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9.2.1 Fossil fuels provide both energy and raw materials such as ethylene for the
production of other substances
Identify the industrial source of ethylene from the cracking of some of the fractions from the refining of petroleum
Industrial source: crude oil/ petroleum; either as by-product of petrol refining or ‘cracking’ of higher boiling point fractions. Cracking is the chemical process of breaking large hydrocarbons into smaller ones. Catalytic cracking is the process in which high molecular weight fractions from crude oil are catalytically broken into lower molecular weight substances (an alkane and alkene), like petrol, to increase output of these high demand products.
Identify that ethylene, because of the high reactivity of its double bond, is readily transformed into many useful products
Ethylene has a double bond since it’s an alkene, which makes it reactive. Ethylene undergoes an addition reaction in which the double bond opens out to form two single bonds thus linking molecules together, i.e. (CH2=CH2)n→ (-CH2 – CH2-)n. It forms
useful products such as ethanol as a result.
Identify that ethylene serves as a monomer from which polymers are made
Ethylene monomers polymerise to form low or high density polyethylene.
Low density polyethylene is made from the gas phase process – high pressure, high temperatures and an organic peroxide initiator.
High density polyethylene is made using Zeigler-Natta process – just above atmospheric pressure, temperatures around 60°C and zeolite catalyst.
Identify polyethylene as an addition polymer and explain the meaning of this term
Polyethylene is called an addition polymer. This means that it forms by molecules adding together without the loss of any atoms. Each double C=C bond opens out to form single bonds with adjacent molecules thus linking molecules together.
Outline the steps in the production of polyethylene as an example of a commercially and industrially important polymer
1. Catalyst (for HDPE) or initiator (for LDPE) attaches to ethylene molecule → creates activated species.
Z + CH2=CH2 → Z – CH2 – CH2
2. Ethylene molecules attach to the species, expanding the chain.
Double bond
Catalyst/ Initiator
Ethylene molecule
2
C6H5
Z – CH2 – CH2 + CH2=CH2 → Z – CH2 – CH2 – CH2 – CH2
3. Polymerisation stops when two activated species collide, forming a stable polymer. Z – (CH2 – CH2)x + Z – (CH2 – CH2)y → Z – (CH2 – CH2)x+y – Z
Note: Since activated chains of various lengths can collide, polymer molecules formed have different chain lengths and different masses, therefore, in a polymer sample there is a distribution of molecular weights, but the average is 46000.
Identify the following as commercially significant monomers:
- vinyl chloride
- styrene, both by their systematic and common names Vinyl Chloride Styrene
Systematic Name: chloroethene Systematic Name: phenylethene/ ethenylbenzene Common Name: vinyl chloride Common Name: Styrene
Describe the uses of the polymers made from the above monomers in terms of their properties
Polymer Structure Properties Uses
Low density Polyethylene
Extensive chain branching; no stiffening side groups; no cross linking
Soft; flexible; low melting point; transparent; not strong
Cling wrap, carry bags; squeeze bottles
High density Polyethylene
No chain branching; no chain stiffening side groups; no cross linking
Hard; brittle; high melting point; translucent Kitchen utensils; food containers; milk bottles; rubbish bins Poly (vinyl chloride)
Considerable chain stiffening Cl side groups; polar C-Cl bonds produce strong intermolecular forces
Hard; inflexible; vulnerable to UV attack (inhibitor added to absorb UV light, preventing degradation) Electrical wiring insulation; garden hoses; drainage and sewerage pipes
Polystyrene Large phenyl stiffening side groups; minimal chain
branching; C-C and C-H bonds
Crystalline PVC -
transparent; resistant to UV attack; hard; rigid
Expanded PVC – light weight; spongy; moulded easily; good insulator; soft
Disposable drink glasses; foam packing material
Identify data, plan and perform a first-hand investigation to compare the reactivities of appropriate alkenes with their corresponding alkanes in bromine water
Background Information: Alkanes and alkenes are both non-polar molecules with weak dispersion (intermolecular) forces. Alkanes undergo substitution reactions (reactions in which an atom in a molecule is replaced by another atom or group of atoms). Alkenes undergo addition reactions. Aim: To compare the reactions of cyclohexene and cyclohexane in bromine water
Chlorine replaces hydrogen of regular ethene molecule Benzene ring replaces hydrogen of regular ethene molecule
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Safety Precaution: Cyclohexane, cyclohexene and bromine water are toxic by all routes of exposure; protect yourself by using small quantities, wearing safety glasses and avoiding inhalation by using a fume cupboard.
Method: In two test tubes, place ten drops of cyclohexene and cyclohexane. To each sample add 10 drops of bromine water. Shake vigorously. Observe for a colour change.
Results: cyclohexene decolourises bromine water; cyclohexane shows no reaction Conclusion:
Cyclohexene undergoes an addition reaction: C6H10 + Br2 → C6H10Br2 (cyclohexane + bromine water →
1, 2 – dibromocyclohexane)
Cyclohexane does not undergo a chemical reaction (however, in the presence of UV light, a substitution reaction occurs)
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9.2.2 Some scientists research the extraction of materials from biomass to reduce our
dependence on fossil fuels
Discuss the need for alternative sources of the compounds presently obtained from the petrochemical industry
There is only a finite supply of crude oil; all reserves will be used up within the next few decades
There is pressure to reduce energy use and greenhouse gas emissions
As oil supplies diminish, oil prices rise, one day to such an extent that oil will be too expensive a fuel source
When oil reserves run empty, if there were no alternatives, there would be no fuel or plastic For these reasons and more, ethanol and cellulose are being researched as alternative fuel sources.
Explain what is meant by a condensation polymer
Condensation polymers are chains of linked monomers that form when a functional group of one monomer reacts with the functional group of another monomer, joining and eliminating a small molecule (often water).
Describe the reactions involved when a condensation polymer is formed
Using glucose as an example,
… HO – C6H1004 – OH HO – C6H1004 – OH HO – C6H1004 – OH…
→ O – C6H10O4 – O - C6H10O4 – O - C6H10O4 + xH2O
When two glucose monomer molecules react through two -OH hydroxyl groups, a H-OH (water) molecule is condensed out, leaving an -O- linking the two monomer molecules. The first two glucose molecules to join condense out an H-OH, and every glucose molecule added to the growing chain then condenses out another H-OH.
Describe the structure of cellulose and identify it as an example of a condensation polymer found as a major component of biomass
Cellulose is a long, linear molecule because of the alternating CH2OH groups on either side of the chain
and C–O–C bond angles. Hydrogen bonds make cellulose difficult to break down, rigid and strong. The OH groups can not interact with water, making cellulose insoluble. Cellulose has potential as a
biopolymer because it is a major component of biomass (organic material derived from living organisms, e.g. crops, animal waste).
