Yr 12 Notes (Other Half)

9.2 – Production of Materials:

1. Fossil fuels provide both energy and raw materials such as ethylene, for the production

of other substances:

Identify the industrial source of ethylene from the cracking of some of the fractions from the

refining of petroleum:

Petroleum (crude oil) is a complex mixture of hydrocarbons consisting mainly of alkanes and

smaller quantities of unsaturated hydrocarbons including alkenes.

It includes gas, petrol, kerosene, fuel and diesel oil, lubricating oils, crude oil vapours and bitumen.

Crude oil is separated into its different components using fractional distillation

Petroleum refining consists of distilling crude oil to divide it into a series of fractions according to

their boiling point ranges. Each fraction contains several different hydrocarbons of similar molecule

mass

Ethylene (C2H4) is one of the most useful substances in the petrochemical industry, and is in

extremely high demand

Cracking is the chemical process of which hydrocarbons of higher molecular mass are converted to

hydrocarbons of lower molecular mass.

It is an endothermic reaction because during the process, chemical bonds within the hydrocarbon

molecules are broken

Reasons for cracking – increase output of high demand products

o In refineries, the output of product does not match the economic demand; ethylene is in very high

demand, but it only makes up a very small percentage of crude oil

o To match the demand for ethylene, low, demand, long-chain hydrocarbons are cracked and

ethylene is produced

Catalytic cracking

o Process carried out in a cat cracker

o Alkanes with 15-25 carbon atoms are broken into two smaller molecules, an alkane and an alkene

o Catalyst used for cracking alkanes are inorganic compounds called zeolites: crystalline

aluminosilicates (compounds of aluminium, silicon and oxygen with some metals ions attached)

o Carried out at 500°C in the absence of air, with pressure just above atmospheric pressure

o By using a catalyst, the cracking process can occur at lower temperatures with considerable

savings in energy. However, it cannot decompose large molecules completely into ethylene, so it is

insufficient in meeting the demands and current needs in industry.

o C15H32 C10H22 + C5H10

o C5H10 C2H4 + C3H6

Thermal cracking

o Also known as steam cracking – non catalytic process, however requires high temperatures

o Process in which a mixture of alkanes with steam is passed through very hot metal tubes ( 700-

1000°C), at pressure above atmospheric, to decompose the alkanes completely into small alkenes

such a ethylene, propane and butene, and some hydrogen.

o Steam is present as an inert diluent to allow easy flow of the gases through the hot tubes, while

keeping the concentrations of the reacting gases low enough to ensure that the desired reactions

occur.

.

Identify that ethylene, because of the high reactivity of its double bond, is readily

transformed into many useful products:

Ethylene has a highly reactive double-bond due to the high electron density of the double bond. One

of the bonds readily break, creating two new bonding sites on the molecule

The reactive bond in ethene means that it can easily be converted into a range of useful products

such as ethanol and the starting materials for several important plastics.

Addition reactions when substances which react with alkenes by opening out the double bond to

form two single bonds

Main advantage is that ethylene can undergo polymerisation

Hydrogenation: Hydrogen is reacted with ethylene, using a platinum catalyst at 150 °C, creating

ethane

o C2H4 (g) + H2 (g) C2H6 (g)

Hydration: Ethylene is reacted with water, using phosphoric as a catalyst, producing ethanol.

Industrially important reaction, solvent and pharmaceutical.

o C2H4 (g) + H2O (l) C2H5OH (l)

Halogenation: Reactive molecules from the halogen group - group 7, such as Fl, Cl and Br, react

with ethylene.

o C2H4 (g) + Cl2 (l) C2H4Cl2 (l) (1,2 – dichloroethane)

Hydrohalogenation: Hydrohalogen or hydrogen halide such as HCL, HFI and ethylene reacts to

form a halo-ethane

o C2H4 (g) + HCl (g) C2H5Cl (g) (chloroethane)

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

Polymerisation is the chemical reaction in which many identical small molecules combine to form

one very large molecule

Monomer – small molecules that can be linked together

Polymer – large product molecule formed from many identical small molecules combine together to

form one large molecule.

Alkanes Alkenes

Substitution reaction Addition reaction

An atom in a molecule is replaced by another atom or group

of atoms

Opening of double bond to form two single bonds

with no loss of any atoms

Often requires UV light Spontaneous

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

Additional polymer – formed by molecules adding together without the loss of any atoms.

Each double bond ‘opens out’ to form single bonds with the neighbouring molecules

Polyethylene is an additional polymer, as the ethylene molecules combine with each other.

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

industrially important polymer:

Initiation: Heated initiator or catalyst splits to form free radicals, and is added to ethylene. the

initiator activates on ethylene molecule by attaching to it, breaking its double bond, and attaches to

only one bonding site, creating an ethylene initiator radical

Propagation : Another ethylene monomer attaches to this radical, opening another bonding site, and

another ethylene molecule attaches it, forming a new activates species. Polymerisation continues by

ethylene molecules one after the other adding to the growing chain

Termination: The reaction stops when two activate chains collide and the two radicals react and join

together forming a stable, longer polymer molecule/chain.

This is a random process so the length of polyethylene chains can vary greatly.

Low Density Polyethylene High Density Polyethylene

Crystalline linear chains

Amorphous (non-crystalline) branched

Chains.

Structure Properties Uses Conditions for production

Low density

Polyethylene

(LDPE)

A lot of chain branching

so that the molecules are

unable to get close to one

another and dispersion

forces between chains are

weakened.

Non- crystalline and

amorphous

Extensive chain-

branching

Lack of stiffening side

groups

Lack of cross-linking

Low density

Low melting

point

Greater

flexibility and

softness

Not strong

Wrapping

materials

Disposable

shopping bags

Flexible toys

Milk bottles

Squeeze bottles

in both the home

and laboratory

High pressure (1000-3000

times atmospheric)

High temperatures

(3000°C)

Initiator (organic peroxide,

a compound containing a

– O–O – group or oxygen

gas

High Density

Polyethylene

Long unbranched chains

so that the chains are able

High density

High melting

Kitchen utensils

and containers

Pressure of only a few

times atmospheric

(HDPE) to intertwine and align

closely. This leads to an

orderly arrangement and

it is crystalline

Crystalline linear

chains

No chain branching

point

Relatively hard

and tough

Plastic crates

Variety of

building products

Rubbish bins

Pipes for

transporting

natural gas to

households.

Temperatures of about

60°C

Catalyst – Ziegler - Natta

(mixture of titanium (III)

chloride and a

trialkylaluminium

compound)

Identify vinyl chloride and styrene as commercially significant monomers by both their

systematic and common names:

Structure of

monomer

Systematic name

(monomer)

Common name

(monomer)

Structure of

polymer

Common name and

(systematic) of polymer

Ethene Ethylene

Polyethene

(polyethene)

Chloroethene Vinyl Chloride

Polyvinyl chloride

(polychloroethene)

(PVC)

Phenylethene

OR

Ethenyl benzene

Styrene

Polystyrene

(Polyethenyl benzene)

OR

(polyphenylethene)

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

properties:

– Low-Density Polyethylene (LDPE):

Uses Related to Properties :

Plastic cling wrap because it is flexible, clear and non-toxic

Cling wrap – low density chains don’t pack closely together so its flexible and transparent,

allowing it to be used to wrap around bowls and sandwiches.

Milk bottles; as it is non-toxic, cheap, un-reactive and recyclable.

– High-Density Polyethylene (HDPE):


Kitchen utensils and containers; as it is strong and non-toxic

Rubbish bins; as it is rigid, only slightly flexible and hard

Pipes and other building materials, as it is rigid, hard, and un-reactive

– Polyvinyl Chloride (PVC):


Garden hoses; it can contain UV inhibitors; it is relatively un-reactive, flexible, and durable.

Can be softened with plasticisers.

Pipes and guttering, it is very rigid and hard, and un-reactive. It is also easily shaped – is a

thermoplastic as it can be repeatedly melted and reshaped.

– Crystal Polystyrene:


CD cases and cassette tapes; used because polystyrene is clear, hard, rigid, easily shaped, and

is a good insulator.

Screw driver handles and kitchen cupboard handles; very durable and strong, hard and

inflexible.

– Expanded Polystyrene:


Packaging, and disposable cups; it is slight (full of air), cheap, and it is a thermal insulator

due to the bubbling of the gas through the mixture

Sound-proofing; it is a shock absorbent material, light, easily shaped.

PRACTICAL – Identify data, plan and perform a first-hand investigation to compare the

reactivities of appropriate alkenes with the corresponding alkanes in bromine water:

– In this experiment an alkene (cyclohexene) and its corresponding alkane (cyclohexane), were placed

in a solution of yellow bromine water.

– METHOD:

1. Place 2 medium test tubes on test tube rack

2. Drop 10 drops of cyclohexane in one test tube and cyclohexene in the other test tube

3. Place test tube rack in fume cupboard and add 10 drops of 0.1mol of bromine water into each test

tube

4. place a stopper one ach test tube

5. Shake test tube vigorously for 30 seconds

6. Observe for any colour changes or observations in the bottom bromine water layer. If the

solution decolourises the bromine water then the solution is an alkene.

– RESULT: It was observed that cyclohexene turned the bromine water colourless (decolourised),

whereas the cyclohexane solution remained yellow.

– Thus ONLY cyclohexene reacted with the bromine water, and thus the alkene was said to be more

reactive than its corresponding alkane; this is due to the double bond of the alkene.

– JUSTIFY the method:

Cyclohexene and cyclohexane were used, instead of ethylene or propene because C1 to C4 are

gases at room temperature, and would be hard to manage; cyclohexene is liquid at room

temperature.

Also cyclohexene/ane was used instead of hexene/ane because cyclic hydrocarbons are more

stable than their linear counterparts.

– SAFETY precautions:

Bromine water is highly toxic if ingested, and is slightly corrosive.

Cyclohexene and cyclohexane are poisonous if ingested, and both give off fumes, as they are

highly volatile and highly flammable. Use small quantities in a fume cupboard.

PRACTICAL – Analyse information from secondary sources such as computer simulations,

molecular model kits or multimedia resources to model the polymerisation process:

– In this experiment, molecular modelling kits were used to show how polyethylene is produced

through the polymerisation of ethylene.

– The class was divided into groups, and each group was provided with a kit.

– 3 ethylene monomers were created by each group, with black balls representing carbons and smaller,

white balls representing hydrogen.

– Then the monomers were ‘polymerised’: each group combined their monomers with every other

group until a large chain was created – a section of polyethylene.


The models created a 3D representation of the chemical process, which led to greater

understanding of polymerisation.

The use of ball-and-stick models, depicting the double-bond with flexible rubber rods, greater

increased understanding of the process.

– LIMITATIONS of the method:

The model only provided a very limited section of a polyethylene molecule, as there were limited

numbers of kits.

The use of catalysts (such as Zeigler-Natta catalysts) was not shown in the process, and thus it

was not completely accurate.

2. Some scientists research the extraction of materials from biomass to reduce our

dependence on fossil fuels:

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

petrochemical industry:

There is an overwhelming need for alternative sources of compounds that are presently derived from

the petrochemical industry

Crude oil is the current source of fuels and petrochemicals including plastics

Crude oil is a fossil fuel, and hence it is a non renewable source and is finite, and there are debates to

how long current supplies will last

Some argue that s oil supplies diminish, costs will increase and oil will become too expensive to use

as a fuel, overall energy will fall and alternative fuels will become cost effective.

One of the most appealing replacements for crude oil is cellulose, because it contains all the carbon-

chain structures needed for the production of materials, and it is so remarkably abundant on Earth.

Explain what is meant by a condensation polymer:

Condensation polymer is a reaction between two different functional groups in which a small

molecule (usually water) is eliminated and the functional groups become linked together. Reactive

groups on both ends of each monomer react with one another

Example: cellulose and starch

Describe the reaction involved when a condensation polymer is formed:

Example: when two glucose monomer molecules react though two hydroxyl groups –OH, an H-OH

molecule is condensed out, leaving an –O– linking the two monomer molecules. The first two

glucose molecules to join condense out an H-OH, and every glucose molecule added to the growing

chain then condenses out another H-OH.

β – glucose

C6H12O6

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

found as a major component of biomass:

Biomass is the material produced by living organisms mainly plant material, but it also includes

animal excreta.

Cellulose is a natural condensation polymer.

It is formed when β – glucose monomers link together by beta – 1,4 – glycosidic bonds to form a

long, un-ranched chain polymer.

The glycosidic bond (C-O-C) forms when two –OH groups condense and water is split out.

It is a long polymer chain made of repeating glucose monomers units which FLIP for every alternate

glucose.

Hydroxyl groups in molecule result in strong hydrogen bonding that holds the cellulose chain

together

Produces a water insoluble polymer with great strength and rigidity

A flat, straight and rigid molecule

In cellulose the –CH2OH groups alternate on opposite sides of the chain forming flat ribbon-like

very linear strands. Due to the presence of many hydroxyl (-OH) groups, strong hydrogen bonds

form between the polymers strands giving cellulose its great strength and rigid structure.

However in the structure, very few –OH groups can interact with water molecules and so cellulose is

insoluble in water.

C2H12O6 (s) + C2H12O6 (s) C12H22O11 (s) + H2O (l)

Identify that cellulose contains the basic carbon-chain structures needed to build

petrochemicals and discuss its potential as a raw material:

Cellulose is a major component of biomass

Biomass needs to be considered as an alternative source of chemical energy to replace non-

renewable fossil fuels such as coal and oil.

The long chain carbon structure of cellulose gives it the potential to replace petroleum as the

feedstock for petrochemicals.

As each glucose unit of cellulose has 6 carbon atoms joined in a chain, this is the basic structure for

starting molecules for petrochemicals.

Because the basic structure used to make petrochemicals is short-chained alkenes like ethylene (2C),

propene (3C) and butene (4C), therefore glucose has the potential to be transformed into those

compounds.

The Potential of Cellulose as a Raw Material :

Although cellulose can provide limitless amounts of renewable raw materials, it is currently too

expensive and impractical

As the conversion of the long chains of cellulose into glucose monomers is achieved by acid

hydrolysis (acidic decomposition) or by enzymes (bacterial digestion). The glucose is then

fermented by yeast to produce ethanol

Ethanol can then be dehydrated using concentrated sulfuric acid to form ethene, which can be used

to manufacture a wide range of petrochemicals.

This is a lengthy and expensive process, hence cellulose has great potential, but it is not

economical.

REPORT – Use available evidence to gather and present data from secondary sources and

analyse progress in the recent development and use of a named biopolymer. This analysis

should name the specific enzyme(s) used or organism used to synthesise the material and an

evaluation of the use or potential use of the polymer produced related to its properties:

– Name of Biopolymer: Polylactic acid (PLA)

– Organism Used: Alcaligenes eutrophus

– Production:

1. Starch is harvested from corn, wheat etc

2. Extraction of sugar from plants e.g. corn to produce corn starch by adding amylase and

glucoamylase to break down starch to sugars.

3. Bacteria (alcaligenes eutrophus) placed in solution of sugar/starch. Bacteria ferments sugar into

lactic acid

4. Lactic acid is extracted from alcaligenes eutrophus

5. Catalyst is added to dehydrate the lactic acid into lactide. (Lactic acid is chemically treated to

cause it to link up into long chains or polymers)

6. Lactide acts as monomer and is converted to PLA

– Properties:

Appearance of a transparent thermoplastic

Highly flexible and can be made to be heat resistant to a temperature of 120°C, so it can be

used to contain boiling water and be microwave safe.

