentration o This value is compared to a function (absorbance vs concentration) produced by measuring known concentrations to find conc. Why useful: o Relies on absorption rather than emission, nearly 100% atoms in ground state absorb as opposed to 80

•

Ions (O2+, NO+), O

Contains most of earth’s gases, organisms inhabit this zone, weather events Contains ozone layer (25 km), temperature increases with altitude and gives stability Coldest layer (down to -100 celsius) Temp rises with altitude, ionic and atomic gas particles, important in radio communications since radio waves reflect off

Identify the main pollutants found in the lower atmosphere and their sources Pollutant CO Airborne lead CFC’s SO2

Chem notes

Source Burning fossil fuels, forest fires Lead smelters, leaded fuels Foaming agent, refrigerant-air conditioner coolant, propellant Combustion (fuel impurities), metal extraction from sulfide ores, chemical manufacturing David Lee BHHS 2007

54

Oxides of nitrogen (NO + NO2) Particulates •

Combustion (vehicles and power stations) Combustion, bush fires, industrial processes such as mining

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

Allotrope – a different physical form of the same element in the same phase Ozone is: o An allotrope of element oxygen o Naturally present in atmosphere; only 0.02 ppm at ground level in clean air, 2 – 8 ppm in stratosphere o Detrimental in lower atmosphere – poisonous to many organisms; causes breathing difficulties, fatigue and headache in humans o Beneficial in upper atmosphere - filters our short wavelength UV light which can damage living tissue Produced in upper atmosphere through UV light: 3O2(g) -- (UV light)  2O3(g) • •

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

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

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

Property Boiling point (oC) Odour Colour Density Reactivity

Solubility in water Oxidation ability

Chem notes

O2 -193

O3 -111

None None Slightly denser than air Highly reactive with many metals and nonmetals

Sharp, irritating Pale blue 1.5 times denser than air Very highly reactive, attacks double bonds on alkenes

Sparingly soluble

More soluble than O2

Lower

Higher

David Lee BHHS 2007

Reason Polar bonds between molecules means intermolecular forces stronger (O3) One more O atom per molecule The O-O bonds in O3 are less strong than the double covalent bond in oxygen O2 is non-polar, O3 is bent and thus is polar O involved only in coordinate covalent has greater electron affinity 55

•

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

Free radical – a neutral species with an unpaired electron which can be formed by splitting a molecule into two neutral fragments Property O2 O Reactivity Less reactive (full outer Very reactive (unpaired electron) valence shell) Oxidation Lower Higher (unpaired electron, high ability tendency to take electrons to complete valence shell) •

Identify the origins of CFCs and halons in the atmosphere CFCs (contain chlorine, fluorine and carbon) – developed as refrigerant in 1930’s to replace ammonia, also used as propellant, solvent and foam blowing agent. Through use, gas released into atmosphere Halons (carbon and halogens) – were used in fire extinguishers, recently use has been drastically reduced

o

o

•

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

•

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

Discuss the problems associated with the use of CFC’s and assess the effectiveness of steps taken to alleviate the problem Write the equations to show the reactions involving CFC’s and ozone to demonstrate the removal of ozone from the atmosphere Identify alternative chemicals used to replace CFC’s and evaluate the effectiveness of their use as a replacement for CFC’s

• •

Problems: Removal of stratospheric ozone o CFC’s released into the troposphere are not washed out by rain (non-soluble) and not destroyed by sunlight/oxygen at low altitudes o Diffuse into the stratosphere and short wavelength UV breaks a chlorine off e.g.

CCl 3 F +UV  →Cl • +CCl 2 F o

Chlorine reacts with ozone

Cl + O3  → •ClO + O2 o

ClO reacts with free oxygen atoms

ClO + O  → •Cl + O2 o

o o

Net result is conversion of O3 and O to two O2 Chlorine molecule unchanged at end, can continue to react and remove ozone This occurs on average a few thousand time before chlorine radical reacts with another chemical which removes it e.g. methane