Identify that cellulose contains the basic carbon chain structures needed to build petrochemicals and discuss its potential as a raw material
Cellulose contains the basic carbon structures required by the plastics industry to be a source of chemicals or a biopolymer. Plastics made from cellulose biodegrade into fungi and bacteria. Cellulose → thermochemical pretreatment + hydrolysis → breaks down into constituent sugars → fermented → ethanol → polymerised to form useful products
However, the breaking down of cellulose requires much energy (as it is hydrogen bonded) and is more expensive than cracking crude oil→ form ethylene → hydrated to form ethanol. Nevertheless, research on cellulose needs to continue to see through its potential as a raw material.
5 Use available evidence to gather and present data from secondary sources and analyse progress in the recent development of a named biopolymer. This analysis should name the specific enzyme used or organism used to synthesise the material and an evaluation of the use or potential use of the polymer produced related to its properties
Polylactic acid (PLA) is a biodegradable thermoplastic derived from renewable plant material such as corn starch.
Corn kernels are milled → extract starch → enzymes break down starch in dextrose → lactic acid bacteria converts dextrose to lactic acid → polymerised → PLA
Uses:
Plant pots, mulch film and disposable nappies because of its biodegradability and compostability.
Its transparency, rigidity and crack-resistance makes PLA suited for use as food containers and drink cups, since the food would be visible and containers would not break.
Advantages of PLA:
Biodegradable; compostable
Sustainable since it is made from renewable resources – corn Less greenhouse gas emissions, no toxic gases
Requires less energy for production than conventional plastics Disadvantages of PLA:
Only biodegradable under ‘controlled composting environmental conditions’ which is not readily accessible by consumers
PLA breaks down into lactic acid, which demands a lot of oxygen. However, research is being conducted into anaerobic digesters so that PLA can break down without oxygen.
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9.2.3 Other resources such as ethanol are readily available from renewable resources
such as plants
Describe the dehydration of ethanol to ethylene and identify the need for a catalyst in this process and the catalyst used
C2H5OH → C2H4(g) + H2O(l) in the presence of heat and concentrated sulfuric acid which acts as a
dehydrating agent and catalyst
Describe the addition of water to ethylene resulting in the production of ethanol and identify the need for a catalyst in this process and the catalyst used
C2H4(g) + H2O(l) → C2H5OH in the presence of heat and dilute sulfuric acid which acts as a hydrating
agent and catalyst
Describe and account for the many uses of ethanol as a solvent for polar and non-polar substances
The ethanol molecule consists of two parts – the non-polar alkyl (-CH2CH3) end and the polar
hydroxyl (-OH) end.
Outline the use of ethanol as a fuel and explain why it can be called a renewable resource
Ethanol has been proposed as an alternative fuel source because:
It undergoes complete combustion and unlike petrol, burns efficiently → does not release soot or carbon monoxide.
It is a renewable resource since it is made through fermentation of plant material, and so would reduce our dependency on non-renewable crude oil.
Ethanol is made from carbon dioxide, water and sunlight and plant material and when it is burnt it returns back to carbon dioxide and water, which can be reconverted into ethanol.
Assess the potential of ethanol as an alternative fuel and discuss the advantages and disadvantages of its use
Advantages of ethanol as a fuel:
At 10-20% concentration in petrol, it is a ‘petrol extender’ and vehicle engines do not need to be modified to utilise it.
It undergoes complete combustion [see previous dot point]
It is ‘greenhouse neutral’ since the carbon dioxide released during fermentation and combustion is used up in photosynthesis.
i.e. Fermentation: C6H12O6 → 2C2H5OH + 2CO2 (g)
+ Combustion: 2C2H5OH + 3O2 (g)→ 4CO2 (g) + 6H2O(l)
= Photosynthesis: 6CO2 (g) + 6H2O (l) → C6H12O6 + 6O2
Disadvantages of ethanol as a fuel:
Large areas of agricultural land need to be devoted to growing crops for fuel instead of for food
Non-polar alkyl end forms dispersion forces with non-polar substances, thus dissolving them
Polar hydroxyl end forms dipole-dipole or hydrogen bonds with polar substances, thus dissolving them
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It does not release as much energy as petrol → vehicles receive fewer kilometres (-1367kJ in ethanol vs. -5464kJ in petrol)
Smelly wastes present environmental problems
Process information from secondary sources to summarise the processes involved in the industrial production of ethanol from sugar cane
Process information from secondary sources to summarise the use of ethanol as an alternative car fuel, evaluating the success of current usage
In Australia, ethanol is currently added to petrol to form E10 Unleaded which is 10% ethanol/90% petrol blend. At this concentration, vehicle engines need not be modified; car manufacturers claim that higher concentrations corrode vehicle engines. Ethanol is more expensive and less efficient than petrol as a car fuel because and vehicles receive fewer kilometres because ethanol contains less energy. Furthermore, there are no reliable studies to show that ethanol produces less greenhouse gas emissions than petrol (although it does combust cleanly and produces less soot and carbon monoxide).
Describe conditions under which fermentation of sugars is promoted
Fermentation is a process in which glucose is broken down into ethanol and carbon dioxide by the action of enzymes in yeast. For fermentation:
Suitable grain or fruit is mashed up with water Yeast is added
Air is excluded
The mixture is kept at 37°C (body temperature)
Sugarcane Harvested and
transported to mill Clarified to remove impurities Shredded and crushed to extract sugar-rich juice Mixed with enzymes
and water to convert sugarcane into pure
sugars Heated to turn starch
into liquid
Cooled; yeast is added to convert sugars to ethanol and
carbon dioxide Carbon dioxide is removed then remaining mixtureis fractionally distilled Pure ethanol
8 Summarise the chemistry of the fermentation process
Present information from secondary sources by writing a balanced equation for the fermentation of glucose to ethanol
1. Starch/ sucrose are mixed with enzymes (a biological catalyst) to convert it into glucose. 2. Glucose mixture is clarified to remove impurities and waste, cellulose.
3. Yeast is added to convert mixture into carbon dioxide and ethanol (fermentation). i.e. C6H12O6 → 2C2H5OH + 2CO2 (g)
4. Fermented further to produce ethanol at 15% concentration (any higher would kill yeast, ceasing further fermentation)
5. 15% ethanol mixture is fractionally distilled → 95% ethanol
Solve problems, plan and perform a first-hand investigation to carry out the fermentation of glucose and monitor mass changes
Method:
1. Measure and record the mass of the conical flask.
2. Add an aqueous mixture of glucose and measure and record the mass of the flask plus the glucose mixture. Calculate the mass of the glucose mixture.