Hard, colourless and transparent

Biodegradable, although needs exposure to UV light to decompose

– Uses in Relation to Properties:

Sutures, stents, loose-fill packing, compost bags, food packaging, disposable tableware since it is

flexible and biodegradable.

Fibrous PLA – nappies and disposable garments since its biodegradable and flexible

– Advantages:

Biodegradable

Renewable resources to produce

– Disadvantages:

Agricultural impact, less crops being produced for food

Land diverted for production of plastics instead of other uses

Significantly more expensive to produce, therefore not very successful at replacing cheap

petrochemical plastics.

3. Other resources, such as ethanol, are readily available from renewable resources such

as plants:

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

process and the catalyst used:

Ethylene is made from ethanol by dehydration

Dehydration is a chemical process in which water is removed from a compound

Ethanol is dehydrated by heating it with concentrated sulfuric acid which acts as a catalyst.

An acid catalyst is needed because the acid breaks the C-OH and C-H bonds. Allowing the

formation of a double-bond and water (also reduces activation energy)

Dehydration of ethanol:

ethanol ethylene + water

C2H5OH (l) C2H4 (g) + H2O (l)

Describe the addition of water to ethylene resulting in the production of ethanol and identify

the need for a catalyst in this process and the catalyst used:

The hydration of ethylene is the chemical process whereby a water molecule is added to ethylene

forming ethanol.

Ethylene is hydrated by heating it with dilute sulfuric acid which acts as a catalyst.

The acid catalyst, opens the double bond, allowing water to attach, forming ethanol.

Hydration of ethylene:

ethylene + water ethanol

C2H4 (g) + H2O (l) C2H5OH (l)

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

substances:

A range of substances including polar, non-polar and some ionic compounds dissolve readily in

ethanol

It dissolves polar substances The ethanol molecule consists of two parts:

o The polar hydroxy end (-OH)

o The non-polar alkyl end (-CH2CH3)

The polar hydroxy end is able to form dipole-dipole forces or hydrogen bonds with other polar

substances and hence ethanol is able to form hydrogen bonds with polar water molecules and so

dissolve water. This is because its OH end is very polar; the C-O and O-H bonds are polar, because

O is much more electronegative than C or H

The alkyl end of the ethanol can form dispersion forces with other non-polar substances and hence it

is able to dissolve the non-polar hexane.

Outline the use of ethanol as a fuel and explain why it can be called a renewable resource:

Combustion is the reaction whereby a hydrocarbon reacts with oxygen to form carbon dioxide and

water vapour, releasing energy.

Ethanol is able to undergo combustion, so it can be used as a fuel

C2H5OH (l) + 3O2 (g) 2CO2 (g) + 3H2O (g)

It has also been used as fuel-additives in automobiles, up to 20% ethanol

As a renewable resource:

Ethanol can be a renewable resource as it can be derived from non-fossil fuel sources such as

the fermentation of glucose

The raw materials its made from can be replenished each year by the sun

It is made from carbon dioxide, water and sunlight (via glucose)

This glucose can be derived from bacterial decomposition of cellulose (an abundant material)

or from starch

Describe the conditions under which fermentation of sugars is promoted:

Fermentation is the process in which glucose is broken down to ethanol and carbon dioxide by the

action of enzymes present in yeast

Conditions for fermentation:

Suitable grain or fruit mashed up with water

Yeast is added

Air is excluded (anaerobic environment)

Mixture is kept at about 37°C

Summarise the chemistry of the fermentation process:

Yeast is added to mashed grain and water

The yeast and other microbes break down the large carbohydrates 9starch or sucrose) into simpler

sugars (glucose) which are then fermented

sucrose + water glucose + fructose

C12H22O11 (aq) + H2O (l) C6H12O6 (aq) + C6H12O6 (aq)

In an anaerobic environment, the yeast uses their enzymes to break down the sugars, forming ethanol

and CO2 as products.

C6H12O6 (aq) 2C2H5OH (aq) + 2CO2 (g)

When ethanol concentration reaches 15%, the yeast die and fermentation STOPS.

Distillation is used to obtain higher ethanol concentrations (95-100%).

Define the molar heat of combustion of a compound and calculate the value for ethanol from

first-hand data:

The molar heat of combustion is the heat released when one mole of a substance undergoes

complete combustion with oxygen at standard atmospheric pressure with the final products being

carbon dioxide gas and liquid water.

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

disadvantages of its use:

Advantages of using ethanol as a fuel

It is a renewable resource and so would reduce the use of non-renewable fossil fuels

Easily mixed with existing fuels

Creates more job opportunities

Economic benefits because it reduced wastes from agricultural and human wastes

Improve energy security from foreign fuels

Burns cleanly as it undergoes complete combustion

It is ‘greenhouse neutral’ because the carbon dioxide produced from the production of

ethanol and combustion is proportional to the intake of carbon dioxide by plants to

photosynthesis

Fermentation: C6H12O6 (aq) 2C2H5OH (aq) + 2CO2 (g)

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

Photosynthesis: 6CO2 (l) + 6H2O (l) C6H12O6 (aq) + 6O2 (6)

Disadvantages of using ethanol

Large areas of agricultural land would need to be devoted to growing suitable crops with

consequent environment problems such as soil erosion, deforestation, fertiliser runoff and

salinity.

Less land devoted to growing crops and food, making food prices increase

Disposal of the large amounts of smelly waste fermentation liquors after removal of ethanol

would also prevent major environmental problems

Costs more to produce than fossil fuels

Produces less energy than some fuels, so more ethanol would be needed to go the same

distance compared to petrol

Engines must be modified to run on fuel containing more than 20% ethanol

Engines wear down faster due to the need for higher engine compression ratios needed for

ethanol combustion

Identify the IUPAC nomenclature for straight-chained alkanols from C1 to C8:

– Alkanols are a group of alkanes where one or more hydrogens have been replaced by the hydroxyl

(–OH) functional group

– When naming alkanols, there are specific rules:

The number of carbons determines the prefix of the name:

#C’s 1 2 3 4 5 6 7 8

Prefix methane- ethane- propane- butane- pentane- hexane- heptane- octane-

Process information from secondary sources to summarise the processes involved in the

industrial production of ethanol from sugar cane:

Sugar cane

Harvested and transported to processing mill

Shredded and crushed by a hammer mill to extract sugar rich juice

(remaining fibre cellulose waste bagasse burnt in the mills boiler furnaces to produce electricity

that is used for processing operations within mills)

Heated under pressure and lime added (to get dust and dirt out)

Passed through a clarifier where clarified juice is drawn off from the top – impurities settle to

bottom.

Cool and add yeast, the sugar undergoes fermentation – glucose and fructose (

Fermentation stops when ethanol is greater than 15% as yeast dies

Distillation of mixture to produce ethanol (fractional distillation to produce 95% ethanol/water

mixture)

Add drying agent to remove remaining water

Process information from secondary sources to summarise the use of ethanol as an

alternative car fuel, evaluating the success of current usage:

Advantages :

Ethanol can be used as an additive in some petrol blends E10 as it blends well with petroleum,

reducing carbon monoxide emissions as well as reducing ozone formation

Since 1975 25% of Brazil’s cars ran on 100% ethanol

It’s cleaner than petroleum, reduces carbon levels

Comes from renewable resource –biomass

Carbon neutral and greenhouse neutral (equations)

Complete combustion of ethanol requires less oxygen than octane, therefore complete combustion is

more likely with ethanol

C2H5OH (l) + 3O2 (g) 2CO2 (g) + 3H2O (g)

C8H18(l) + 15/2 O2 (g) 8CO2 (g) + 9H2O (g)

Ethanol has lower ignition temperature than petrol so it is more readily and easier to ignite

Creates jobs and sustains rural industry

Disadvantages:

Australia’s land is not arable (sufficient quality to grow crops) as the land is dry and has high

salinity. Therefore finding land to grow biomass for fuel is difficult in Australia

Federal government has imposed a 10% cap on ethanol content of transport fuel sold in Australia

Car manufacturers have widely claimed that amounts above 10% could damage car engines and may

void warranties. Problems include possible perishing and swelling of materials in fuel systems and

corrosion of engine components.

Cars need to be modified in order to run on ethanol >10%

Ethanol has higher flash point than petrol, so its harder to combust in colder conditions

Even when completely combusted petrol can produce more energy than complete combustion of

ethanol, therefore combustion of ethanol can only be economically attractive if ethanol is cheaper

than petrol since it will need to be replaced frequently

JUDGEMENT

PRACTICAL – Solve problems, plan and perform a first-hand investigation to carry out the

fermentation of glucose and monitor mass changes:

– A 250 mL side-arm conical flask with a rubber stopper was used. A plastic hose was connected to

the side arm, and the end of the hose was placed in another conical flask in a solution of limewater.

No gas was allowed to escape the apparatus:

– 100 ml of 0.15 M sucrose solution was placed in the conical flask. ONE gram of yeast was placed

into the sucrose. This was mixed thoroughly.

– The stopper firmly put on, and the flask was WEIGHED with an electronic scale.

– Yeast beaker was placed in incubator (37°C).

– Both flasks were weighed daily for 5 days.

– RESULTS:

The yeast flask turned foamy and smelt clearly of alcohol, while the limewater turned cloudy;

this proved that CO2 and ethanol were produced, and that fermentation occurred.

The mass of the yeast flask also steadily decreased by about half a gram each day; this is due to

the carbon dioxide escape; the limewater flask also gained approximately the same mass.


A “closed” system was used to ensure an accurate experiment.

Limewater was employed to prove CO2 was produced.

The incubator ensured that the most optimal fermentation occurred.

– LIMITATIONS of method:

The combined masses of both flasks steadily decreased as well; this was due to inevitable

leakages of gas.

The atmosphere in the flasks was not anaerobic (oxygen-free) and this could have hampered the

fermentation process.

Present information from secondary sources by writing a balanced equation for the

fermentation of glucose to ethanol:

– Fermentation of glucose:

4. Oxidation-reduction reactions are increasingly important as a source of energy:

Explain the displacement of metals from solution in terms of transfer of electrons:

A displacement reaction is a reaction in which a more reactive metal changes a less reactive metal’s

IONS into solid ATOMS. That is, the less reactive metal’s ions are ‘displaced’ out of solution and

neutralised into atoms.

Displacement reactions are actually electron transfer reactions, where one substance donates

electrons to another.

Metal displacement reactions are an example of electron transfer reactions. In these reactions a more

reactive solid metal will displace the ions of a less active metal from solid. This means that the more

active solid will dissolve (oxidise) to form ions and the less active metal ions are reduced forming a

solid deposit.

Oxidation Is Loss, Reduction Is Gain of ELECTRONS.

Hence, the 2 half-equations above can be labelled as oxidation or reduction

Zn Zn2+ + 2e¯ (Oxidation; zinc LOSES electrons)

Cu2+ + 2e¯ Cu (Reduction; copper GAINS electrons)

The species that is oxidised is the reductant (thus, zinc is the reductant)

The species that is reduced is the oxidant (thus, copper is the oxidant)

Identify the relationship between displacement of metal ions in solution by other metals to

the relative activity of metals:

– Only a more reactive metal will displace a less reactive metal.

– Using this fact, a table of reactivity, or ‘metal activity series’ can be formed, with the more reactive

metals on the left, and the less reactive metals on the right.

– THE METAL ACTIVITY SERIES:

– This ‘series’ was deduced through experimentation; it should be learnt.

– Thus, a metal on the series can displace out of solution ANY metal on its right, but cannot displace

any metal on its left; hydrogen is included as a standard.

– What the ‘metal activity series’ implies:

The metals on the left are very reactive and hence LOSE electrons easily, and are thus likely to

be oxidised; most of the time they are reductants.

The metals on the right are very unreactive; but when they are ions, they GAIN electrons very

easily, and thus are easily reduced; most likely oxidants.

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

electrons:

Changes in the oxidation state of a species indicates whether it has been oxidised or reduced

Oxidation involves an INCREASE in oxidation state

Reduction involves a DECREASE in oxidation state

Sun of oxidation states = overall charge on the ion

Species Oxidation state

Elements in natural occurring state 0

Cations and anions Overall charge on ion

Combined oxygen - 2

Combined hydrogen + 1

Describe and explain galvanic cells in terms of oxidation/reduction reactions:

Galvanic cell is a battery

It uses chemical reactions to produce electrical energy spontaneously

It consists of two different half equations – the oxidation half cell and the reduction half cell – where

each half cell consists of an electrode (metal of carbon) in an electrolyte

Has an external circuit, a wire connecting the anode to cathode (the electrons) and allows

ELECTRONS to flow from the anode to the cathode.

Has salt bridge that allows for the MIGRATION OF IONS between the half cells and so maintains

the electrical neutrality of the half cells.

Outline the construction of galvanic cells and trace the direction of electron flow:

– Electrically neutral solutions are needed for optimal electricity production. Role of the salt bridge:

The salt bridge completes the circuit

The salt bridge maintains electrical neutrality; this means that it keeps the charges in both the

half-cells at zero, by allowing the flow of ions.

Define the terms anode, cathode, electrode and electrolyte to describe galvanic cells:

Anode: the electrode at which oxidation occurs: negative electrode in galvanic cell

Cathode: the electrode at which reduction occurs: positive electrode in galvanic cell

Electrolyte: the salt solution or paste which conducts electricity

Oxidant: the species that causes the reduction of another species and so it itself is oxidised; the

reducing agent

Reductant: the species that causes the oxidation of another species and so it itself is reduced; the

oxidising agent

Salt bridge: an electrolyte soaked filter paper that forms a connection between the two half cells.

AN OX – Anode is OXIDISED (it is the “negative” side).

RED CAT – Cathode is REDUCED (it is the “positive” side).

o In galvanic cells, electrons flow from anode to cathode.

PRACTICAL – Perform a first-hand investigation to identify the conditions under which a

galvanic cell is produced:

– A 2 cm strip of zinc and copper were cut from metal strips. Using wire leads and crocodile clips, the

zinc strip was connected to the NEGATIVE terminal of a voltmeter, and the copper strip connected

to the POSITIVE terminal

– The zinc was then placed in 50 mL of 1M solution of zinc sulfate, and the copper in 50 mL of 1M

solution of copper sulfate.

– A strip of filter paper was soaked in potassium nitrate; the two cells were then connected using this

‘salt bridge’

Solve problems and analyse information to calculate the potential E° requirement of named

electrochemical processes using tables of standard potentials and half-equations:

The standard electrode potentials of half cells are listed as REDUCTION half equations

The strongest reductants (species that prefer to be oxidised) are at the top – more reactive

The strongest oxidants (species that are reduced) are at the bottom

Rules for predicting a spontaneous redox reaction

Select the reduction half equation from the table (states are important) and record the E°

Select the oxidation half equation by reversing the appropriate reduction half equation (write

it from right to left) and reverse the sign of the E °

Balance the electrons in the two half equations by multiplying by an appropriate coefficient,

but DO NOT change E°

Add the two half equations cancelling out the electrons and add the E ° .