Cl + CH 4  → HCl + •CH 3 Chem notes

David Lee BHHS 2007

56

The problem: o Removal of stratospheric ozone reduces filtering of short length UV radiation, meaning:  Increased incidence of sunburn and skin cancer  Increased damage e.g. brittleness to synthetic materials such as PVC  Increased risk of eye cataracts  Reduced plant growth for some species (e.g. rice) due to UV interference with photosynthesis mechanisms o CFC’s are greenhouse gases, and enhance global warming Alleviation: o Agreements to phase out use of CFC’s (e.g. Montreal Protocol, cease use in developed by 1996) o Agreements to phase out use of halons by 2010 o Assistance to poorer countries to phase out CFC use o Replacement with safer alternatives  HCFC’s – contain C-H bonds decomposable by radicals and atoms in troposphere and are decomposed to significant extend. However, ozonedestroying capacity is still significant (phase out by 2030 – Montreal Protocol). Only useful as temporary substitute  HFC’s – no C-Cl bonds, do not form Cl atoms in atmosphere, no ozonedestroying capacity. Useful as permanent substitute, but more expensive  Air being used to replace as foaming agent Assessment: o Adherence to agreements will ensure ozone layer returns to pre-CFC state since damage is reversible o The pace of CFC withdrawal means it will be many decades before the above happens, meaning the effects of increased UV radiation will be felt o Relies on co-operation of countries, will be less effective if countries withdraw o Replacements such as HFC’s are more expensive than CFC’s, may be a burden on lower countries until better alternatives found Overall, it is an effective long-term solution but prolongs problems in the short term. It relies heavily on cooperation of countries and this may be a downfall. •

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

Evidence/information CHECK: o Measurements of total ozone in a column of atmosphere have been conducted since 1957 o In springtime of 1980-1984, a severe depletion of ozone above Antarctica was detected by the British Antarctic survey o By 1985 it was approximately 30%, and in some places it had been completely destroyed o A net decrease of 3% per decade was recorded for the period 1978-1991, factoring in natural variations Method of collection: Data collected by range of ground and airborne instruments o Dobson spectrophotometer (up to 48 km, groundbased)  Developed in 1924, only source of long-term data Chem notes

David Lee BHHS 2007

57

Can be used to measure both total column ozone and profile ozone, currently used to calibrate measurements by other methods  Measures the intensity of four different wavelengths of UV radiation reaching it; two are strongly absorbed by ozone, the others are not  The ratio between the two intensities is determined and used to calculate total ozone Disadvantages/advantages  Strong affected by aerosols and pollutants  Measures only over a small area 

•

o

LIDAR (10 – 50 km, groundbased)  Relies on absorption of laser light by ozone  Telescope used to collect UV light scattered by two laser beams, one which is absorbed by ozone (308 nm) and one which isn’t (351 nm)  By comparing these values, a profile of ozone concentration vs altitude can be measured

o

Balloons (up to 40km, airborne)  Various instruments can be mounted onto balloons  Electrochemical concentration cells measure current produced by chemical reactions with ozone  Photospectroscopy utilises film or electric sensors sensitive to UV light to measure wavelengths affected by ozone Disadvantages/advantages  Can provide many days of continuous coverage  Inexpensive  Unpowered, flight path cannot be controlled

o

TOMS (from space)  Observes incoming solar energy and backscattered UV radiation at six different wavelengths  Gas molecules in the atmosphere scatter some EM radiation back, while some is absorbed by ozone  By comparing the intensity of backscattered radiation to incoming radiation, amount of ozone can be obtained Disadvantages/advantages  Provides global coverage  Constant, accurate coverage  Coverage in variety of weather and geophysical conditions  More expensive than other methods Identify that water quality can be determined by considering: o Concentrations of common ions o Total dissolved solids o Hardness o Turbidity o Acidity o Dissolved Oxygen and Biochemical oxygen demand

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

David Lee BHHS 2007

58

o

Concentrations of common ions 1. Cations include Na+, Mg2+, Ca2+, K+, Al3+ and heavy metals such as Hg2+, Pb2+, Cd2+ and arsenic. Common anions are Cl-, SO42- and HCO32. Hardness (e.g. calcium, magnesium), detract from taste and appearance (iron), Choke marine life(Al3+), toxic to humans (Hg2+, Pb2+), promote eutrophication (nitrates, phosphates) which forms undesirable decomposition products e.g. ammonia and detracts from taste/appearance 3. Water percolating through soil or underground estuaries, agricultural runoff 4. a. AAS used to analyse cations b. Gravimetric analysis used for anions e.g. precipitating Cl- as AgCl