3. Add 1 gram of yeast and secure flask opening with lid containing a pipe which leads to a test tube containing limewater.
4. Incubate overnight.
5. Measure mass of conical flask with fermented ethanol mixture. Calculate mass changes. 6. Record observations of limewater.
Results: 5 grams loss, limewater turns milky as carbon dioxide is produced during fermentation.
Note: Additionally, you can test for ethanol by adding potassium permanganate (KMnO4) to a sample
of the final mixture, it should turn colourless.
Define the molar heat of combustion of a compound and calculate the value for ethanol from first-hand data
The molar heat of combustion of a substance is the heat liberated when one mole of the substance undergoes complete combustion with oxygen at standard atmospheric pressure with products being carbon dioxide and water. Value for ethanol: -1367 kJ mol-1
Identify the IUPAC nomenclature for straight-chained alkanols from C1 to C8
1 – meth 2 – eth 3 – prop 4 – but 5 – pent 6 – hex 7 – hept 8 – oct Formula for alkanols: CnH2n+1OH
Formula for alkanes: CnH2n+2
Formula for alkenes: CnH2n
Identify data sources, choose resources and perform a first-hand investigation to determine and compare heats of combustion of at least three liquid alkanols per gram and per mole
Calculating Heat Capacity: ∆H= -mC∆T
i.e. Heat Capacity (in joules) = mass in grams of water X 4.18 X temperature change Heat of Combustion per gram: ∆H (in kilojoules) ÷ mass of alkanol used
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9.2.4 Oxidation-reduction reactions are increasingly important as a source of energy
Explain the displacement of metals from solution in terms of transfer of electrons
Reactions between a metal and a solution containing the ions of a different metal are displacement or redox reactions. The metal is oxidised and dissolves and the ions of the metal in solution are reduced to form a solid metal deposit.
Oxidation: loss of electrons (LEO) Reduction: gain of electrons (GER)
Identify the relationship between displacement of metal ions in solution by other metals to the relative activity of metals
Displacement reactions are electron transfer reactions. In such reactions, a more active solid metal oxidises and will displace the ions of a less active metal in solution (which is reduced).
Note: On the standard electrode potentials table, the metal higher up is oxidised.
Account for changes in the oxidation state of species in terms of their loss or gain of electrons
Oxidation → increase in oxidation state Reduction → decrease in oxidation state
Elements in their naturally occurring state have an OS of 0.
For ions, the OS is the charge/ valency of the ion (e.g. SO42- has OS -2)
Examples:
1. Find the OS of the underlined species in Cr2O7
2-Sum of oxidation states of constituent elements=overall charge of entire species Therefore, 2x + 7×-2 = -2 → x=6: the OS of Cr2 is 6.
2. If MnCl3 is converted to MnO2, determine whether oxidation or reduction has occurred.
OS of Mn in MnCl3 is +3; OS of Mn in MnO2 is +4 → increase in OS → oxidation has occurred Describe and explain galvanic cells in terms of oxidation/reduction reactions
A galvanic cell is an electron pump that produces electricity by pumping electrons out of the anode, where oxidation occurs, into an external circuit (a metallic conductor) and draws them back into the cathode, where reduction occurs.
Outline the construction of galvanic cells and trace the direction of electron flow
A galvanic cell consists of:
Two different half cells, consisting of an electrode in electrolyte solution
An external circuit, which allows the flow of electrons from the anode to cathode A salt bridge, which allows the migration of ions and maintains electrical neutrality
Define anode, cathode, electrode and electrolyte to describe galvanic cells
Anode: electrolyte at which oxidation occurs Cathode: electrolyte at which reduction occurs
Electrode: conductor of a cell (metal or carbon) which gets connected to the external circuit Electrolyte: a substance which in solution conducts electricity
10 Perform a first-hand investigation to identify the conditions under which a galvanic cell is produced
Method: Set up a simple galvanic cell and observe the voltage. Observe what happens to the voltage when: the salt bridge is removed; and electrode is removed; two identical half cells are used.
Results: 0.7 V with 2 different half cells; 0 V without salt bridge; 0 V without electrode; 0 V with two identical half cells.
Conclusion: The essential features of a galvanic cell are: two different half cells; two different electrodes; a salt bridge.
Perform a first-hand investigation and gather first-hand information to measure the difference in potential of different combinations of metals in an electrolyte solution
Method:
Set up a beaker of sulfuric acid. Attach a piece of copper to a
voltmeter and place this in the beaker. Attach different metals (Zn, Fe, Pb and Mg) to the other end of the voltmeter and place this end in the beaker as well. Measure the potential difference of various
combinations of metals against copper.
OR an alternate method: Set up a standard galvanic cell with a voltmeter and one half cell of copper in copper sulfate solution; let the other half cell be metal x in its corresponding solution. Measure and record the potential difference. Repeat process with different combinations of metal electrodes (Zn, Fe, Pb and Mg). The reliability of the investigation can be increased by the use of repeat trials for the various combinations you have chosen.
Results: Greatest Potential Difference → Least Potential Difference: Mg → Zn → Fe → Pb
Solve problems and analyse information to calculate the potential E° requirement of named electrochemical processes using tables of standard potentials and half equations
Process with worked example:
Calculate the standard cell potential of Co I Co2+ II Ag+ I Ag
1. Determine which metal is oxidised and which is reduced by referring to the standard electrode potentials table; the metal higher up is oxidised.
Co is oxidised; Ag is reduced
2. Write the reduction half equation and record E°. 2Ag+(aq) + 2e- → 2Ag(s) E°: +0.80
3. Write the oxidation half equation by reversing the appropriate reduction equation and record E° and also reversing its sign.
Co(s) → Co2+(aq) + 2e- E°: +0.28
4. Add the two half equations to form a redox equation (the electrons should cancel out). Add the E° potentials (a positive E° indicates a spontaneous reaction)
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2Ag+(aq) + Co(s) → 2Ag(s) + Co2+(aq) E°: 0.80 + 0.28 = 1.08
Gather and present information on the structure and chemistry of a dry cell and evaluate it in
comparison to the button cell in terms of: chemistry, cost and practicality, impact on society, impact on environment
Dry Cell Button Cell
Structure
Consists of a zinc anode, an aqueous paste of ammonium chloride and a mixture of powdered carbon, manganese, and a carbon graphite rod as the cathode.
Consists of a zinc anode, silver oxide cathode with a conductive substance C or Ag mixed with Ag2O, and an alkaline
electrolyte solution of potassium hydroxide.