A positive E° indicates a spontaneous redox reaction

5. Nuclear chemistry provides a range of materials:

Distinguish between stable and radioactive isotopes and describe the conditions under

which a nucleus is unstable:

Radioactivity is the spontaneous emission of radiation from certain atoms

Stable isotopes don’t spontaneously change into another element

Atomic number: All atoms with more than 83 protons are radioactive (Z>83)

Ratio for light elements Z<20 has a one neutron : one proton ratio 1:1 (n:p)

Proton-Neutron Ratio: The ratio of protons to neutrons determines whether an atom will be stable or

not. Outside the zone of stability n:p ratio is greater

As atomic number increases the stability ratio of neutrons to protons increases.

As an atom is small, it has the same number of protons and neutrons

As it gets larger and larger the number of protons increases and as they are positive, they repel each

other, so more neutrons are needed as they are spaces that are placed between the protons.

Describe how transuranic elements are produced:

Transuranic elements are artificial elements with atomic number greater than 92 – uranium (Z>92)

which has the largest value of the natural occurring elements.

Transuranic elements are produced in two ways:

o Neutron bombardment (in nuclear reactors): In nuclear reactors, the fission chain reaction produces

large amounts of neutrons. When atoms are placed inside the reactors, they are bombarded by

these neutrons (slow thermal neutrons). Occasionally the atom absorbs one of these neutrons;

however, it is unstable and undergoes beta decay. Hence the proton number increases, and a

transuranic element can be created.

o Fusion reactors (in particle accelerators): the production of larger transuranic elements in achieved

by colliding heavy nuclei with high speed positive particles. The positive particles need to be at

very high speeds to overcome the positive repulsive force of the heavy nuclei and fuse with them.

Particle accelerators are used to bring these particles to the high speed required in order to

overcome the electrostatic repulsion.

In a linear accelerator, positive particles are accelerated in a straight time along the axes of a

series of cylinders made alternatively positive and negative so the particles are always being

pushed from behind by a positive cylinder and pulled from in front by a negative one.

Cyclotrons use in addition a strong magnetic field to constrain the particles to a spiral path.

Describe how commercial radioisotopes are produced:

Commercial radioisotopes are isotopes that are produced on a regular basis for medical, industrial or

other uses.

Many are produced by bombardment of a target nucleus with a specific particle resulting the

absorption of that particle by the target nucleus.

Identify instruments and processes that can be used to detect radiation:

Cloud chamber : can be used to detect alpha, beta and gamma. The ionising radiation travels

through cold supersaturated alcohol vapour ionising the molecules to produce ions. Vapour

condenses on these ions and creates small droplets showing the path of the radiation.

o Ionising power

o Radiation passes through supersaturated alcohol vapour, ionising some of the air

o Ions formed act as nuclei upon which droplets of liquid are formed, path of radiation is made

o Alpha: dense straight lines

o Beta: dense zig-zag tracks

o Gamma: fainter tracks

Geiger-Muller counter:

o Ionising power

o Rays enter through thin end window of Geiger tube, hits argon gas molecules

o Ionises it by knocking an electron out

Metal tube filled with argon gas

When charged particles hit argon, the atom ionises to produce an electric current

Electric current is measured and shows up on measuring instrument

Photographic film:

o Darkening of photographic film

o Laboratory workers wear radiation badges and t he amount of darkening of the film is a measure

of the amount of radiation that worker has received/been exposed to.

Scintillation counter:

o Causes particular substances to emit light

o Substances containing alpha, beta and gamma, emit a flash of light which is collected and

amplified in photomultiplier

o Electrical signal generated operates an electronic counter

Identify one use of a named radioisotope in industry and in medicine and describe the way in

which the above named industrial and medical radioisotopes are used and explain their use

in terms of their chemical properties:

Medicine: Cobalt-60

Radiation therapy, treating cancer by irradiating the areas of the body with gamma ray which kill

cancers cells along with good cells

Made by neutron bombardment of Co-59

Half life of 5.3 years so machinery can last quite long and doesn’t need to be frequently replaced

and maintained

Gamma radiation which can penetrate deep into body tissues where the cancer may be, destroys

their DNA and kills cancer cells and stops growth and spread of cell

Healthy cells are also killed, staff must be continually protected from exposure

Industry: Sodium-24

Detects leaks in underground water or oil pipes by adding liquid and scanning along the pipe. If

tracer leaks into the soil surrounding the pipe it can be detected and leak is located.

Formed by neutron bombardment of sodium-23

Half life of 15 hours so that after leak is detected, radioactivity quickly decays and water or oil is

safe to use again

Gamma radiation so that is can penetrate though the ground

Process information from secondary sources to describe recent discoveries of elements:

Hassium -265: was produced in 1984 by firing Fe-58 at Pb-208 targets

Use available evidence to analyse benefits and problems associated with the use of

radioactive isotopes in identified industries and medicine:

Benefits:

Industrial: allows for more sufficient processes and previously impossible things

Medical: created a new wide range of non-invasive diagnostic techniques

Problems:

Causes dangerous mutations to cell DNA and caners in those

Those working with radioisotopes must wear protective clothing and take precautions to

avoid over exposure – allows for their safe use

Radioisotopes need to have their wastes disposed of and transported properly as the leakage

of radioactivity could have devastating consequences on environment.

Would lead to increase only if there was contamination of land, food supply and population

Careful planning and proper provisions must be made

Nuclear reactors produce considerable amounts of nuclear waste, which we have no way of

disposing safely, and which last for thousands of years

The storage of radioactive material presents a problem, as they must be kept in shielded

containers to prevent radiation leakages

Doses of radiotherapy must be extremely carefully controlled, to balance between the

benefits of killing cancer cells and the risk of harm

Nuclear technicians and other workers must be continually protected and avoid any form of

irradiation as diseases such as cancer or radiation poisoning can result from this

9.3 – The Acidic Environment:

1. Indicators were identified with the observation that the colour of some flowers depends

on soil composition:

Classify common substances as acidic, basic or neutral:

– Common acids:

Vinegar (acetic acid), vitamin C (ascorbic acid), lemon juice (citric acid), aspirin (acetyl salicylic

acid), ‘fizzy’ drinks (carbonic acid), car battery fluid (sulfuric acid).

– Common bases:

Drain cleaners (sodium hydroxide), household cleaners (ammonia), antacid tablets (calcium

carbonate), baking powder (sodium bicarbonate), washing powder (sodium carbonate).

– Common neutral substances:

Pure water, table salt (sodium chloride; but NOT all salts are neutral), milk, oils and other fats,

sugars.

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:

Indicator Colour Approximate pH

range

To test whether a

solution is:

Methyl red Pink / yellow 4.4 – 6.2 Highly or slightly acidic

Phenolphthalein Colourless / pink 8.3 / 10.0 Highly or slightly basic

Bromothymol blue Yellow / blue 6.0 / 7.6 Acidic, neutral or basic

Methyl orange Red / yellow 3.1 / 4.4 Highly or slightly acidic

Litmus solution Red / blue 5.0 / 8.0 Acidic, neutral or basic

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

acidity/basicity

– Testing Soil pH:

Some plants only grow within narrow pH ranges, so the pH of the soil needs to be regularly

tested.

Mix soil with indicator until it turns into a thick paste

Add white powder and wait approximately 1 minute. Read from colour chart to determine pH

Powder absorbs soil water and indicator colour can be seen against white background

– Testing pH of Pools:

Pool water must be near neutral to avoid health problem such as irritating eyes and skin

A few drops of indicator is placed in a sample of the pool water; place cap and shake to mix.

– Monitoring pH of Chemical Wastes:

Wastes that are produced from laboratories or photographic film centres tend to be highly acidic.

The pH of the wastes must be neutralised before they can be safely disposed.

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

Red cabbage was chopped to fine pieces

Boiled in hot water for 10 minutes

Collected liquid formed (light purple colour)

Added 3 drops of red cabbage solution into nitric acid, distilled water and sodium hydroxide

Acid (HNO3): red

Neutral (water): yellow/green

Base (NaOH): purple

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

range of indicators:

2. While we usually think of the air around us as neutral, the atmosphere naturally

contains acidic oxides of carbon, nitrogen and sulfur. The concentrations of these acidic

oxides have been increasing since the Industrial Revolution:

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

they act as acids:

Oxides are compounds that contain oxygen

Acidic oxides:

Non-metals

Reacts with water to form an acid

Reacts with bases to form salts

Can be neutralised by strong bases to form salts and water

CO2 (g) + H2O (l) H2CO3 (aq) (carbonic acid)

Basic oxides:

Metals

Reacts with acids to form salts

Does not react with alkali solutions

Reacts with water to form bases

Na2O (s) + H2O (l) 2NaOH (aq) (sodium hydroxide)

Amphoteric oxides:

Oxides that can act as both acids and bases

E.g. Beryllium, aluminium, zinc, tin and lead

Analyse the position of these non-metals in the Periodic Table and outline the relationship

between position of elements in the Periodic Table and acidity/basicity of oxides:

– The acidic oxides are on the RIGHT side of the periodic table (non-metals).

– The basic oxides are on the LEFT side of the periodic table (metals).

– The amphoteric oxides are in-between; NOBEL gases have no oxides.

– Increase in acidic character as you go towards right

Define Le Chatelier’s principle:

Equilibrium:

Reversible reactions are considered to be composed of a forward reaction and a reverse reaction

When reactants are first mixed in a closed system the rate of the forward reaction is usually high

as the number of collisions between reactants is high

The rate of the reverse reaction is initially low as the concentration of the products is low at the

start of the reaction

Overtime the forward reaction rate decreases and the reverse reaction rate increases. Eventually

these opposing rates become equal and the reaction has reached equilibrium

Chemical equilibrium in a closed system at constant temperature is characterised by:

Constant macroscopic properties such as colour and pressure

Equal but opposing rates of reaction

Le Chatelier’s principle states that when a closed system at equilibrium is disturbed it will counteract

the change in order to minimise the disturbance.

Identify factors which can affect the equilibrium in a reversible reaction:

Change in CONCENTRATION

Increasing the concentration by adding more of that species, causes the equilibrium to shift to

the side which uses up the added species

Adding a solid (which is present as a solid in the equilibrium reaction) or any pure liquid

like water does NOT increase its concentration and hence has NO EFFECT on the

equilibrium position

Decreasing the concentration of a species by removing that species, causes the equilibrium to

shift to the side that produces the removed species

Removing a substance which is present as a solid in the equilibrium reaction) does NOT

affect the equilibrium.

Change in GAS PRESSURE (or change in VOLUME)

Increase the volume of a system involving gas results in a decrease in the overall pressure of the

gases. It causes the reaction to shift to the side producing the most gaseous molecules in order to

increase pressure

Decrease the volume of a system involving a gas results in an increase in the overall pressure of

the gases. It causes the reaction to shit to the side producing the least gaseous molecules in

order to reduce pressure

Change in TEMPERATURE

Increase temperature by adding heat causes the reaction to shift in the endothermic direction,

that is to the side that uses up the added heat

Decrease temperature by removing heat or cooling causes the reaction to shit in the exothermic

direction, that is to the side that produces the removed heat

Describe the solubility of carbon dioxide in water under various conditions as an

equilibrium process and explain in terms of Le Chatelier’s principle:

The reaction is exothermic

Shift to the left release carbon dioxide

Shift to the right dissolved carbon dioxide

If an acid is added, the concentration of H+ will increase and the reaction will shift to the left,

producing more gas

If a base is added, the OH- will react with the H+, producing water and reducing the concentration of

the H+, so the reaction will shift to the right to produce more H+.

If the soda water is warmed, the reaction will shift to the left, releasing carbon dioxide, to use up the

added heat and thus the drink becomes less fizzy.

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

Sulfur dioxide

NATURAL

Bacteria may decompose organic matter to produce H2S, and when that oxidises, sulfur

dioxide is formed

Volcanic gases released during eruptions and from geysers

Smoke from bushfires

INDUSTRIAL

Burning of fossil fuels, coal, petroleum and industrial plants involving smelting of sulfide

ores that contain sulfur impurities, which when burnt, it oxidises to produce sulfur dioxide

Oxides of nitrogen

NATURAL

Lightning, which produces the high temperature condition

Nitrous oxide – action of certain bacteria on nitrogenous materials in soil

Anaerobic respiration of bacteria

INDUSTRIAL

Combustion, both in power stations and in vehicles.

At high temperatures in the combustion chambers, oxygen and nitrogen from air combine to

form nitric oxide, which is slowly converted to nitrogen dioxide

Use of nitrogenous fertiliser which provides more raw material for bacteria

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

chemical reactions which release oxides of nitrogen:

– Sulfur Dioxide:

When organic matter decomposes it produces hydrogen sulfide (H2S), which then oxidises

(reacts with oxygen) to produce sulfur dioxide:

2H2S (g) + 3O2 (g) 2SO2 (g) + 2H2O (l)

The burning of sulfur-rich coal and other fossil fuels directly combines sulfur with oxygen:

S (s) + O2 (g) SO2 (g)

The extraction of metals from metal sulfides also releases sulfur dioxide. E.g. smelting of zinc:

2ZnS (s) + 3O2 (g) 2ZnO (s) + 2SO2 (g)

– Oxides of Nitrogen:

Nitric oxide is produced either when lightning, with its high temperatures combines nitrogen and

oxygen:

N2 (g) + O2 (g) 2NO (g)

The same reaction occurs in the high temperatures of engines or power plants, also combining

nitrogen and oxygen.

Nitrogen dioxide is formed when nitric oxide reacts with oxygen in the air:

2NO (g) + O2 (g) 2NO2 (g)

Assess the evidence which indicates increases in atmospheric concentration of oxides of

sulfur and nitrogen:

Increasing of burning fossil fuels after Industrial Revolution lead to increase of oxides of sulfur,

evidence is the air quality of major industrial cities such as London, deteriorated greatly.

Power stations increased their burning of fossil fuels, which provided the high temperatures for both

nitrogen and oxygen gas to combine to form nitric oxide, which then reacts with oxygen to form

nitrogen dioxide.

N2 (g) + O2 (g) 2NO (g)

2NO (g) + O2 (g) 2NO2 (g)

In 20th century, motor vehicles increases and amount of oxides of nitrogen increase as motor engines

provided the high internal combustion for production of nitric oxide.

Quantitative analysis of trapped air bubbles in Antarctic ice and measurement of carbon isotopes in

old trees, grass seeds in museum collections and calcium carbonate in coral

Analysis of gas found in ice-core samples excavated from Antarctica shows the levels of N2O in

atmosphere has increased by 10%

NO2 leads to formation of photochemical smog, a direct indicator of excessive levels of nitrogen

oxides in atmosphere

Increase of acid rain in previous years supports increase of oxides of sulfur and nitrogen

Analyse information from secondary sources to summarise the industrial origins of sulfur

dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the

environment:

CONCERNS:

Acid rain: Both sulfur dioxide and nitrogen oxides are acidic oxides that react with water to form

acid rain, which is very destructive; it can destroy forests, corrode limestone buildings and

structures, disrupt natural ecosystems and alter the pH levels of oceans, lakes and rivers.

Health problems: Sulfur dioxide irritates the respiratory system and causes breathing difficulties

at concentrations as low as 1 ppm. It triggers asthma attacks ad aggravates emphysema.