o

Total dissolved solids 1. Mass of solids dissolved in unit volume of water (ppm) 2. High TDS reduces crop and plant growth, >500 ppm is not suitable for human consumption, >1000 ppm is unsuitable for irrigation 3. Underground aquifers, flowing through farming/grazing areas with disturbed soil 4. a. Evaporation, but it produces very small masses of solid, easy to lose through turbulent bubbling/spitting (use large volumes of water, 1 litre round-bottom flask is suitable) b. Conductivity – most solid dissolved substances are ionic and conduct electricity; can be measured by conductivity meter to determine dissolved solids

o

Hardness 1. Concentration of Ca2+ and Mg2+ ions in water (ppm CaCO3) 2. Hard water forms a precipitate with soap, reducing cleaning power. Under high temperatures such as in kettles, Ca2+ forms an insoluble precipitate with sulfate and carbonate ions, reducing kettle efficiency 3. (see below) 4. a. Titration with EDTA, which forms stable complexes with these ions. Indicator Eriochrome Black T. In solution is blue but forms red coloured complex with Mg2+. Endpoint when it turns blue, indicating no more Mg2+ in solution

o

Turbidity 1. Measure of suspended solids in water 2. Undesirable appearance and taste, reduces sunlight penetration for plant photosynthesis, can absorb IR and raise water temperature 3. Clay, silt, plankton, industrial wastes 4. Secchi disk, visually

o

Acidity 1. pH of water 2. At extreme ranges, it can reduce survivability for aquatic organisms 3. Decomposition of organic matter, acid rain, exposure of sulfide ores in mining, fertiliser run-off 4. a. Universal indicator solution or paper b. pH meter

o

Dissolved Oxygen (ppm), Biochemical oxygen demand

Chem notes

David Lee BHHS 2007

59

1. Concentration of dissolved oxygen in water / capacity of organic matter to consume oxygen 2. Regular levels of O2 (~10 ppm) indicate high quality water, since low levels indicate high BOD and other factors such as: a. Heat pollution which reduces O2 solubility b. Excess of organic wastes such as sewerage which take up O2 in aerobic bacterial decomposition c. Eutrophication by excessive growth of aquatic plants 3. Dissolved from atmosphere, produced by photosynthesis by aquatic plants, algae / Aerobic organisms such as bacteria, fish, worms 4. a. (dissolved oxygen) Titration (see later) b. (BOD) Addition of nutrient to sample and incubating at 20oC in sealed, air-free container in dark for 5 days, then measuring residual dissolved oxygen. Difference is BOD •

Identify factors that affect the concentrations of a range of ions in solution in natural bodies of water such as rivers and oceans Pathway from rain to water body – rainwater collects ions before it runs into natural bodies of water • Bushland contains small amounts of nitrates and phosphates from natural nutrients on surface • Rainwater soaking into ground collects Ca2+, Mg2+, sulfate and chloride from soil and rocks it flows through • Percolation into deep underground aquifers results in collection of Fe3+, Mn2+ among others Human activity • Removal of natural vegetation or irrigation can increase salinity and thus NaCl in rivers • Agricultural fertilisers contribute nitrates and phosphates through runoff or dumping • Discharge of sewage increases nitrates/phosphates, and various ions such as Cl• Acid rain caused by industry is better able to leech certain cations e.g. Ca2+ and Mg2+ from soil • Motor car emission can increase lead Frequency of rain – more rain means more dissolved ions entering water bodies Bushfires – bushfires unlock nutrients and ions such as nitrates from plants, water picks this during runoff Water temperature – higher water temperature increases evaporation and thus increases concentrations of all ions in solution

o

o

o o o

•

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

Prac – Qualitative analysis 1). A sample of catchment water was taken and visually inspected for colour and hydrocarbons (hydrocarbons produce rainbow effect on surface) 2). Some water was poured into a small conical flask and shaken for one minute, then observed for bubbles which indicate detergent 3). Temperature of the water was recorded with a thermometer 4). A pH meter connected to a data logger was used to measure the pH of the solution 5). A turbidity tube was lowered into the sample until the bottom disappears, and reading recorded 6). Silver nitrate was added, and a white precipitate indicated the presence of chloride ions 7). A fresh sample was taken and Heavy metals (See below) Chem notes

David Lee BHHS 2007

60

8). The above steps were repeated for tap water and distilled water Results: Colour PO43SO42Hydrocarbons Turbidity (NTU) Detergents pH Heavy metals ClTDS

Distilled water Clear No No No