Chemistry O: Zn(s) → Zn2+(aq) + 2e-
R: 2NH4+(aq) + 2MnO2(s) → Mn2O3(s) +
2NH3(aq) + H2O(l)
O: Zn(s) + 2OH-(aq) → Zn(OH)2(s) + 2e-
R: Ag2O(s) +H2O(l) + 2e- → 2Ag(s) + 2OH- Cost and
Practicality
Cheap; mass produced; useful in devices requiring a small current, e.g. torches, portable radios, clocks; leaking problems common because zinc casing erodes during operation
Expensive due to silver content; but practical because of its small size and constant voltage (1.5 V)
Impact on Society Made the use of portable appliances
such as radios possible and feasible.
Has allowed for use of miniature appliances such as hearing aids and watches; its long life means it does not need frequent replacement → practical.
Impact on Environment
Small quantities of zinc, ammonium and carbon are harmless; but not rechargeable and large in size → increase in landfills.
Little environmental impact as it takes up little space in landfill; contents are less likely to leak and less toxic.
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9.2.5 Nuclear chemistry provides a range of materials
Distinguish between stable and radioactive isotopes and describe the conditions under which a nucleus is unstable
Isotopes are atoms of one element that differ by having different numbers of neutrons but the same number of protons, hence have a different mass number. An isotope is unstable if:
Its atomic number is greater than 83
Its ratio of neutrons to protons places it outside the zone of stability
The nucleus of a large radioactive atom has high energy, and therefore is unstable; energy is reduced by expelling alpha, beta or gamma particles, and as a result the atom is more stable.
Describe how transuranic elements are produced
Transuranic elements are artificial, man made elements with atomic numbers greater than 92. Early transuranic elements were made in nuclear reactors where existing elements were bombarded with slow thermal neutrons which were absorbed to produce new elements. Later transuranic elements were made in particle accelerators where heavy nuclei would be bombarded with positive particles at high speed; since both the target and particles are positively charged, fast speed is needed to
overcome electrostatic repulsion.
Describe how commercial radioisotopes are produced
Radioisotopes are made in the same way as transuranic elements, i.e. in nuclear reactors or particle accelerators; the only difference is that the target nuclei need not be heavy.
Process information from secondary sources to describe recent discoveries of elements
Elements 104-112 (not 108) have been produced using particle accelerators. For example, Element 106, Seaborgium, is made in a particle accelerator involving fusion of an isotope of Californium 249Cf with an oxygen isotope:
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8 O + 249 98Cf → 267106Sg + 4 1 0n
Identify instruments and processes that can be used to detect radiation
Photographic film: radiation is monitored by measuring the extent of the darkening of photographic film.
Cloud chamber: as radiation passes through a supersaturated vapour, it ionises the air, forming water droplets; the path of these droplets indicates the type of radiation: alpha – straight, dense tracks, beta – less dense, wavy tracks, gamma – faint, random droplets.
Geiger- Muller counter: as radiation passes the Geiger tube, it hits gas molecules and ionises them. Audible electrical pulses are produced as gas molecules are ionised, the rate of these pulses indicates the amount of radiation.
Scintillation counter: a flash of light is emitted as substances are irradiated with radiation.
Identify one use of a named radioisotope: in industry; in medicine
Technetium-99m is used in medicine for diagnosis of blood, heart, brain and thyroid abnormalities. Cobalt-60 is used in medicine for cancer therapy and in industry as a thickness gauge.
Describe the way in which the above named industrial and medical radioisotopes are used and explain their use in terms of their chemical properties
13 Radioisotope Radiation
Emitted
Half-life Uses in terms of Chemical Properties
Technetium-99m Gamma 6 hours Medical diagnosis – gamma rays are highly penetrable so can be detected at the body’s surface without invasion; short half life ensures it leaves the body quickly, leaving minimal damage; easily bonds with other chemicals, so it combines with tin to form a serum which is injected into the body and inside the body it binds with red blood cells to detect circulation disorders.
Cobalt-60 Beta; Gamma
5.3 years
Thickness gauge – beta and gamma rays penetrate through metal sheets (but only to a certain degree); relatively long half life makes it suited in machinery since the radioactive source does not need frequent replacement.
Cancer therapy – gamma rays kill cancer cells because they contain high energy; half life is short enough to expel reasonable bouts of radiation at moderate intensity to kill cancer cells.
Use available evidence to analyse the benefits and problems associated with the use of radioisotopes in identified industries and medicine
Field Benefits Problems
Medicine Opens a wide range of non-invasive
diagnostic procedures
Radiation therapy is, in most cases, the most effective treatment for cancer
Harmful to people and life forms Causes tissue damage: skin burns,
nausea, radiation sickness
Can cause cancer: leukaemia and lung cancer
Genetic damage: deformities in offspring
Industry Monitoring equipment are more
sensitive and accurate
Enables examination of weld and structural faults in buildings and machinery which otherwise can not be detected
Radiation from equipment can stray and leak if not carefully monitored or stored in well shielded containers
Writing Nuclear Equations: x+y
a+bM ↔ xaP + ybR
Symbols: Neutron: 10n
Proton: 11p
Electron/ beta particle: 0-1e
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9.3.1 Indicators were identified with the observation that the colour of some flowers
depends on soil composition
Classify common substances as acidic, basic or neutral
Common acids: vinegar, lime/ lemon juice, aspirin, vitamin C Laboratory acids: hydrochloric, sulfuric and nitric
Neutrals: water
Common bases: ammonia, washing soda, antacid tablets, oven/drain cleaners Laboratory bases: all hydroxides and oxides, e.g. NaOH, Mg(OH)2, Fe2O3
Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothymol blue can be used to determine the acidic or basic nature of a material over a range, which is identified by colour changes
Identify data and choose resources to gather information about the colour changes of a range of indicators
Common Indicator Colour in Acid Colour in Base
Litmus Red (<5.0) Blue (>8.0)
Phenolphthalein Colourless (<8.3) Pink (10.0)
Methyl Orange Red (<3.1) Yellow (>4.4)
Bromothymol Blue Yellow (<6.0) Blue (>7.6)
Perform a first-hand investigation to prepare and test a natural indicator
Method: Heat red cabbage leaves in a beaker under a Bunsen burner. Stop burner after water turns purple and allow to cool. Pour purple water in three test tubes - one containing acid (HNO3), one with
water and one with a base (NaOH).
Results: Colour in: acid - red; water - purple; base - yellow
Identify and describe some everyday uses of indicators including testing of soil acidity/basicity Pool Ranger Home Swimming Pool pH Tester
Test: Collect pool sample in tube then add 5 drops of phenol red and shake. Compare tube colour with adjacent colour markings.
Importance: Acid or basic pool water results from a buildup of bacteria and pollutants. Pools need to maintain a relatively neutral pH as acidic or alkaline water causes irritations to the skin and eyes.