Nitrogen dioxide irritates the respiratory tract and causes breathing discomfort at concentrations

above 3 -5 ppm, as it begins to destroy tissue, as it forms nitric acid.

Photochemical smog: NO2 causes the formation of photochemical smog which is an air pollution

in which sunlight reacts with nitrogen dioxide, hydrocarbons and oxygen to form ozone, other

pollutants such as PANS and haze, which is poor visibility due to small particles in air.

Explain the formation and effects of acid rain:

Acid rain is rain that has a higher hydrogen concentration than normal, pH<5

FORMATION of acid rain:

Sulfur dioxide reacts with rain in the atmosphere forming sulfurous acid:

SO2 (g) + H2O (l) H2SO3 (aq)

Sulfurous acid then reacts with oxygen; this is catalysed by air particles:

2H2SO3 (aq) + O2 (g) 2H2SO4 (aq)

Nitrogen dioxide also reacts with rain, making nitric and nitrous acids:

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

Nitrous acid then reacts with oxygen, again catalysed by air particles:

2HNO2 (aq) + O2 (g) 2HNO3 (aq)

Acid rain is formed when sulfur oxides and nitrogen oxides are absorbed by water droplets

EFFECTS of acid rain:

Environmental damage:

o Increasing acidity of lakes, which has detrimental effects on fish populations, as fishes are

unable to survive and reproduce, and the mucus on their scales are destroyed.

o Destroys forests and vegetation, as plant roots and leaves are damaged and lost and trees are

denuded.

o Soil chemistry can change, leading to death of important micro-organisms and release of

normal insoluble aluminium and mercury into soil water, which contribute to crop failure,

as their salts are more soluble at low pH’s

Urban and structural damage include the corrosion of marble statues and buildings (marble

contains carbonates which acids readily react with), also with metal structures.

Limestone/marble: CaCO3(s) + 2H+(aq) Ca2+

(aq) + CO3(g) + H2O(l)

Steel structures: Fe(s) + 2H+(aq) Fe2+

(aq) + H2(g)

PRACTICAL – Identify data, plan and perform a first-hand investigation to decarbonate

soft drink and gather data to measure the mass changes involved and calculate the volume of

gas released at 25˚C and 100kPa.

– A 250mL can of soft drink was decarbonated by opening the cap; no drink was allowed to spill.

– Another 250mL can of water was opened as a control, to compare the amount of water evaporated or

loss during that time.

– The bottle was weighed at the beginning, and the end.

– It was assumed that any mass loss was due to carbon dioxide loss.

– RESULTS: mass of water loss – mass lost in soft drink


The shaking affected the CO2/H2CO3 equilibrium, and this forced the gas out.

No liquid was allowed to spill to try to keep a high level of accuracy.

Electronic scales were also used for accuracy.


The accuracy of this method is limited, as there may be loss of water as gas, or there may still be

CO2 dissolved.

3. Acids occur in many foods, drinks and even within our stomachs:

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

– An acid is a substance that releases H+ ions. In water, acids ionise (separate into its ions)

– When acids react with other substances, the H+ ion is transferred to another species, proton-donor

Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic),

hydrochloric and sulfuric acid:

Common name for acid Acetic acid Citric acid Hydrochloric acid

Systematic name Ethanoic acid 2 – hydroxypropane -

1,2,3 – tricarboxylic acid

Muriatic

Chemical formula CH3COOH C6H8O7 HCl

Strength Weak acid Weak – ish acid Strong acid

Degree of ionisation 1% 8% 100%

Ionisation equation

Diagram modelling the

ionisation of the acid

solution

Found / use Present in vinegar Occurs in citrus fruit Produced in stomach to

aid in digestion

Describe the use of the pH scale in comparing acids and bases:

The pH scale is used to determine the acidity or alkalinity of a substance

The pH scale allows us to compare acids and bases, and their strengths

Describe acids and their solutions with the appropriate use of the terms strong, weak,

concentrated and dilute:

– A strong acid is an acid that releases ALL its H+ ions when in solution (2 particles present only)

Ionises completely, has a high dissociation constant, donates protons freely

The solution contains only hydronium ions and the anions of the acid; there are no neutral

acid molecules present

Concentration for strong acids can be found directly from thr concentration of the acid using

the mole ratio for the ionisation reactions as 0.1mol L-1 HCl contains 0.1mol L-1

Its molecules are completely dissociated in solution

Solutions of strong acids are better conductors of electricity than weak acids

Examples : sulphuric acid (H2SO4), hydrochloric (HCl) , hydrobromic (HBr), hydriodic (HI),

nitric (HNO3)

Equation shows complete ionisation; only the ions are present

EG: HCl (g) + H2O (l) H3O+ (aq) + Cl־

(aq)

– A weak acid is an acid that does NOT completely release all its H+ ions (3 particles present)

Ionises partially – some of its molecules remain intact in solution, poor proton donor

Concentration for weak acids can only be found if degree of ionisation of acid is known

EG: 0.1mol L-1 acetic acid, degree of ionisation is 1% so solution contains 0.1mol L-1 H+

The solutions is an equilibrium between the neutral acid molecules and the hydronium ions

and anions of the acid, equations shows equi8librium that exists, all species are present

Examples : R – COOH acids, carbonic acid (H2CO3), sulfurous (H2SO3), nitrous (HNO2),

phosphoric (H3PO4)

Its ionisation reaction with water is a reversible reactions that reaches equilibrium when a

certain number of H+ ions are released

EG: CH3COOH (s) + H2O (l) H3O+ (aq) + CH3COO־

(aq)

Can describe the difference between strong and weak acid in terms of the equilibrium between its

molecules and its ions. Strong – equilibrium lies to right, all molecules ionise. Weak – equilibrium

lies to left, only small percentage of molecules ionise, rest remain as molecules of acid]

The terms concentrated (high molarity) and dilute (low molarity) refer ONLY to the amount of acid

molecule present in the solution, they have nothing to do with ionisation

CONCENTRATED solution: total concentration of solute species is high – more than 5 mol/L of

solute. Contains a large amount of solute in a given amount of solution

DILUTE solution: total concentration of solute species is low – has less than about 2 mol/L of solute.

Contains a small amount of solute in a given amount of solution

Gather and process information from secondary sources to explain the use of acids as food

additives:

– Acids are added to food for 2 reasons: as preservatives, and to add flavour.

– PRESERVATIVES:

Ethanoic acid (acetic acid in the form of vinegar) is used as a preservative in ‘pickling’ as it

lowers the pH, preventing the growth of bacteria, which some may be harmful.

Sulfur dioxide is added to many foods as a presercative of the gas, as a solution in water. It is

used to prevent dried fruits from being attached by moulds and other microbes

Citric acid is a natural preservative, often added to jams and conserves. Gives a tart taste and the

acidity and high sugar content also prevents the growth of microbes. Also used as an antioxidant

and prevents microbe growth in canned foods

– FLAVOURINGS:

Carbonic acid is added to soft drinks to add ‘fizz’.

Phosphoric acid is also added to soft drinks to add ‘tartness’ of flavour.

Ethanoic acid, as vinegar, is also used as flavouring.

– Acids are added to processed foods to:

increase nutritional value and to enhance flavour of food

prevent spoilage (acids reduce pH to a level where the microorganisms that cause spoilage can

no longer reproduce)

– Lemon juice is used to coat cut apple, pear and bananas to prevent it from ‘browning’. It inactivates

natural enzymes in the food that cause the fruit to discolour.

– Purpose of cream tartar in baking powder and self raising flower: when it becomes moist the baking

soda reacts with the acid to form carbon dioxide gas which helps make the cake ‘rise’

– Acids can be used to preserve food because many microorganisms are sensitive to an acidic

environment. Adding acids to food changes the pH so that it is outside the range in which the

microorganism can survive.

Identify data, gather and process information from secondary sources to identify examples of

naturally occurring acids and bases and their chemical composition:

– NATURAL ACIDS:

Hydrochloric acid: Aqueous HCl; it is produced naturally by the lining of our stomachs. It aids

in the digestion of food.

Citric acid: C6H8O7; occurs naturally in large quantities in citrus fruits.

Ethanoic acid: CH3COOH; it is found naturally in vinegar, which is produced by the natural

oxidation of ethanol.

Lactic acid: C3H6O3; it is formed in the body during strenuous exercise, muscle tissues, milk and

yoghurt

– NATURAL BASES:

Ammonia: NH3; it is present in the stale urine of animals. It is also formed through the anaerobic

decay of organic matter. Household cleaners, removes grease from floors or to clean windows.

Fertilisers and extraction of nickel.

Metallic oxides: E.g. iron(III) oxide, copper(II) oxide and titanium(IV) oxide. These insoluble

oxides are solid bases found in minerals.

Calcium carbonate: CaCO3; it is found naturally as limestone. Used in plaster and cement

PRACTICAL – Solve problems and perform a first-hand investigation to use pH

meters/probes and indicators to distinguish between acidic, basic and neutral chemicals:

– Part A - Testing Substances With Methyl Orange:

10 test-tubes were set up with 5 mL of acidified solution (0.1 M HCl).

2 drops of methyl orange were placed in each test-tube.

A range of substances was dropped in each test-tube, from various carbonates, metal oxides,

sulfates, and common household substances.

Any change, such as indicator colour change or bubbles, was recorded.

– RESULTS:

It is known that methyl orange is red in strongly acidic solutions, and yellow in slightly acidic to

highly alkaline solutions.

Any substance that caused the acidified solution to turn yellow was classified as a base. This

included:

Carbonates: Calcium carbonate and sodium hydrogen carbonate caused the solution to turn

yellow and formed bubbles.

Metal Oxides: Calcium oxide and zinc(II) oxide turned the solution yellow, without any

bubbles.

Hydroxides: Magnesium hydroxide rapidly turned the solution yellow.

The other substances caused no change in the solution; these were labelled acidic or neutral.

These included sucrose, salt (NaCl), lemon juice, vinegar and magnesium sulfate.

– Part B – Testing Common Substances Using a pH Meter:

5g of the substance was dissolved in 20 mL of water and the pH tested.

They were then classified as acids, bases or neutral.

Neutral: Both table salt and milk were neutral.

Bases: Washing powder, the antacid tablet and toothpaste.

Acids: Soil, vitamin C tablet and aspirin.


Methyl orange is only red in strongly acidic solutions, so even a very weak base would have

been able to elicit a colour change.

A wide range of substances was used to portray the wide range of possible bases. This showed

that bases are not only limited to metal hydroxides.

A pH meter gave instantaneous accurate results.

Equal amounts of each substance, dissolved in 20 mL of water made for a fair test of the

substance’s pH.


The methyl orange test was not able to distinguish acidic and neutral substances, as both caused

no colour change from the red.

PRACTICAL – Plan and perform a first-hand investigation to measure the pH of identical

concentrations of strong and weak acids:

– 20 mL of 0.1 M solutions of sulphuric, citric, hydrochloric and acetic acid were placed in four

separate beakers.

– A pH meter was used to measure the pHs of the solutions.

– RESULTS:

Sulfuric acid; pH = 0.7

Hydrochloric acid; pH = 1.2

Citric acid; pH = 2.9

Acetic acid; pH = 3.3

– Sulfuric acid has the lowest pH of all the solutions as it is a diprotic strong acid, where was

hydrochloric acid is a monoprotic strong acid. Citric and acetic acid are weak acids.

4. Because of the prevalence and importance of acids, they have been used and studied for

hundreds of years. Over time, the definitions of acid and base have been refined:

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

and Arrhenius:

Scientist

Date

Observations made Theory of acids Limitations of theory and

examples

Lavoisier Dissolving acids of non-metals

in water formed acids

“All acids contain Macro level observations

CO2 (g) + H2O (l)

1780 Oxides of non-metals when

dissolved in water form acids, all

acids contain oxygen

oxygen”

Presence of oxygen in

compounds formed from

non-metals causes acidity.

H2CO3 (aq)

SO3 (g) + H2O (l)

H2SO4 (aq)

Davy

1815

Decomposed HCl by electrolysis

and found that it did not contain

oxygen.

Acids reacted with metals

forming salts

“All acids contained

hydrogen which could be

replaced by metals”

All acids contain

replaceable hydrogen

Macro level observations

2HCl (aq) + Mg (s)

MGCl2 (aq) + H2 (g)

Arrhenius

1884

Noticed that hydrogen gas was

evolved at the negative electrode

when an electric current was

passed through acids like

hydrochloric acid. Acids

conducted electricity

“Acids were substances

that ionised in solution to

produce hydrogen ions”

All acids produce

hydrogen ions in solution

ie they ionise to produce

H+

Conceptual definition

HCl (g) + H2O (l)

H3O+ (aq) + Cl- (aq)

CH3COOH (l) H+

(aq) + CH3COO־ (aq)

Bronsted /

Lowry

1923

An acid is a substance that in

solution tends to give up protons

(hydrogen ions) and a base is a

substance that tends to accept

protons

Acid – proton donor

Base – proton acceptor

An acid-base reaction is a

proton transfer reaction

HA + H2O A- + H3O+

acid base

B + H2O HB+ + OH-

base acid

Gather and process information from secondary sources to trace developments in

understanding and describing acid/base reactions:

Scientist(s) Acid Definition Base Definition History

Lavoisier Corrosive substances that contain oxygen -

Worked with metal oxides that form the common oxyacids with

water.Davy Corrosive substances

with hydrogen -Worked with hydrohalic acids such

as HCl and HBr, disproving Lavoisier.

Arrhenius Substances disassociate into H+

ions in solution

Substances that disassociate into

OH- ions in solution

Made theories of acid/base ionisation. Only aqueous (water)

solutions were considered.

Brönsted-Lowry

Proton (H+) donor Proton (H+) acceptor

Extended acid/base reactions to those without water. Acids must contain hydrogen; no solution

required

Outline the Brönsted-Lowry theory of acids and bases:

– B-L theory states that:

Acids are proton donors

Bases are proton acceptors

– When an acid donates a proton it forms its conjugate base, the conjugate base is a species which has

one less H and one more negative charge than its acid

– When a case accepts a proton it forms its conjugate acid. The conjugate acid has one more H and its

charge is increased by 1

– If a substance has a greater tendency to give up protons than the solvent, then that substance in that

solvent will be an acid

HA + H2O A- + H3O+

– If a substance has a greater tendency to accept protons than the solvent has, then the substance will

be a base in the solvent.

B + H2O HB+ + OH-

Identify conjugate acid/base pairs:

H2S H+ (aq) + Cl־

(aq)

HF H+ (aq) + HF־

(aq)

HCOOH H+ (aq) + COOH־

(aq)

Identify amphiprotic substances and construct equations to describe their behaviour in

acidic and basic solutions:

– Amphiprotic substances are substances that can act both as a proton donor and as a proton acceptor

– It can act as both an acid and a base and their behaviour depends on the environment they are placed

in

– NH3 NH2 + H+

– NH3 + H+ NH4+

Identify neutralisation as a proton transfer reaction which is exothermic:

– Neutralisation reactions are reactions between acids and bases

– As acids are proton-donors and bases are proton-acceptors then neutralisation reactions between

acids and bases are PROTON-TRANSFER reactions

– Protons (H+) are transferred from the acid to the base

– All neutralisation reactions are exothermic; they all liberate heat energy

– The ΔH ≈ -56 kJ/mol, depending on the strength of the reactants.