Aquasonic Home Aquarium pH Test Kit
Test: Collect water from aquarium in tube and add 3 drops of bromothymol blue. Compare test tube colour with chart provided. Add appropriate chemicals (if needed) to maintain pH between 6.0 - 7.8. Importance: pH is lowered when there is a buildup of bacterial wastes. A pH between 6.0 - 7.8 provides the ideal conditions for freshwater aquaria. Extreme changes in pH → reproductive abnormalities, dissolving of scale mucus membrane.
Soil pH Test Kit
Test: Place a teaspoon of sample soil on test plate and add indicator liquid and barium sulfate white powder then stir. Compare colour of sample with colour card provided.
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Importance: Different plants prefer/grow best in different soils with varying pH, e.g. azaleas prefer slightly acidic, vegetables prefer slightly alkaline, no plants grow if pH < 4. Inadequate pH → death, impaired plant growth.
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9.3.2 While we usually think of the air around us as neutral, the atmosphere naturally
contains acidic oxides of carbon, nitrogen and sulfur. The concentration of these
acidic oxides have been increasing since the Industrial Revolution
Identify oxides of non-metals which act as acids and describe the conditions under which they act as acids
CO2 – carbon dioxide, SO2 – sulfur dioxide and NO2 – nitrogen dioxide are oxides of non-metals which
dissolve in water to form acids and react with bases to form salts and water.
Analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and their acidity/basicity of oxides
Acidic character of elements increases across a period. Elements of the left (metals) form basic oxides; elements in the middle form amphoteric oxides (i.e. they display both acidic and basic character); elements on the right (non-metals) form acidic oxides; noble gases are inert so do not form any oxides.
Define Le Chatelier’s principle
If a system at equilibrium is disturbed, then the system adjusts itself as to minimize the disturbance.
Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle
CO2(g) + H2O(l) ↔ H2CO3(aq) ↔ H+(aq) + HCO3-(aq) ∆H: negative
More CO2 → shift to the right → more H+(aq) + HCO3-(aq)
Less CO2 → shift to the left → less H+(aq) → less acidic
Add H+ or HCO
3- → shift to the left → more CO2(g) + H2O(l)
Add base → reacts with H+ → shift to the right
Add heat→ shift to the left (since reverse reaction absorbs heat) *vice versa+
Note: If ∆H is negative, the forward reaction is releasing heat, if ∆H is positive, forward reaction is absorbing heat.
Identify factors which can affect the equilibrium in a reversible reaction
Increase in concentration → reaction shifts to the side which uses up the added species
Decrease in concentration (removed) → reaction shifts to the side which produces the removed species
Increase in volume → decrease in pressure → shift to the side which produces the most
gaseous molecules. Note: No change if both sides produce same number of gaseous molecules Decrease in volume → increase in pressure → shift to the side which produces the least
gaseous molecules
Increase temperature → shift to the endothermic side that uses up the added heat Decrease in temperature → shift to the exothermic side that produces the removed heat
Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen
Source Sulfur Dioxide Oxides of Nitrogen
Natural Geothermal hot springs
Volcanoes
Smoke from bushfires
High temperatures from lightning (which forms nitric oxide, NO)
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Bacterial decomposition in soil (which forms nitrous oxide, HNO2)
Industrial Burning of fossil fuels (such as coal,
which contains sulfur)
Extraction of metals from sulfide ores (smelting)
Combustion of nitrogen and oxygen in vehicle chambers and power stations (which forms nitric oxide, as in lightning)
Describe, using equations, examples of chemical reactions which release sulfur dioxide and oxides of nitrogen
As bacteria decompose, it forms hydrogen sulfide (H2S) which oxidises to form sulfur dioxide:
2H2S(g) + 3O2(g) → 2SO2(g) + 2H2O(g)
Coal contains sulfur, so when it is burnt, it reacts with oxygen to form sulfur dioxide: S [in compound] + O2(g) → SO2 (g)
High temperatures from lightning or vehicle chambers causes nitrogen and oxygen to combine forming nitric oxide, which further oxidises to form nitrogen dioxide:
O2(g) + N2(g) → 2NO(g) then 2NO(g) + O2(g) → 2NO2(g)
Assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen
Industrial Revolution → increased burning of coal + extraction of metals → increased SO2 emissions
Electricity + motor car → high temperature combustion → increased levels of nitrogen oxides There is a lack of reliable evidence before 1970 since there was a lack of technology that monitored atmospheric concentrations, and also the technology that existed at the time was insufficient in monitoring such low levels which existed then since the problems were just emerging.
Calculate volumes of gases given masses of some substance in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0°C and 100kPa or 25°C and 100kPa
Mole Ratio Equations:
n = m ÷ fm i.e. number of moles = mass ÷ formula mass
n = V ÷ mV i.e. number of moles= volume ÷ molar volume which is 22.71 at 0°C/100kPa or 24.79 at 25°C/ 100kPa
Process with worked example:
What volume of carbon dioxide gas measured at 0°C/100kPa can be reacted from 15.5g of NaOH to form NaCO3 and water?
1. Write a chemical equation
2NaOH(s) + CO2(g) → Na2CO3 (s) + H2O(l)
2. Find the number of moles of the known substance nNaOH = 15.5 ÷ 39.998 = 0.3875…
3. Determine the mole ratio (i.e. moles of unknown : moles of known) CO2: NaOH = 1:2
4. Find the moles of unknown (i.e. multiply moles of known by mole ratio) nCO2 = ½ × nNaOH = 0.3875÷2
5. Convert moles of known into units asked for V= (0.3875÷2) × 22.71 = 4.4 L
Explain the formation and effects of acid rain
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i.e. SO2(g) + H2O(l) → H2SO3(aq) then 2H2SO3(aq) + O2(g) → 2H2SO4(aq)
Similarly, NO2 dissolves in the air to form nitrous acid and nitric acid.
i.e. 2NO2(g) + H2O(l) → HNO2(aq) + HNO3(aq)
Sulfuric acid and nitric acid together form acid rain. Effects of acid rain:
Increases the acidity of lakes → detrimental effect on fish as it upsets their reproductive processes and strips mucus membrane on their scales
Damages pine forests and vegetation → corrodes leaves, changes soil concentration and pH Erodes marble, limestone and metal structures (carbonates in these materials react with acid →
weathering and erosion), e.g. CaCO3(s) + 2H+(aq) → Ca2+(aq) + CO2(g) + H2O(l)
Analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment
Buildup in atmosphere of sulfur dioxide and oxides of nitrogen → formation of acid rain + photochemical smog → detrimental health effects on population (e.g. breathing difficulties) + detrimental effects of acid rain (see previous syllabus dot point). Therefore, the rate of emission of these chemicals needs to be monitored and regulated.
Identify data, plan and perform a first-hand investigation to gather data to measure the mass changes involved and calculate the volume of gas released at 25°C and 100kPa
Method: Weigh an unopened soda can on an electronic balance. Similarly, weigh another soda can that has been refilled with water – this is a control. Open the soda can. Leave both cans overnight. Reweigh add record the mass changes. Calculate the volume of CO2 released.
19
9.3.3 Acids occur in many foods, drinks and even within our stomachs
Define acids as proton donors and describe the ionisation of acids in water
An acid ionises in water to form hydronium ions and donates a proton to form a conjugate base: Acid (aq) → H3O+ + Conjugate Base
Identify acids such as acetic (ethanoic acid), citric (2-hydroxypropane-1,2,3-tricarboxylic acid), hydrochloric and sulfuric acid
Acetic acid (vinegar), citric acid (in fruits and as a preservative), hydrochloric (in the stomach and made industrially) and sulfuric acid (in acid rain) are common acids.
Describe the use of the pH scale in comparing acids and bases
The pH scale is used to compare the concentration of hydrogen ions in solutions of acids and bases In a neutral solution, e.g. water, [H+] = [OH-] = 10-7 mol L-1 and so pH =7.
In an acidic solution, [H+] > 10-7 mol L-1 and pH < 7.
In a basic solution, [H+] < 10-7 mol L-1 and pH > 7.
pH increases as [H+] decreases
Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute
Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids
In a strong acid, all acid molecules ionise, there are no neutral acid molecules.
In a weak acid, only a small percentage of acid molecules ionise, most remain as neutral molecules. The concentration of an acid refers to its molarity; concentrated if it is above 5 mol L-1 and dilute if it is less than 2 mol L-1.
Identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+] To calculate pH given [H+]: pH = -log10 [H+]
To calculate [H+] given the pH: [H+] = 10-pH To calculate [H+] given [OH-]: [H+] = 10-14 ÷ [OH-]
As pH increases by 1, the concentration of the hydrogen ions, i.e. [H+], decreases by a factor of ten, or ten fold. E.g. if pH = 1, [H+] = 10-1 = 0.1 but if pH = 2, [H+] = 10-2 = 0.01
20 Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules
Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids
HCl: strong acid → ionises completely → high *H3O++ → low pH
Citric and Acetic: weak acids → only partially ionise → less *H3O++ → higher pH
Describe the difference between a strong and a weak acid in terms of equilibrium between the intact molecule and its ions
An aqueous solution of a strong acid contains only hydronium ions and the anions of the acid; there are no neutral acid molecules, i.e. the ionisation reaction goes to completion.
An aqueous solution of a weak acid is at equilibrium between the neutral acid molecules and hydronium ions and anions of the acid [see ionisation equations in table above].
Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals
Using a pH meter/ probe is a non destructive means of measuring the pH of chemicals. Using indicators is a destructive means.
Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids
Method: Using a pH probe measure and record the concentration of a strong acid, such as HCl. Rinse the probe with distilled water to avoid cross contamination. Repeat with a weak acid, such as acetic acid.
Gather and process information from secondary sources to explain the use of acids as food additives
Acids are added to foods to:
Preserve food – acids lower the pH to a range outside one which microorganisms can survive in, thus enzyme reactions are inhibited or slowed down. e.g. citric acid is a preservative in jams; acetic acid preserves canned beetroot and pickled onions.
Enhance flavour and/or nutritional value – provides a tart, sour taste. e.g. phosphoric acid in soft drinks; citric acid and ascorbic acid are antioxidants.
Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition
Acid or Base Name Chemical Formula Where is it Found?
Hydrochloric Acid HCl In the stomach
Lactic Acid CH3CH(OH)CO2H Milk and Yoghurt
Ammonia (base) NH3 Formed in the anaerobic decay of organic matter
Calcium Carbonate (base) CaCO3 Limestone
Acid Acetic Citric Hydrochloric
Ionisation Equation
CH3COOH (aq) + H2O (l) ↔
H3O+ (aq) + CH3COO- (aq)
C6H8O7 (aq) + 3H2O (l) ↔ C6H5O73- (aq) + 3H3O+ (l) HCl (g) + H2O (l) → H3O+ (aq) + Cl- (aq) pH 2.9 2.1 1.0 Degree of Ionisation 1% 8% 100%
21 Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations
Example:
Calculate the pH of 2.0 mol L-1 solution of sulfuric acid [H2SO4] = 2.0 H2SO4 (l) → 2H+ (aq) + SO42- (aq) ratio of H2SO4 : H+ = 1 : 2 [H+ ] = 2 × 2.0 = 4.0 pH = -log10 4.0 = -0.6
Process with Harder Worked Example:
[A neutralisation reaction] 50 mL of 0.100 mol L-1 hydrochloric acid is added to 75 mL of 0.050 mol L-1 sodium hydroxide solution. Calculate the pH of the resulting solution.
1. Equation states mole ratio. Calculate the moles of H+ from acid information. VHCl = 0.050; CHCl = 0.100
C = n ÷ V
nH+ = C × V = 0.100 × 0.050 = 0.005 moles
2. Calculate the moles of OH- from basic information. VNaOH = 0.075; CNaOH = 0.050
nOH- = 0.050 × 0.075 = 0.00375 moles
3. Determine which is in excess and by how many moles. H+ is in excess
Excess = (moles of H+)– (moles of OH-) = 0.005 – 0.00375 = 0.00125 4. Find the concentration of excess H+ or OH-.
[H+] = moles ÷ total volume = 0.00125 ÷ (0.050 + 0.075) = 0.01
Note: If OH- is in excess, calculate its excess, then calculate [H+] using [H+] = 10-14 ÷ [OH-] 5. Calculate pH.
22
9.3.4 Because of the prevalence and importance of acids, they have been used and
studied for hundreds of years. Over time, the definitions of acid and base have
been refined
Outline the historical development of ideas about acids including those of:
Lavoisier Davy Arrhenius
Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions
Scientist by Date
Observations Theory of Acids Example
Lavoisier (1780)
Non-metal oxides dissolve in water to produce acids
Acids contain oxygen CO2 (g) + H2O (l) →
H2CO3 (aq)
Davy (1815)
Decomposed HCl and found that it did not contain oxygen
Acids contain replaceable hydrogen (ability to be replaced by metals) Zn (s) + 2HCl (aq) → ZnCl2 (aq) + H2 (g) Arrhenius (1884)
When an electrical current was passed through acid, hydrogen gas evolved at the anode
Acids ionise in solution to produce hydrogen ions
CH3COOH (l) →
H+ (aq) + CH3COO(aq)
The Bronsted-Lowry definition [see below] expanded our understanding of acids/bases as it proposes that the acidity of a substance depends on its properties relative to those of the solvent, and not just its structure. This concept furthers our understanding of acid-base equilibrium and pH calculations.