Analyse information from secondary sources to assess the use of neutralisation reactions as

a safety measure or to minimise damage in accidents or chemical spills:

– It is important to immediately neutralise any chemical spills involving strong acids and bases, as

they are corrosive and can be extremely dangerous

– Neutralisation reactions are widely used as safety measures in cleaning up after such incidents

– When neutralising an acid or a base, the following procedure is followed:

The most preferred agents of neutralisation has the properties of being stable easily transported,

solid (powdered), cheap and amphiprotic ( so it can act as a weak acid or a weak base) since it

can neutralise both acids and bases, even if an excess is used, it is very weak and so does not

pose any safety risks

The neutralised product is then absorbed using paper towels and disposed.

The speed of the reaction for neutralising the spilt material

The need for a reagent that will not have any harmful effect if an excess of it is used

The safety in handling and storing the reagent

The cost of the reagent

The possibility of the one reagent being able to neutralise both acid and alkali spills

– The most common substance used to neutralise spills in laboratories is powdered sodium hydrogen

carbonate; this is because

It is a stable solid which is easily and safely handled and stored

Cheapest alkali available

If too much of it is used, there is less danger than others

The hydrogen carbonate ion (HCO3-) is an amphiprotic species and it is cheap and readily

available substance

– Strong acids and bases must never be used to neutralise spills, if an excess is used, the spill will be

dangerous again

Describe the correct technique for conducting titrations and preparation of standard

solutions:

– Titration is a chemical technique used to experimentally determine the unknown concentration of a

solution through a chemical reaction.

– Acid/base titration uses a neutralisation reaction to determine this concentration.

– The Chemical Theory of Acid/Base Titration:

Titration is also known as volumetric analysis.

The point of acid/base titration is to determine the concentration of an unknown solution by

slowly reacting a certain volume of this solution with another solution of known concentration,

until the endpoint is reached.

The ENDPOINT of a chemical reaction occurs when all available molecules have reacted, and

the reaction comes to an end.

The volumes of the reactants at the endpoint are carefully measured; using the knowledge of

these volumes, and the original concentration of the standard solution, the concentration of the

unknown solution can be calculated.

– Primary Standards:

There are certain criteria chemicals need to satisfy before they can be used to create standard

solutions; these suitable chemicals, called primary standards, must have the following

properties:

VERY HIGH PURITY:

This is to produce accurate results, untainted by chemical impurities.

CHEMICAL STABILITY (low reactivity):

Standards must be chemically stable so they do not react violently with the water solvent,

or with gases in the atmosphere (e.g. CO2).

NON-HYGROSOPIC AND NON-EFFLORESCENT:

Hygroscopic substances absorb water from their surroundings, while efflorescent

substances release water into their surroundings. Both these processes change the

concentration of solutions, resulting in imprecise primary standards.

HIGH SOLUBILITY:

Primary standards need to dissolve completely into their solutions.

HIGH MOLECULAR WEIGHT:

The high molecular weight of primary standards sets off any errors in measurement that

may have occurred.

Hence, from the above example, we can see that sodium carbonate is indeed a suitable primary

standard; it can easily be obtained at extremely high purities, in a solid powdered form, and when

it forms solutions, they are stable.

Unsuitable chemicals to use as primary standards include sodium hydroxide, which in solution

will react with gases in the air, hydrochloric acid, which is efflorescent, and sulfuric acid, which

is severely hygroscopic.

– Preparing Standard Solutions:

The typical laboratory glassware in which standard solutions are made are called volumetric

flasks; 250 mL flasks are most commonly used.

Using the above example, we will determine the amount of primary standard needed to create

250 mL of the 0.05 M standard solution of sodium carbonate.

CALCULATIONS:

MR(Na2CO3) = 2 (22.990) + (12.011) + 3 (15.999) = 105.987

n = c × v = 0.05 x 0.25 = 0.0125 mol

m = n x MR = 0.125 x 105.987 = 1.325 g

Hence, to make up 250 mL of 0.5 M solution, 1.325 g are needed.

– METHOD :

1. Firstly, the primary standard must be as pure as possible, as this means it

must free of moisture. The primary standard must be placed in an

oven, and cooled in a dessicator to remove all traces of moisture.

2. Thoroughly rinse a 250 mL volumetric flask, a small beaker and a

glass funnel with distilled water. Place the funnel in the neck of the

volumetric flask, and place the beaker on an electronic scale.

3. Zero the scale, and using a very clean spatula, transfer as accurately as

possible, 1.325 g of sodium carbonate into the beaker.

4. Using a wash-bottle of distilled water, transfer the powder into the flask by ‘washing’ it into the

funnel. Ensure that the entire beaker is washed, and all water that touches the beaker flows into

the funnel, to ensure all solute is transferred. Wash the funnel, allowing the water to flow into

the flask, and then remove it.

5. Fill the flask half-way up to the 250 mL graduation mark, and gently swirl the flask until all the

solute has dissolved. Set the flask on the bench.

6. Using the wash-bottle, fill the flask with water until it is just under the graduation mark. Using a

Pasteur pipette, add distilled water until the meniscus sits exactly on the 250 mL mark.

7. Cover the flask with its glass stopper and invert the entire flask three times. If the meniscus has

lowered, add a few more drops. If not, the standard solution is complete.

– Determining The Location of the Endpoint:

The endpoint of a chemical reaction is the point where the reaction STOPS, because all the

species (of at least one reactant) have reacted.

In acid/base titration, the reaction is a neutralisation reaction and hence we use our knowledge of

pH to determine when the endpoint has been reached.

At the endpoint, all the species have reacted, and there is only the aqueous salt left; however, this

does NOT mean that the pH at the endpoint is 7.

The pH at the endpoint depends on the acid and base being used.

EG: For the titration above, nitric acid is titrated into the sodium carbonate:

Before titration occurs, for sodium carbonate, pH is much greater than 7.

As nitric acid is slowly added, the pH decreases steadily, until the endpoint is reached, at

22.2 mL of acid.

BUT the salt formed between a strong acid and a weak base is slightly acidic; hence at the

endpoint, the solution in the flask is slightly acidic.

This is the titration curve of the above reaction:

– INDICATORS:

To determine the pH, to determine the endpoint, we use indicators.

However, a suitable indicator must be used.

For strong-acid/weak-base titrations, as above, methyl orange is the most suitable indicator. This

is because it changes from yellow to pale orange/pink within the slightly acidic range,

corresponding with the endpoint of pH = 2.5:

For strong-base/weak-acid titrations, phenolphthalein is suitable.

– The Correct Technique for Conducting Titrations:

There is a very precise and specific technique to titration, which uses a variety of calibrated

glassware. These include:

Volumetric Flask: This is used to prepare and hold standard solutions.

Conical Flask: This is used to hold the reactants during titration. Its shape prevents the

reactants from spilling as they are swirled together.

Burette: The burette is a piece of cylindrical glassware, held vertically, with volumetric

divisions on its full length and a precision tap (stopcock), on the bottom. It is used to

dispense precise amounts of a liquid reagent in titration. Burettes are extremely precise and

accurate to ±0.05 mL.

Pipette: The pipette is a glass tube used to transfer precise volumes of liquid reagents.

Pipettes are usually designed to transfer one measurement of volume, such as only 25 mL.

The reagent is drawn up the pipette using a pipette filler (e.g. a rubber bulb).

Before titration can begin, all glassware must be RINSED appropriately. The proper rinsing

technique for the different glassware is:

Volumetric flasks (including glass stoppers) are rinsed thoroughly with distilled water,

preferably multiple times. Close the flask with the stopper until it is going to be used; it is left

wet.

Conical flasks are rinsed thoroughly with distilled water and left wet.

Burettes are rinsed with a specific technique; first the distilled water is filled into the burette

and the tap opened. Water is allowed to flow out and thoroughly rinse the tip. More water is

then added, and the entire glass tube is swirled in your hands to wash the sides of the burette.

The water is then poured out from the top. Burettes MUST be rinsed THREE TIMES with

distilled water and then ONCE more with the solution it is going to contain. No, it’s not

insane, it’s standard laboratory procedure.

Pipettes are also rinsed THREE times with distilled water, and then ONCE with the solution

it is going to contain.

After everything is rinsed appropriately, the glassware is filled.

Usually, the acid is placed in the burette, and the base in the conical flask, but it really doesn’t

matter much. Using the above example again:

Using a funnel, the nitric acid is poured into the burette until ABOVE the zero mark. Hold a

white card behind the zero mark, and open the tap slowly until the meniscus sits JUST on top

of the mark. The white card makes the meniscus clearer.

Using the pipette, a fixed volume of sodium carbonate (say 25 mL) is drawn from the

volumetric flask and deposited into the conical flask.

A few drops of methyl orange indicator are added to the conical flask, and the solution turns a

clear yellow colour.

– THE TITRATION:

Finally, the titration can be performed.

The conical flask is placed on a white tile (to make the solution’s colour clear) under the burette,

which is held in a retort stand.

The tap is slowly opened, and the conical flask is continuously swirled.

When the first colour change is noticed, the tap is immediately closed. A swirl of the conical

flask will likely return the solution back to its original colour.

Very slowly open the tap so that solution flows out in drops, and stop when the endpoint is

reached, as shown by the colour of the indicator:

In the case of the above example, the endpoint is pale pink.

The first titration performed is a rough draft and often overshoots the endpoint. This first titration

is rejected.

Titration is performed multiple times to achieve the accurate results.

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

system:

– A buffer solution is a solution that contains comparable amounts of a weak acid and its conjugate

base and which is therefore able to maintain an approximately constant pH even when an acid or

base is added to it

– Comparable amounts of a weak acid and its conjugate base

– Maintains constant pH when an acid or base is added to it

– A natural occurring buffer is the carbonic acid system which occurs in lakes and rivers and

maintains a constant neutral pH needed for life to exist

Buffer: H2CO3 / HCO3- comparable amounts

As rainwater falls through the air it forms a dilute solution of carbonic acid

CO2 (g) + H2O(l) H2CO3

HCO3- comes from dissolving of carbonate rocks e.g. limestone in surrounding areas around

the lake

Buffer equilibrium that exists:

PRACTICAL – Choose equipment and perform a first-hand investigation to identify the pH

of a range of salt solutions:

– In this practical, the pH of a range of salt solutions was estimated by examining the parent acids and

bases.

– The actual pH was then measured using a pH meter.

– Salts :

Sodium acetate - CH3COONa:

Parent acid is CH3COOH (weak) and parent base is NaOH (strong).

Predicted pH is slightly basic.

True pH = 8.9

Sodium nitrate - NaNO3:

Parent acid is HNO3 (strong) and parent base is NaOH (strong)

Predicted pH is neutral.

True pH = 7

Other salts:

NaHCO3 - pH = 8

MgCl2 - pH = 3

PRACTICAL – Perform a first-hand investigation to determine the concentration of a

domestic acidic substance using computer-based technologies:

– 50 mL of household vinegar was placed in a small beaker.

– A pH probe was attached to the laboratory computer; the probe was rinsed with distilled water, and

then rinsed with left-over solution.

– The probe was then placed in the solution, and the pH measured.

– The probe was rinsed again, and pH was measured twice more.

– The average pH was calculated, and using this, the concentration of H+ ions found

– Given the fact that food-grade ethanoic acid only has about 0.4% ionisation, the concentration of

ethanoic acid was then calculated.

– RESULTS:

The average pH measured was 2.5

[H+] = 10-pH = 10-2.5 = 0.00316 mol/L

But [H+] = 0.4% × c ; (where c = concentration of ethanoic acid)

Hence c = 0.00316 ÷ 0.004 = 0.79

Therefore vinegar is 0.79 M ethanoic acid.

5. Esterification is a naturally occurring process which can be performed in the

laboratory:

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

compounds:

–

Explain the difference in melting point and boiling point caused by straight-chained alkanoic

acid and straight-chained primary alkanol structures:

–

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

equations, examples of esterification:

–

Describe the purpose of using acid in esterification for catalysis:

–

Explain the need for refluxing during esterification:

–

Identify the IUPAC nomenclature for describing the esters produced by reactions of straight-

chained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8:

–

Outline some examples of the occurrence, production and uses of esters:

–

Process information from secondary sources to identify and describe the uses of esters as

flavours and perfumes in processed foods and cosmetics:

–

PRACTICAL – Identify data, plan, select equipment and perform a firsthand investigation

to prepare an ester using reflux:

9.4 – Chemical Monitoring and Management:

Δ. Construct word and balanced formulae equations of all chemical reactions as they are

encountered in this module:

EQUATIONS:

– Different products under different conditions:

Complete combustion:

propane + oxygen carbon dioxide + water

C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (g)

Incomplete combustion:

propane + oxygen carbon + carbon monoxide + water

C3H8 (g) + 3O2 (g) C (s) + 2CO (g) + 4H2O (g)

– Haber Process:

nitrogen + hydrogen ammonia + heat

N2 (g) + 3H2 (g) 2NH3 (g) ∆H = -92 kJ/mol

– Sulfate Content in Fertilizer:

barium chloride + sulfate barium sulfate + chloride

BaCl2 (aq) + SO42-

(aq) BaSO4 (s) + 2Cl- (aq)

– How Ozone Protects Us From UV Radiation:

Oxygen reacts with UV, forming 2 radicals:

O2 + UV radiation 2O·

Radical reacts with oxygen, forming ozone:

O· + O2 O3

Ozone reacts with UV, forming oxygen and a radical:

O3 + UV radiation O· + O2

Radical reacts with ozone, creating oxygen:

O· + O3 2O2

– All the CFC-Related Equations:

Formation of chlorine radical:

CCl2F2 (g) + UV radiation ·Cl (g) + ·CClF2 (g)

Reaction of ozone:

·Cl (g) + O3 (g) ·ClO (g) + O2 (g)

Regeneration of chlorine:

·ClO (g) + O (g) ·Cl (g) + O2 (g)

Removal of chlorine radical:

·Cl (g) + CH4 (g) HCl (g) + ·CH3 (g)

Removal of chlorine monoxide radical:

·ClO (g) + NO2 (g) ClONO2 (g)

Formation of molecular chlorine:

HCl (g) + ClONO2 (g) Cl2 (g) + HNO3 (g)

Decomposition of molecular chlorine:

Cl2 (g) + UV radiation 2·Cl (g)

– The Heavy-Metal Sulfide Test:

The sulfide-test is based on the following equilibrium:

S2- (aq) + 2H3O+

(aq) H2S (aq) + 2H2O (l)

1. Much of the work of chemists involves monitoring the reactants and products of

reactions and managing reaction conditions:

Outline the role of a chemist employed in a named industry or enterprise, identifying the

branch of chemistry undertaken by the chemist and explaining the chemical principle that

the chemist uses:

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

Describe an example of a chemical reaction such as combustion, where reactants form

different products under different conditions and thus would need monitoring:

Gather, process and present information from secondary sources about the work of

practising scientists identifying:-

2. Chemical processes in industry require monitoring and management to maximise

production:

Identify and describe the industrial uses of ammonia:

Identify that ammonia can be synthesized 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 that the use of a catalyst will lower the reaction temperature required and identify

the catalyst(s) used in the Haber process:

–

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 why the Haber process is based on a delicate balancing act involving reaction

energy, reaction rate and equilibrium:

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

discuss the monitoring required:

Gather and process information from secondary sources to describe the conditions under

which Haber developed the industrial synthesis of ammonia and evaluate its significance at

that time in world history:

3. Manufacture products, including food, drugs and household chemicals, are analysed to

determine or ensure their chemical composition:

Deduce the ions present in a sample from the results of tests:

Describe the use of atomic absorption spectroscopy (AAS) in detecting concentrations of

metal ions in solutions and assess its impact on scientific understanding of trace elements:

Gather, process and present information to interpret secondary data from AAS

measurements and evaluate the effectiveness of this in pollution control:

Gather, process and present information to describe and explain evidence for the need to

monitor levels of specific ions in substances used in society:

PRACTICAL – Perform first-hand investigations to carry out a range of tests, including

flame tests, to identify the following ions:-

PRACTICAL – Identify data, plan and select equipment and perform first-hand

investigations to measure the sulfate content of lawn fertiliser and explain the chemistry

involved:

Analyse information to evaluate the reliability of the results of the above investigation and to

propose solutions to problems encountered in the procedure:

4. Human activity has caused changes in the composition and structure of the atmosphere.

Chemists monitor these changes so that further damage can be limited:

Describe the composition and layered structure of the atmosphere:

– The atmosphere is a layer of gas, 200-300 km thick, that surrounds the Earth.

– COMPOSITION:

The atmosphere’s main composition of gases is:

78% nitrogen.

21% oxygen.

0.9% argon.

Other gases that make up an extremely small percentage of the atmosphere are carbon dioxide,

oxides of nitrogen (NOx), sulfur compounds (SO2 and H2S), carbon monoxide, neon, helium,

hydrogen, ammonia, ozone (O3) and volatile organic compounds, such as methane.

The atmospheric levels of these less abundant gases is measured in the unit called ‘ppm’

(parts per million).

The levels of water vapour in the atmosphere varies greatly according to weather, location, and

other factors, and hence is often disregarded.

– STRUCTURE:

The atmosphere is divided into layers, each with its own characteristics.

The first layer is the troposphere, followed by the stratosphere, then the mesosphere,

thermosphere, ionosphere and then finally the exosphere.

For this course, we will focus our atmospheric study on the troposphere (called the ‘lower

atmosphere’) and the stratosphere.

Troposphere:

Extends up from the surface of the Earth, to 15 km high.

At about 15 km high is the tropopause; it marks the end of the troposphere, and is the

beginning of the stratosphere.

The troposphere is the region where ‘weather’ is experienced.

Temperature Profile:

Temperature decreases as altitude increases (as you go up).

Stratosphere:

Extends from 15 km (the tropopause) to 50 km.

At about 50 km high is the stratopause, where the mesosphere begins.

The stratosphere has no ‘weather’.

The ozone layer is in the stratosphere.

Temperature Profile:

Temperature increase as altitude increases.

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

– POLLUTATANTS:

Carbon monoxide: From motor cars, bush fires, cigarettes, gas stove-tops.

Oxides of nitrogen (NO + NO2): From vehicles and power stations.

Sulfur dioxide: combustion and metal extraction (from sulfide ores).

Hydrocarbons: From vehicles and industrial solvents.

VOCs (volatile organic compounds): industrial plants, domestic solvents.

Particulates (soot, asbestos, etc.): combustion, mining, bushfires.

Airborne lead: lead smelters and lead-based paint from old houses.

– GAS MIXING:

Hot air rises and cold air falls.

For the troposphere, its bottom is warmer than its top (because temperature decreases as altitude

increases); hence, the air from the bottom is always rising to the top. That is, there is constant

mixing of gases (convection).

This means that pollutants released at the Earth’s surface by human activity are rapidly

dispersed throughout the stratosphere.

This is what causes the lower atmosphere to be so rapidly polluted.

For the stratosphere, its bottom is colder than its top (because temperature increases as altitude

increases); hence, there is little movement of gases. The cold air at the bottom and hot air at the

top simply stays there.

This ‘protects’ the stratosphere from pollutants released at the surface, as gas mixing stops at

the tropopause. The only way pollutants enter the stratosphere is by slow diffusion of gases.

Describe the formation of a coordinate covalent bond:

– RECALL:

The valence shell is the outermost electron shell of an atom.

Chemical reactions occur because of the activity of valence electrons; for example, if valence

electrons are transferred from one atom to another, this is a chemical reaction resulting in the

formation of an ionic bond.

To achieve chemical stability, all atoms seek to have an octet (8) of electrons in their valence

(outermost) shells.

– In a normal covalent bond, two atoms SHARE electrons in order for both to have a complete shell of

eight electrons:

EG: 2 oxygen atoms can share electrons to form a covalent bond and make an oxygen molecule

(O2). Notice that each oxygen contributes 2 electrons to the covalent bond (the ‘bubble’):

TIP - To count the electrons that belong to an atom, count BOTH electrons in bonds (all of

them), as well as the electrons not in bonds.

For example, the left oxygen atom in the O2 molecule as 4 electrons not in a bond, as well as

4 electrons in a bond. Hence it has a total of 8 electrons, which makes it a stable atom.

– However, in a coordinate covalent bond between two atoms, one atom will donate a pair of electrons

to the other to form a bond, so both can have an octet:

Further explained in the next dot-point.

Demonstrate the formation of coordinate covalent bonds using Lewis electron-dot

structures:

– RECALL:

In LEWIS electron-dot structures, atoms are represented by their chemical symbol, surrounded

by only valence electrons; e.g. oxygen will be represented as an ‘O’ surrounded by 6 electrons

(see ABOVE).

Use different symbols for the electrons of different atoms (e.g. crosses/dots).

– A coordinate-covalent bond involves one atom donating a pair of electrons to the other to form a

bond, so both can have an octet; all the electrons in the bond come from one atom:

EG: In the OZONE (O3) molecule, one of the oxygen atoms forms a coordinate covalent bond

with an oxygen atom:

TIP – When asked to draw a Lewis diagram of ozone, check that all oxygen atoms have only 6

electrons of their own, but 8 electrons bonded altogether:

Left : 4 unbonded electrons + 4 electrons in covalent bond = 8

Middle : 2 unbonded + 4 in covalent bond + 2 in coordinate covalent = 8

Right : 6 unbonded + 2 in coordinate covalent = 8

Compare the properties of the oxygen allotropes O2 and O3 and accounts for them on the

basis of molecular structure and bonding:

– RECALL:

Allotropes are different structural forms of the same element.

For example, allotropes of carbon are graphite and diamond.

Allotropes can exhibit very different physical AND chemical properties.

– Oxygen has 2 stable allotropes; oxygen gas (O2) and ozone gas (O3).

– These two allotropes exhibit different many different properties, and these can be explained in terms

of molecular structure and bonding:

Colour: O2 is colourless, while O3 is a pale blue gas:

No explanation available.

Boiling Point: The boiling point of O2 is –183°C, while O3 has one of –111°C:

The boiling point of O2 is lower than that of ozone as O2 has a lower molecular mass, hence

requiring less energy in the boiling process.

Solubility in Water: O2 has very low solubility in water compared to O3:

Non-polar O2 does not form strong intermolecular forces in the polar water. Ozone has a bent

structure, which creates a slight polarity in the molecule, allowing it to have intermolecular

interactions with water.

Chemical Stability: Ozone is much less stable than O2:

To decompose oxygen, its double bond has to be broken; this requires considerable amounts

of energy. However, the single bond (coordinate covalent bond) in ozone requires much less

energy to be broken, and hence ozone is much less stable (readily decomposes to O2).

Oxidising Strength: O2 is a moderately strong oxidising agent, while O3 is an extremely strong

oxidising agent:

The oxidising strength of ozone comes from the weakness of the single bond; it easily

releases an oxygen which can then oxidise a compound.

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

– Free radicals are atomic or molecular species with unpaired electrons.

They are NOT ions.

– EG: The oxygen free radical:

The oxygen free radical has two pairs of electrons, as well as two unpaired electrons (see

diagrams); these unpaired electrons are highly reactive.

It is basically an oxygen atom; they have the same electron configuration.

– The oxygen free radical can be made either by passing electrical

current through oxygen gas to decompose it, or by exposure to U.V.

radiation:

O2 2O·

– The oxygen radical is very short-lived, and will instantly react with other radicals, and so it cannot

be compared to oxygen gas in terms of physical properties such as boiling point or colour.

– However, a comparison can be made in terms of reactivity:

The oxygen free radical is much more reactive than oxygen gas.

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

and a lower atmosphere pollutant:

– The action of ozone (as well as its value) depends greatly on where it is located.

– In the upper atmosphere (stratosphere) ozone is found as the ‘ozone layer’. It is extremely crucial

part to life on Earth.

– However, in the lower atmosphere (troposphere), ozone is a serious air pollutant.

– Ozone in the Upper Atmosphere:

Ozone in the stratosphere, in the form of an ozone layer, protects us from harmful ultraviolet

radiation (UV light):

There are 3 forms of UV light; UV-A, UV-B and UV-C.

The ozone-layer blocks the harmful UV-B and UV-C rays from passing through the

atmosphere; these can cause many cancers and severe sunburn.

The useful UV-A (needed for photosynthesis) can still pass through.

The following equations show how ozone is formed and destroyed as well as how it protects us

from harmful UV radiation:

O2 + UV radiation 2O·

O· + O2 O3

O3 + UV radiation O· + O2

O· + O3 2O2

Every time an oxygen/ozone reacts with UV light, it absorbs it.

Hence ozone can be said to be an upper atmosphere UV radiation shield.

– Ozone in the Lower Atmosphere:

However, when ozone is found in the lower atmosphere (troposphere), it is considered a serious

air pollutant.

This is because ozone can cause serious health and environmental problems.

HEALTH ISSUES:

Ozone is poisonous to humans; as a strong oxidant, it can react with body tissue, especially

with sensitive mucous membranes when breathed in.

Ozone causes breathing difficulties, aggravates respiratory problems and produces headaches

and premature fatigue.

PHOTOCHEMICAL SMOG:

When ozone is found as a component of smog it is a serious pollutant, and can be very

hazardous for health.

Sunlight splits nitrogen dioxide and the free radical produced joins with oxygen to form

ozone: NO2 NO + O·. Also, hydrocarbons and PAN (peroxyacyl nitrates) can be

present in the smog.

– Thus ozone is both as a major pollutant, and a UV radiation shield.

Identify and name examples of isomers (excluding geometrical and optical) of haloalkanes

up to eight carbon atoms:

– Haloalkanes are compounds formed when one of the hydrogens of an alkane is replaced by a

halogen atom (F, Cl, Br or I).

– Naming Haloalkanes:

There is a systematic IUPAC method of naming haloalkanes.

EG: Name this haloalkane:

Firstly, count the number carbons in the longest carbon chain; in this case, there are 3, so

the parent alkane is PROPANE.

Next identify, name and number the halogens, using ‘fluoro-’, ‘chloro-’, ‘bromo-’ and

‘iodo-’ as prefixes instead of chemical names. Say we take the left-most carbon as carbon-

1 (C1); then, there are 3 chlorines (2 on C1 and 1 on C3), and there are 3 fluorines (1 on C1

and 2 on C3). For multiple halogens, use the prefixes di-, tri-, and tetra-. Hence, so far we

have propane, 1,1,3-trichloro- and 1,3,3-trifluoro.

To make the full name, place the parent alkane at the end, and place the halogens in

alphabetical order: 1,1,3-trichloro-1,3,3-trifluoropropane.

Now CHECK; say instead we took the right-most carbon as C1. In that case, the name

would be 1,3,3-trichloro-1,1,3-trifluoropropane. You MUST take the name with the lower

sum of numbers.

In this case, both names have an equal sum. If this occurs, you give the lower numbers to

the more electronegative halogen (F > Cl > Br > I). Hence, the correct name is 1,3,3-

trichloro-1,1,3-trifluoropropane.

– Isomers :

Isomers are compounds that have the same chemical formula, but different structural formula.

Carbon compounds, such as haloalkanes are able to form many different isomers; the longer the

chain, the more isomers are possible.

EG: Draw 4 isomers of C4H7ClF2 and name them using IUPAC nomenclature:

2-chloro-1,3-difluorobutane.



1-chloro-1,1-difluoro-2-methylpropane.

Identify the origins of chlorofluorocarbons (CFCs) and halons in the atmosphere:

– CFCs:

Chlorofluorocarbons are compounds that contain ONLY carbon, fluorine and chlorine; NO

hydrogen atoms (e.g. CFC-12 is CCl2F2).

CFCs were introduced in the 1930’s (known as freons) as replacements for ammonia in

refrigeration:

This was because had the required pressure-dependent properties that refrigerants needed, as

well as that they were odourless, non-flammable, non-toxic and inert, much unlike the toxic,

foul-smelling ammonia.

CFCs were very widely used as:

Refrigerants in fridges and air-conditioners.

Propellants in aerosol spray cans.

Foaming agents in the manufacture of foam plastics like polystyrene.

Cleaning agents in electronic circuitry.

These many uses released CFCs directly into the lower atmosphere.

As they were very inert and insoluble in water (rain), they remained in the troposphere, where air

convection spread them throughout the atmosphere.

Very slowly, the CFC’s began to diffuse through the tropopause and into the stratosphere, where

problems began to occur.

– Halons:

Halons are compounds that contain carbon and bromine, as well as other halogens; they are

basically CFC’s that also contain bromine (no hydrogen).

They are dense, non-flammable liquids that were widely used as effective fire-extinguishers

(called BCF fire-extinguishers).

As they were used onto fires, the halons were released directly into the atmosphere, where they

too slowly diffused into the stratosphere.

Discuss the problems associated with the use of CFCs and assess the effectiveness of steps

taken to alleviate these problems:

– The biggest problem associated with the use of CFCs is the destruction of stratospheric ozone (i.e.

depletion of the ozone layer):

As stated above, CFCs are very inert, and are not washed out by rain.

As a result, CFCs remain in the troposphere for many years, and eventually diffuse into the

stratosphere, where they deplete the ozone layer.

This leads to more UV radiation reaching Earth, which greatly increases the chances of

mutations and damage (especially cancer) in living things.

– How CFCs Destroy The Ozone Layer:

Firstly, short wavelength UV radiation (that has not been removed by the ozone layer) attacks the

CFC molecule and breaks off a chlorine atom:


This chlorine atom (radical) then reacts with ozone, forming oxygen gas, and a chlorine

monoxide radical (ClO):

·Cl (g) + O3 (g) ·ClO (g) + O2 (g)

The chlorine monoxide radical then reacts with an oxygen radical, and the chlorine radical is

regenerated:

·ClO (g) + ·O (g) ·Cl (g) + O2 (g)

The net result is that an ozone molecule and an oxygen radical have been converted into 2

oxygen molecules AND the chlorine has not been used up.

The chlorine radical can then attack another ozone molecule and repeat the whole process

thousands of times; this is a chain reaction.

A very small amount of CFC in the stratosphere can do significant damage.

– How Chlorine Radicals Are Removed:

According to the above, theoretically, a single chlorine atom could destroy the entire ozone-

layer; however certain natural reactions remove this radical.

The chlorine atom reacts with stratospheric methane, ending the chain reaction:

·Cl (g) + CH4 (g) HCl (g) + ·CH3 (g)

Neither HCl nor the methyl-radical has any effect on ozone.

Another important reaction for stopping ozone depletion involves the ·ClO species reacting with

nitrogen dioxide, forming chlorine nitrate:


– The Antarctic Spring Ozone Hole:

There is a serious periodic depletion of the ozone-layer that occurs every spring over Antarctica;

it is called the ‘Ozone Hole’.

This is due to the conditions of Antarctica in winter, as well as spring.

Antarctic winters are perpetually dark; the cold conditions, as well as solid particulate catalysts

in the air, encourage the following reaction to occur:


This has zero effect on ozone levels during the winter.

However during early spring, the Sun begins to rise, and the situation changes dramatically;

sunlight is able to split chlorine molecules:


Hence in spring there is another source of chlorine radicals to destroy more ozone; the

concentration of ozone is reduced dramatically, causing a hole.

Eventually, the fixed amount of Cl2 created over the winter is blown away by Antarctic winds

and the ozone layer slowly regenerates.

– Dealing With The CFC Problem:

The only way to stop ozone depletion is to STOP releasing CFCs of any form.

INTERNATIONAL AGREEMENTS:

The main way the CFC problem is being dealt with is by international agreements based on

the common goal of phasing out CFCs.

The Montreal Protocol on Substances That Deplete the Ozone Layer (1987) is an

international treaty designed to protect the ozone layer by phasing out the production of a

number of substances believed to be responsible for ozone depletion.

Its goals include ceasing the manufacturing and banning the use of CFCs and certain

haloalkanes by 1996, the end of halon use by 1994, the phasing out of HCFCs, as well as the

provision of financial assistance to developing nations in order to help them reach the goals

of Montreal.

CFC REPLACEMENTS:

Finding alternative compounds to fulfil the roles of CFCs is a major step forward in

preventing ozone depletion.

This is examined in greater detail below.

Dealing With Increased UV Radiation:

Increasing UV levels have meant that more UV inhibitors need to be included in polymers

(such as PVC) and Cancer Councils have advised the use of only sunscreens with a rating of

SPF 30+ or greater.

– Effectiveness of These Solutions:

The Montreal Protocol is only effective if member nations ratify the protocol and adhere to its

regulations; so far, the Montreal Protocol has been a huge success in international agreement and

environmental health.

Certain CFC replacements are not as effective as the CFCs themselves; future technological

advancement hopes to find better replacements.

There are still, however, significant levels of CFCs in the atmosphere, and current technology

has no way of removing them.

– The Greenhouse Problem:

A less talked about issue of CFCs and HCFCs is that they contribute greatly to the increased

greenhouse effect (thousands of times more than CO2).

This may lead to climate change, and is the focus of the Kyoto Protocol.

Present information from secondary sources to write the equations to show reactions

involving CFCs and ozone to demonstrate the removal of ozone from the atmosphere:

– All the CFC-related equations from above:

Formation of chlorine radical:


Reaction of ozone:

·Cl (g) + O3 (g) ·ClO (g) + O2 (g)

Regeneration of chlorine:

·ClO (g) + O (g) ·Cl (g) + O2 (g)

Removal of chlorine radical:

·Cl (g) + CH4 (g) HCl (g) + ·CH3 (g)

Removal of chlorine monoxide radical:


Formation of molecular chlorine:


Decomposition of molecular chlorine:


Present information from secondary sources to identify alternative chemicals used to replace

CFCs and evaluate the effectiveness of their use as a replacement for CFCs:

– Ammonia:

Large scale (industrial) refrigeration has reverted back to using ammonia as a refrigerant, as was

done prior to the discovery of CFCs.

However, great care is exercised, as ammonia is dangerous and toxic.

– HCFCs:

Hydrochlorofluorocarbons are CFCs that contain hydrogen.

These were the first replacements for CFCs.

HCFCs contain C–H bonds that are susceptible to attack by reactive radicals in the troposphere

and so are decomposed rapidly to a significant extent.

This means that only a very small proportion ever reaches the stratosphere.

HCFCs replaced CFCs in domestic refrigeration, as propellants in spray cans, as an industrial

solvent and as a foaming agent.

Effectiveness :

Small amounts of HCFCs do reach the stratosphere, and hence they are also ozone-depleting

(10% the ozone-depleting potential of CFCs).

They are seen as only a temporary solution.

HCFCs also contribute massively to the greenhouse effect, and so their use is being phased

out (complete ban by 2030).

– HFCs:

Hydrofluorocarbons are compounds that contain only carbon, hydrogen and fluorine (NO

chlorine or bromine).

They are widely seen as a viable CFC and HCFC alternative, as they contain reactive C–H bonds

(so they degrade in troposphere) as well as the fact that they do not contain any chlorine (and

hence cannot form ·Cl radicals).

Their ozone depleting potential is zero.

HFCs are very widely used in refrigeration and air-conditioning applications.

Effectiveness :

As they have zero ozone-depleting potential, HFCs are a good alternative to using CFCs in

terms of atmospheric health.

However, they are not as effective refrigerants as CFCs, and are slightly more expensive.

They are also strong greenhouse gases, and so further research is required.

Analyse the information available that indicates changes in atmospheric ozone

concentrations, describe the changes observed and explain how this information was

obtained:

– How Are Ozone Levels Measured?

Stratospheric ozone levels are measured from ground-based instruments, from instruments in

satellites and from instruments in weather-balloons.

Ozone levels are measured in Dobson Units (DU).

The measurements made indicate that changes in ozone levels have occurred.

GROUND-BASED:

The ground-based instruments used are UV spectrophotometers that are able to measure the

intensity of light at specific wavelengths.

These instruments are pointed directly upwards towards the sky and are set to measure light

intensity at the wavelengths of light at which ozone absorbs (such as UV-B and UV-C) and at

wavelengths on either side (for light which ozone does not absorb).

A comparison of these 2 measurements gives a measure of the total ozone in the atmosphere

per unit of area of Earth surface at that location.

BALLOON-BASED:

These spectrophotometers can also be placed in high-altitude weather balloons that can rise

above the stratosphere.

The instruments are pointed downwards, and measure ozone from above.

SATELLITE-BASED:

An instrument called the TOMS (total ozone mapping spectrophotometers) have been placed

on several US satellites.

They work similarly to the UV spectrophotometers as above, but as the satellites orbit the

Earth, the TOMS is able to scan the entire globe and measure ozone concentrations as a

function of altitude and geographical location.

– The Changes Observed:

Measurements of the total amount of ozone in a column of atmosphere have been recorded since

1957.

The main depletion of ozone has occurred over the Antarctic.

Scientists identified that a dramatic decline in springtime ozone occurred from the late 1970s

over the entire Antarctic. The decline reached approximately 30% by 1985. In some places, the

ozone layer had been completely destroyed.

The ozone decline over Antarctica during springtime is now not so dramatic, but often exceeds

50%.

PRACTICAL – Gather, process and present information from secondary sources including

simulations, molecular modelling kits or pictorial representations to model isomers of

haloalkanes:

– In this experiment, molecular modelling kits were used to show different isomers of a haloalkane.

– The class was divided into groups, and each group was provided with a kit.

– Each group was provided with a haloalkane, which they were require to draw structural formula for,

and then using the kit, 3 different isomers were formed.


The models created a good visual 3D representation of a chemical property; that is, isomerism.

– RESULTS:

Four isomers of C4H7ClF2 with their IUPAC names:




1-chloro-1,1-difluoro-2-methylpropane.

5. Human activity also impacts on waterways. Chemical monitoring and management

assists in providing safe water for human use and to protect the habitats of other

organisms:

Identify that water quality can be determined by considering:-

concentrations of common ions:

total dissolved solids:

hardness:

turbidity:

acidity:

dissolved oxygen and biochemical oxygen demand:

– Concentration of Common Ions:

Usually, the concentration of all the ions in a solution is considered as a whole; this is called the

total dissolved solids.

However, certain common ions have their concentrations individually measured; the results of

these tests indicate different aspects of water quality.

Important ions whose concentrations are individually measured include metal cations such as

sodium (Na+), magnesium (Mg2+), calcium (Ca2+) and anions such as chloride (Cl-) and the

polyatomic phosphate (PO43-) and nitrate (NO3

-).

The concentrations of sodium and chloride are important indicators of the salinity of water

system; any significant change in salinity (whether increase or decrease) can greatly affect any

aquatic life.

Magnesium and calcium ions are measured to indicate water hardness, which is covered in more

detail below.

Phosphate and nitrate ions are essential to aquatic life; however, excess levels of these ions leads

to eutrophication and algal blooms, destroying waterways.

MEASURING IONS:

The concentration of the metal cations in samples is very quickly and easily measured using

spectroscopic methods such as AAS.

Chloride levels can be measured by titrating the sample against silver nitrate, with potassium

chromate as the indicator.

The methods used to determine phosphate and nitrate levels in water samples are covered

below (see ‘Monitoring Eutrophication’).

– Total Dissolved Solids:

TDS (total dissolved solids) is the total mass of all solids dissolved in a given volume of water. It

is given as either mg/L or ppm (both mean the same thing).

The dissolved solids are mainly composed of salts (ionic compounds).

TDS is related to the quality of water because clean water is relatively free of contaminants; a

large amount of dissolved solids indicates unclean water.

MEASURING TDS:

A gravimetric method for determining TDS is by first filtering the water sample to obtain a

clear solution, then evaporating the sample dryness, and weighing the solids. This method is

very laborious and prone to inaccuracies, and hence electrochemical methods are preferred.

Because most of the dissolved solids are ions, the electrical conductivity of the water sample

can be used to measure TDS. Although conductivity tests give a measure of total dissolved

salts, this is quite an accurate approximation for TDS.

– Hardness:

Hard-water is water with high levels of calcium and magnesium ions.

Hardness is an issue for water quality because hard-water creates ‘scale’ deposits of CaCO3 and

MgCO3 on sinks and bathtubs. Also soap does not work in hard water, and forms scum-

precipitates which can stain fabrics.

MEASURING HARDNESS:

The levels of these ions can be measured by titration against a compound known as EDTA

(ethylene-diamine-tetra-acetic-acid).

Hardness is then expressed as mg/L of CaCO3.

Alternatively, the levels of Mg2+ and Ca2+ can be measured using AAS.

– Turbidity:

Water turbidity means a ‘cloudiness’ of the water.

Turbidity is caused by the presence of suspended solids that are sufficiently small so that they do

not settle upon standing and remain suspended.

Turbidity gives water an undesirable appearance as well as a unpleasant taste.

High levels of turbidity can affect the penetration of light through the water, which then reduces

the rate of photosynthesis, reducing dissolved oxygen.

MEASURING TURBIDITY:

Turbidity is measured using a turbidity tube.

A turbidity tube is a long hollow plastic cylinder that has a mark (usually a black cross)

inscribed onto its flat bottom. Water is poured into the tube until the mark is no longer

visible. The height of water at which this occurs is the turbidity measurement of the water.

Turbidity is measured in NTU (nephelometric turbidity units).

– Acidity:

The pH of a water system is a good indicator of its health.

Anything outside the normal range of 6.5-8.5 indicates a polluted system, caused by the

discharge of some chemical into the water.

Great changes in acidity/alkalinity greatly affect the usability of water.

MEASURING pH:

The pH of water is measured using indicators, pH strips or a pH meter.

– Dissolved Oxygen:

DO (dissolved oxygen) refers to the levels of molecular oxygen (O2) that are dissolved in a water

sample.

Even though O2 has very low solubility in water, this small amount is crucial to aquatic life; this

oxygen is extracted and used for respiration.

Low oxygen levels (lower than 5 ppm) will cause many aquatic species to die. The water body

will decay and develop undesirable characteristics.

MEASURING DO:

Dissolved oxygen can be measured using a chemical titration known as the Winkler Method.

In this method, the dissolved oxygen oxidises manganese from Mn2+ to Mn(IV) in alkaline

solution. Mn(IV) then oxidises iodide (I-) ions to iodine (I2) in acidic conditions. Iodine is

then titrated against sodium thiosulfate (Na2S2O3) using starch as in indicator.

Alternatively, an electronic oxygen sensor can be used.

– Biological Oxygen Demand:

The BOD (biological oxygen demand) of a water body is a measure of the concentration of

dissolved oxygen that is needed for the complete breakdown of the organic matter in the water

by anaerobic bacteria.

That is, it measures how fast bacteria (or other organisms) use up oxygen.

High BOD levels indicate large amounts of organic matter in the water, which will drain the

water of oxygen; sewage has very high levels of BOD.

MEASURING BOD:

Two samples of water must be taken; the dissolved oxygen is measured in one of the samples

immediately, while the other sample is placed in a sealed air-free container and incubated at

20ºC for five days.

After 5-day incubation, the second sample’s oxygen levels are measured and the difference in

dissolved oxygen levels is the BOD.

Identify factors that affect the concentrations of a range of ions in solution in natural bodies

of water such as rivers and oceans:

– The concentrations of a range of ions is bodies of water (such as rivers and oceans) is affected by a

number of factors, both natural and unnatural:

It can be said that the source of water for all water-bodies is rain.

Rain contains very few ions (only very little Na+ and Cl- from sea-spray) and so any increased

levels of ions in water bodies is a result of what occurs in between rainfall and when the water

flows into the water-body.

NOTE that the ion levels in rivers and lakes are much susceptible to change compared to oceans,

which have much greater volumes of water.

– Natural Sources of Ions:

If rain falls on bushland, and then runs-off into streams and rivers, it will only pick up small

amounts of nitrates and phosphates from surface nutrients, as well as some Mg2+ and Ca2+ from

minerals; TDS will be low (≤50).

If rain soaks into the ground and flows into aquifers (layers of permeable rock) and then into

rivers, the water will have increased levels of Ca2+, Mg2+, Cl-, SO42- and CO3

2- which are

dissolved from the soils and rocks the water flows through. TDS will be moderate (≈200).

If water seeps down into rocks (percolates) in deep underground basins (artesian basins) and

only reaches the surface centuries later, then the levels of the above ions will increase massively,

as may contain other ions (such as Fe3+, Mn2+, Cu2+ and Zn2+. TDS values are very high (>1000).

– Land Clearing:

When land is cleared of vegetation, the soil loses the stabilising effects of plant roots and so soil

is easily moved and displaced.

If rain flows over cleared land, it will disturb dirt and sediment and carry it into the water body it

flows into (this will greatly increase TURBIDITY).

The level of dissolved solids and ions (TDS) will increase, due to high levels of mineral ions in

the soils (such as Na+, K+ etc.)

– Agriculture:

When rain flows over land used for the growing of crops and pasture-land it leads to increased

levels of phosphates and nitrates due to fertilizer run-off.

The turbidity will also increase, as well as the BOD, as organic matter (such as animal faeces)

enters the water-ways.

– Raw Sewage and Other Effluents:

Raw sewage, if discharged directly into the water, will greatly increase the levels of many ions

(especially nutrient ions such as nitrates and phosphates).

Sewage can increase a water-way’s TDS by 200 or more; it also greatly increases turbidity, BOD

and pathogen levels.

Stormwater run-off in urban areas can also carry high levels of ions.

Industrial effluent as well as leaching from rubbish dumps can increase the levels of dangerous

HEAVY-METAL ions in water (such as Hg2+ and Cd2+).

Describe the design and composition of microscopic membrane filters and explain how they

purify contaminated water:

– A membrane filter is basically a thin film of synthetic polymer throughout which there are small

pores of uniform size.

– Common polymers used are polypropylene and polytetrafluoroethylene (PTFE).

– Filters are classified according to the size of their pores, as this determines what type of particles can

pass through.

– FILTER DESIGN:

There are many designs in which the filter can be formed; this depends on the purpose for which

the filter was made.

One simple design involves a thin sheet of filter-membrane that is folded around a hollow core

(such as a thin tube) and surrounded by a casing.

When water is passed over the membrane, clean water passes through and exits via the hollow

core.

Another design involves forming the membranes into capillaries (tiny tubes) with a diameter of

about 500 μm; these are called hollow fibre filters.

Large numbers of these capillaries are bounded together to form a filtering unit that has a very

large surface area.

Water is passed over the surface of the capillary and clean water passes through to the middle of

the capillary and flows out.

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

supplies:

– Monitoring Catchment:

The first step to ensure water used for human use is clean is to ensure that the area the water

flows over (the CATCHMENT area) is kept clean.

This involves banning any land-clearing, industry or agriculture in the entire catchment area, to

prevent sediment, animal waste or bacteria to build up in water supplies.

EFFECTIVENESS:

This is a very cheap and effective way of ensuring the purity of water for human use; by

removing the sources of contamination, purity is ensured.

– Screening:

Before the water from catchment areas is allowed to enter treatment plants or storage dams, it is

passed through metal screens that remove large debris such as sticks, leaves, trash and other

large particles which may interfere with subsequent purification steps

EFFECTIVENESS:

This step is effective for its purpose, but more treatment is needed.

– Clarification and Flocculation:

Certain suspended particles (called colloidal particles) cause water to become turbid, but are too

small to be removed by conventional filtration.

In this step, these particles are coagulated together to form large particles, which can then be

removed; this is called ‘flocculation’.

Firstly, the pH of the water is increased, as this encourages the formation of precipitates; this is

achieved by adding lime.

Next, an electrolyte is added to force the particles to precipitate; iron(III) hydroxide or

aluminium chloride is used (although Al3+ ions are preferred as they leave no metallic aftertaste

in the water, as Fe3+ does).

The precipitate is first formed as very tiny particles, but as the water is gently agitated or stirred,

the particles flocculate into larger particles.

EFFECTIVENESS:

Flocculation removes most of the suspended particles, as well as bacteria, which are caught

up in the particle aggregates. It is very cost-effective, and relatively fast.

– Sedimentation:

Water that has been flocculated is then allowed to settle in large tanks; this causes the dirt and

other particles to fall to the bottom of the tank as a sludge, where it is removed.

The clear water is then pumped to the filtration systems.

EFFECTIVENESS:

Sedimentation employs a natural force (i.e. gravity) to separate the sludge from clean water,

and so reduces plant running costs.

The slow speed of sedimentation may affect its effectiveness.

– Filtration:

Water from the settling tank is then passed through a filter bed of fine sand and gravel; this

removes the rest of the particulate matter, as well as the material that did not settle to the bottom

of the tank.

Sometimes anthracite (metamorphic coal) is added to the filter beds, as it adsorbs organic matter

and removes odours.

EFFECTIVENESS:

Sand filtration removes a high proportion of the particulate matter that aggregated during

flocculation, however extremely small particles are not removed (such as some bacteria and

viruses).

It is suitable for providing water to urbanised areas.

– Chlorination:

Lastly, before water is pumped to homes, chlorine gas is bubbled through the water; this forms

the hypochlorite ion (OCl-) which can kill disease causing agents, such as bacteria and some

viruses.

EFFECTIVENESS:

Chlorination is an effective way of removing most pathogenic organisms, but, it is not so

effective at killing viruses

Also, chlorine may impart an unpleasant odour on the water.

– Alternative - Membrane Filters:

All the above are the steps taken to sanitize water in most Australian plants.

Alternatively, membrane filters can be used, which would replace many of the steps used above.

Membrane filters would remove the need for flocculation, sedimentation, sand filtration and also

chlorination, because:

Their pore size is sufficiently small to remove suspended particles.

Membranes are very thin and allow the use of pressurization; this greatly speeds up the

process, and hence removes the need for sedimentation.

Sand filtration is a slow process that membrane filtration supersedes.

Membrane filters can remove all bacteria AND viruses, as these organisms are much too large to

fit through the pores of a membrane filter.

– However, membrane filters are considerably more expensive than current methods used, and so their

use is limited by costs.

Gather, process and present information on the range and chemistry of the tests used to:

identify heavy metal pollution of water

monitor possible eutrophication of waterways

– Heavy Metals:

Heavy metals are the transition metals in addition to lead and arsenic.

The heavy metals that are of the most concern due to their extremely detrimental effects on

health (and hence should be monitored for the most) are mercury, lead, cadmium, chromium and

arsenic.

The levels of these heavy metals is most easily quantitatively measured by using AAS, as

their levels are usually very low.

One of the most common quantitative tests for the presence of heavy metals is the sulfide-test:

A water sample is acidified, and then a few drops of sodium sulfide (Na2S) is added. If a

precipitate forms, then one of the following ions is present:

Lead, silver, mercury, copper, cadmium or arsenic.

If no precipitate forms in acidified conditions, then the sample is made alkaline. If this

produces a precipitate, then one of the following is present:

Chromium, zinc, iron, nickel, cobalt, manganese or aluminium.

CHEMISTRY:

The sulfide-test is based on the following equilibrium:

S2- (aq) + 2H3O+

(aq) H2S (aq) + 2H2O (l)

When in acidic solution, equilibrium lies greatly to the right, and there is only a small amount

of sulfide; this small amount is enough to precipitate the first group of heavy metals, but not

the second.

In alkaline solution, the equilibrium lies to the left; there are large amounts of sulfide, and

hence the 2nd group of heavy metals precipitates (as well as the 1st group; this is why the

acidic test is done first).

Other simple quantitative techniques to identify the presence of heavy metals includes other

precipitation tests, flame tests (e.g. for copper).

Other qualitative tests (besides AAS) includes volumetric and gravimetric analyses, colorimetry

(using colorimeter devices) and chromatography.

– Eutrophication:

Eutrophication is the process by which a water body becomes enriched with nutrients (SO42- &

NO32-) to such an extent that an algal bloom is very likely.

Eutrophication is not the same as algal blooms, but LEADS to algal blooms.

Waterways need to be monitored for possible eutrophication in order to stop algal blooms from

occurring.

Algal blooms need to be prevented because:

Blue-green algae (cyanobacteria) in algal blooms produce poisons that can kill humans as

well as livestock.

Water becomes unsuitable for normal recreational uses as it is clogged up with algae, which

can cover rivers for kilometres

Algae can starve rivers of oxygen, killing all the aquatic life.

To monitor waterways for eutrophication, the levels of nitrates and phosphates needs to be

measured (quantitative testing only).

NITRATES:

Methods to measure the nitrate levels in water usually involve measuring the total nitrogen

levels in the sample.

Two methods are used: Kjeldahl digestion, and colorimetric methods.

The chemistry of Kjeldahl digestion:

The sample is heated with concentrated sulfuric acid to ‘digest’ any organic nitrogen

compounds into ammonium sulfate.

This is then reacted with sodium hydroxide to form ammonia.

The levels of ammonia are then measured by back titration against a standardised

solution of hydrochloric acid.

The chemistry of the colorimetric method:

In this method, the sample is completely digested, as above, except the sample is then

treated with Nessler’s reagent, a chemical that reacts with nitrogenous compounds to

form a yellow compound.

The colorimeter then measures the intensity of the yellow colour to calculate the

concentration of nitrates.

If the levels of nitrate reaches too high, the water is called eutrophic.

PHOSPHATES:

As above, techniques to measure phosphates actually measure the levels of total phosphorus

present in the water sample.

The main technique used is a colorimetric method.

The chemistry of the colorimetric method:

In this method, a measured quantity of ammonium molybdate is added to the sample and

completely dissolved.

A measured quantity of solid (powdered) ascorbic acid is then added, and this forms an

intensely blue complex of a compound know as ‘molybdenum blue’.

The intensity of the blueness is measured by a colorimeter and this places a value on the

concentration of phosphate.

If phosphate levels reach greater than 0.05 ppm (or mg/L), than the water is eutrophic and an

algal bloom is extremely likely.

Gather, process and present information on the features of the local town water supply in

terms of:-

catchment area:

possible sources of contamination in this catchment:

chemical tests available to determine levels and types of contaminants:

physical and chemical processes used to purify water:

chemical additives in the water and the reasons for the presence of these additives:

– For the local Sydney area, the water is supplied by Warragamba Dam.

– CATCHMENT AREA:

Warragamba Dam is Sydney's main water storage dam, and one of the largest domestic water

supply dams in the world.

The dam lies on the Warragamba River, for which it is named.

Catchment areas are areas of land from which rain water drains toward a common water-body.

The Warragamba catchment has an area of about 9000 km2 and extends from south of Goulburn,

north to Lithgow, east to Werombi and Mittagong, and west to part of the Crookwell local

government area.

– SOURCES OF CONTAMINATION:

Land Clearing :

Within the catchment area, there are various logging and land-clearing activities occurring to

make way for more agricultural land.

This has lead to increased turbidity in the water flowing into the dam (especially during

heavy-flow periods).

Higher levels of dissolved solids in the water also occurs.

Agriculture :

There are various patches of land within the catchment area that are used for agricultural

purposes such as growing crops or raising cattle.

Run-off from agricultural land contains high levels of phosphates and nitrates as a result of

leeching fertilizers from the soil and crops.

Low-levels of pesticides (such as organochlorins and organophosphates) have been detected

in the run-off from agriculture.

Water that runs over land used for cattle-grazing (pasture-land) can carry with it animal

faeces (detected by high faecal coliform levels); this contamination leads to the growth of

bacteria, as well as a high BOD.

Mining :

Within the catchment area are abandoned mines which water can flow into.

When water leaks into metal mines and flows out, it leaches out with it certain ions such as

Zn2+ and Cu2+ which are heavy metal pollutants, as well as sulfides from the metal ores,

which make the water acidic.

Coal mining occurs in the outer catchment area; possible contamination of the water involves

disposal of mining wastes directly into water-ways.

Natural Soil :

The natural soil and rock strata around the catchment area have high levels of iron and

manganese in them.

Rain water can leach out these minerals, in the form of Fe3+ and Mn2+ ions.

This leads to water with a coloured tinge and a metallic taste.

Sewage :

At certain places along the sewage line are places called ‘overflow points’.

During times of heavy flooding the sewage treatment plants cannot handle the heavy input of

stormwater flows, and so RAW untreated sewage is allowed to flow out of these overflow

points.

This leads to contamination of the water with bacteria and excess ions.

Animals :

Certain feral and native animals may contaminate the water with their faeces directly, or by

dying and decaying within the water.

This can lead to serious contamination of the water with pathogens.

This was the cause of the Giardiasis and Cryptosporidium scare of 1998.

– TESTING FOR CONTAMINANTS:

The testing for contaminants for Warragamba Dam occurs in a variety of locations; using

multiple testing sites ensures that a wide range of information.

Most chemical tests are used to detect the levels common ions, and ensure that they remain

below safe thresholds; this is especially true for nutrient ions.

The chemical tests available to determine levels and types of contaminants have already been

described in greater detail (see above).

In summary:

The test for common ions uses electrical conductivity; the test for heavy metal ions is the

sulfide-test, but AAS is used to determine quantities; the test for nitrates is Kjeldahl

digestion; phosphates detection involves molybdenum colorimetry and BOD is measured

using the Winkler method.

– WATER PURIFICATION:

The purification of water that is caught in the Warragamba catchment area is purified using the

method described above.

It is a cost-effective method for sanitizing water to an acceptable degree.

In summary:

Water is first screened to remove large debris; flocculation then occurs using Fe3+ or Al3+

electrolytes, followed by shaking to encourage precipitate formation; this is then left in a

sedimentation tank to settle; sludge is scooped out and the clean water is led onto a sand-bed

filter; this filter consists of layers of sand and gravel and the water comes out clean.

– CHEMICAL ADDITIVES:

The 2 main chemical additives in Sydney water are chlorine and fluoride.

Chlorine (Cl2) is added to the water supply as a disinfecting agent; chlorine gas is bubbled

through the water just before it exits the plant. Hypochlorite ions are formed, and these kill

bacteria and some viruses, sanitizing the water.

Fluoride (F-) is added to the water because it is believed to strengthen tooth enamel in growing

children. Fluoride ions are added in the form of sodium fluoride at a controlled concentration of

1 ppm.

PRACTICAL – Perform first-hand investigations to use qualitative and quantitative tests to

analyse and compare the quality of water samples:

– In this practical, water samples were taken from the local Cook’s River and using a ‘StreamWatch’

kit the water quality tests were performed on-site.

– Water samples were taken from various parts of the river, and the results were compared to the same

tests performed on distilled water.

– Pollutants:

This was a qualitative test performed on the water samples to detect pollution.

Any sort of colouration of the water (any deviation from ‘clear’) indicated the presence of a form

of pollution; any ‘rainbow streaks’ meant hydrocarbons.

The water was smelt to detect any odours (another sign of pollution).

A sample bottle was filled half-way with sample water, and another bottle was filled with

distilled water; both were shaken at the same time for equal periods of time. If the bubbles

remained longer on the sample, this indicated the presence of detergents.

The presence of heavy metal pollution was performed in the laboratory; 1 mL of sodium sulfide

was placed in 5 mL of sample water. Any precipitate indicated the presence of a heavy metal.

– Turbidity:

This test was a relatively quantitative test, with turbidity measured in NTU.

Firstly, distilled water was poured into the turbidity tube, and it was clear all the way to the top

of the tube. It was given a turbidity of <10 NTU.

The sample water was collected in a bottle using a telescoping handle, and this water was poured

into the tube; it had a much higher turbidity, and clouded the view of the cross on the bottom of

the tube.

– Temperature:

A thermometer (with a rubber handle) was placed directly into the water way to accurately

measure the temperature of the water; it was found to be within acceptable ranges.

– pH:

pH of water samples was done using strips of pH paper.

Firstly, the pH of distilled water was measured and found to be 7.

The pH of the water sample was then measured and found to be 7.5.

– TDS:

The total dissolved solids was measured using an electrical conductivity meter that measured in

μS/cm (mircosieverts per centimetre).

The sample water had only slightly more ions that distilled water.

– DO/BOD:

The levels of dissolved oxygen was easily measured using an oxygen sensing metre found in the

StreamWatch kit.

Two sample bottles were filled; one was capped air-free and sealed tightly, and the DO level in

the other was measured immediately.

The other sample had the levels of DO measured 5 days later, and the difference was the BOD.

– Hardness:

For hardness, a titration method was used.

To a 100 mL sample, 1 mL of buffer solution (pH 10) was added, and 2-3 drops of Eriochrome

Black T indicator was added.

This red-violet solution was titrated against a solution of EDTA until a permanent blue colour

occurred.

A lack of definite change in colour meant there were no magnesium ions.

Yr 12 Notes (Other Half)

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