Outline the Brönsted-Lowry theory of acids and bases
Bronsted-Lowry (1923) analysed the experiments of the scientists before them to propose their theory that acids are proton donors and bases are protons acceptors.
e.g. HA + H2O → H3O+ + A- where HA is an acid or B + H2O → HB+ + OH-where B is a base
Describe the relationship between an acid and its conjugate base and a base and its conjugate acid
Acid (aq) → H3O+ + Conjugate Base when the acid donates a proton
Base + H+ → Conjugate acid when a base accepts a proton
Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature
Strong Base e.g. H2O, NaOH Weak Base e.g. NH3
Strong Acid
e.g. HNO3, HCl
Neutral salt (neither the conjugate base of the acid nor the conjugate acid of the base significantly react with water to alter the pH)
Acidic salt (as cation reacts)
Weak Acid
e.g. CH3COOH
23 Identify conjugate acid/base pairs
Base Conjugate Acid
CO32- HCO3-
HCO3- H2CO3
NH3 NH4+
OH- H2O
H2O H3O+
Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions
Amphiprotic substances are those which can react as both a proton donor and a proton acceptor. e.g. water - [as a B.L. acid]: H2O → H++ OH-; [as a B.L. base]: H2O + H+ → H3O+
or HCO3- - [as a B.L. acid]: HCO3- → H+ + CO32-; [as a B.L. base]: HCO3- + H+ → H2CO3 Identify neutralisation as a proton transfer reaction which is exothermic
Neutralisation reactions are exothermic proton transfer reactions.
e.g. HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq) – a H+ proton from HCl reacts with OH- in NaCl to form water;
heat is liberated → ∆H < 0 (approx. -56kJ mol L-1)
Describe the correct technique for conducting titrations and preparation of standard solutions
Perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases
Preparing a Standard Solution
1. Measure the mass of the primary standard using a beaker and an electronic balance. 2. Dissolve the standard in 100 mL of water.
3. Transfer dissolved solution to a volumetric flask using a filter funnel.
4. Rinse beaker with distilled water and transfer rinsed solution to volumetric flask. Swirl gently. 5. Fill the volumetric flask until the bottom of the meniscus is in line with the calibration mark. 6. Invert and shake to ensure a homogenous mixture.
7. Calculate the concentration of the solution: n = mass ÷ molar mass then C = n ÷ V (0.250) To be a primary standard, the substance must be:
Of high purity and stability not be volatile, not absorb moisture or react with CO2, as these
would create an impure substance Soluble in water
Of accurately known concentration e.g. Na2CO3 or NaHCO3 but not NaOH Conducting a Titration
1. Collect 200 mL of the base solution (e.g. NaOH) in a beaker.
2. Rinse the burette with 10 mL of the base solution, including running the solution through the stopcock.
3. Fill the burette and attach it to the retort stand using a burette clamp. Record the volume of base solution in the burette.
4. Rinse the pipette with 10 mL of the acid solution (e.g. HCl). 5. Pipette out 25 mL accurately into a clean, dry conical flask.
Acid Conjugate Base
HCl Cl-
H2SO4 HSO4-
NH4+ NH3
H2O OH-
24
6. Add 2 drops of phenolphthalein indicator (or another indicator that will show a colour change as the equivalence point is reached).
7. Run the base solution from burette, at the same time, gently swirling the conical flask. As the equivalence point gets closer, add solution drop by drop. Stop when the solution turns a pale pink colour (for phenolphthalein). [Carry out a rough titration first to get an estimate]
8. Repeat another 2 to 3 times for reliability.
9. Determine the concentration of the acid solution, using CiVi = CfVf.
Titration Calculation Process with Worked Example:
A student performed a titration of 25.00 mL of acetic acid of unknown concentration with 0.123 mol L-1 solution of sodium hydroxide. Her results are shown right:
Calculate the concentration of the acetic acid solution. 1. Write an equation
NaOH (aq) + CH3COOH (aq) → NaCH3COO (aq) + H2O (l)
2. Find the moles of one species.
Average VolumeNaOH = (22.70 + 22.70 + 22.80) ÷ 3 = 22.73
Note: Trial 1 is a rough titration so do not include this result.
nNaOH = 0.123 × 0.02273 = 0.00280
3. Use mole ratio to find moles of required species. Mole ratio of NaOH : CH3COOH = 1 : 1
nCH3COOH = 0.00280
4. Find the concentration of the required species. C = n ÷ V = 0.00280 ÷ 0.02500 = 0.112 mol L-1
Perform a first-hand investigation to determine the
concentration of a domestic acidic substance using computer based technologies
Method: Carry out a titration reaction to determine the mass of acetic acid in vinegar with the aid of technologies such as a magnetic stirrer and computer programs that plot the pH as the titrant from the burette is added.
Qualitatively describe the effect of buffers with reference to a specific example in a natural system
A buffer solution is one that contains comparable amounts of weak acid and its conjugate base, so therefore is able to maintain an approximately constant pH even when significant amounts of strong acid or base are added.
Example of a naturally occurring buffer: H2CO3 (aq) + H2O (l) ↔ H3O+ (aq) + HCO3- (aq)
As rainwater falls, it forms a dilute solution of carbonic acid with a pH of 5.7 – 6.0: CO2 (g) + H2O (l) → H2CO3 (aq)
When it lands in the lake or river, comparable amounts of HCO3- from surrounding carbonate rocks,
such as limestone, provide a natural buffer and react with H3O+ in carbonic acid → shifts the
equilibrium to the left → removes H+ → raises pH to 6.5 – 7.0:
Trial Volume of NaOH used (mL)
1 23.30
2 22.70
3 22.70
25
H3O+(aq) + HCO3-(aq) → H2CO3 + H2O (l)
OH- in water then reacts with H3O+ → shifts equilibrium to the right → removes OH- → restores/
maintains slightly acidic pH:
OH- (aq) + H2CO3 (aq) → HCO3-(aq) + H2O (l)
Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions
Equipment: Test tubes Universal Indicator pH chart A range of 0.1 mol L-1 salt solutions
Method: Add 5 mL a salt solution into a test tube. Add 3 drops of universal indicator and record the colour and corresponding pH with reference to a pH chart.
Conclusion: From the Bronsted-Lowry theory of acids and bases, the ions in the salt solutions act as acids or bases when reacting with water; acidic solutions donate H+; basic solutions accept H+.
Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills
As acids are highly corrosive and bases are highly caustic, it is important to neutralise any spills of these substances if they occur. Amphiprotic substances like HCO3- in NaHCO3 is suitable for neutralising both
acidic and alkaline spills as it can act as a B.L. acid or base: In an acidic spill: H+ + HCO3- → H2CO3
In an alkaline spill: HCO3- + OH- → H2O + CO3
2-NaHCO3 is preferred as it is can be safely handled and stored, it is cheap, if excess is used it is not
26
9.3.5 Esterification is a naturally occurring process which can be performed in the
laboratory
Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds
Alkanols contain a hydroxyl OH- functional group attached to a C atom. General formula: CnH2n+1OH
Alkanoic acids are polar, weak acids that contain a carboxylic –COOH group attached to a C atom. General formula: CnH2n+1COOH
Alkanols are a subset of alcohols and alkanoic acids are a subset of carboxylic acids. Both alkanols and alkanoic acids only contain C, H and O atoms.
Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid and straight-chained primary alkanol structures
Alkanoic acids have higher melting and boiling points than their corresponding alkanols. Alkanols have relatively high boiling points as O – H bonds are able to form strong hydrogen bonds. But alkanoic acids, due to their –COOH group, have the ability to form 2 hydrogen bonds, creating even higher melting and boiling points as more energy is needed to break 2 bonds than 1.
Identify esterification as the reaction between an acid and an alkanol and describe, using equations examples of esterification
Esters (alkyl alkanoates) are formed in a condensation reaction between alkanols and alkanoic acids. i.e. alkanol + alkanoic acid → ester (alkyl alkanoate) + water
e.g. ethanol + propanoic acid → ethyl propanoate + water
Identify the IUPAC nomenclature for describing the esters produced by reactions of straight-chained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8
Describe the purpose of using acid in esterification for catalysis
As esterification is a moderately slow reaction, a few drops of concentrated sulfuric acid is added to speed up the rate of reaction. Sulfuric acid is a dehydrating agent so absorbs the water produced in the reaction
which shifts the equilibrium to the right, thus producing more ester.
Explain the need for refluxing during esterification
Refluxing is the process in which any alcohol vapour, which rises in the heat of the reaction, is condensed using a water cooling condenser to prevent the loss of reactants or products. Thus, the reaction is able to be heated to a higher temperature → pushes the equilibrium to the right (as reaction is endothermic) → produces more ester.
Number of Carbons
Alkyl Part Alkanoate Part
1 Methyl Methanoate 2 Ethyl Ethanoate 3 Propyl Propanoate 4 Butyl Butanoate 5 Pentyl Pentanoate 6 Hexyl Hexanoate 7 Heptyl Heptanoate 8 Octyl Octanoate
+
↔
+ H
2O
27
Refluxing provides a safe alternative to performing the reaction in a closer container, which leads to a toxic build up of gases and possibly an explosion.
Outline some examples of the occurrence, production and uses of esters
Process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics
Esters naturally occur in fruits and flowers and are responsible for their taste and scent. Industrially, esters are artificially produced for practical uses, e.g. ethyl ethanoate is a solvent in nail polish remover; the sweet scent is replicated for use in perfumes, such as apple or pear flavours. Producing esters industrially is cheaper than extracting them from their natural sources.
Identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux
Equipment:
15 mL of acetic (ethanoic) acid 15 mL of ethanol
Round bottom flask with water cooling condenser
3 boiling chips
Concentrated sulfuric acid 500 mL beaker
Wire gauze Tripod
Bunsen burner Method:
1. Place 15 mL of ethanol and 15 mL of acetic acid into the round bottom flask. Add 3 boiling chips (to ensure even boiling) and 10 drops of concentrated sulfuric acid.
2. Heat the mixture using reflux apparatus for 20 to 30 minutes, until 2 layers are visible. Allow to cool.
3. Add 100 mL of water to the mixture and shake. Allow the 2 layers to separate again.
4. Run this mixture through a separating funnel. Discard the aqueous layer.
5. Add 50 mL of Na2CO3 solution to the remaining mixture in the
separating funnel. Shake gently, expelling gas that evolves. Allow the 2 layers to separate again.
6. Run the mixture through the separating funnel again. Discard the lower layer.
7. Repeat steps 5 and 6 to leave a pure sample of the ester, ethyl ethanoate.
28
9.4.1 Much of the work of chemists involves monitoring the reactants and products of
reactions and managing reaction conditions
Outline the role of a chemist employed in a named industry or enterprise, identifying the branch of chemistry undertaken by the chemist and explaining a chemical principle that the chemist uses
Gather, process and present information from secondary sources about the work of practising scientists identifying:
the variety of chemical occupations
a specific chemical occupation for a more detailed study
An analytical chemist working for Sydney Water will use AAS to monitor concentrations of metals, such as Pb and Hg, in water samples which eventually will be supplied to households. The principle of
chemistry used by the scientist is that metal electrons move from one energy shell to a higher one by absorbing electromagnetic radiation of a wavelength specific to that metal, giving each metal a unique emission spectrum metals can be identified and their concentration can be determined even in the presence of other metals. It is important that the scientist monitor and regulate the concentrations of metals in drinking water sources as excessive concentrations can lead to poisoning and illness, e.g. lead poisoning can cause mental retardation, but trace element metals, like Zn, are essential in small
concentrations for body functioning.
Identify the need for collaboration between chemists as they collect and analyse data
Chemists specialising in various fields of chemistry, e.g. analytical, organic, industrial, need to collaborate to solve problems requiring a broad spectrum of chemical knowledge.
Chemists need to regularly exchange viewpoints and have an open minded but critical approach to ensure that our scientific knowledge is constantly improving.
Describe an example of a chemical reaction such as combustion, where reactants form different products under different conditions and thus would need monitoring
Combustion reactions can produce solely carbon dioxide or a mixture of poisonous carbon monoxide and pollutant soot (C), depending on the amount of oxygen present. Using methane as an example:
Oxygen Supply Bunsen Burner Hole Flame Colour
Rate of Combustion Reaction Equation
Sufficient Open Blue Complete combustion;
maximum energy released
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)
Limited Partially
open
Mauve Incomplete combustion 2CH4 (g) + 3O2 (g) → 2CO (g) + 4H2O (l)
29
9.4.2 Chemical processes in industry require monitoring and management to maximise
production
Identify and describe the industrial uses of ammonia
Fertilisers (ammonium nitrate) – important for growing food and crops for growing populations; ammonia is reacted with nitric acid to form ammonium nitrate i.e. NH3 + HNO3 → NH4NO3
Explosives – during WWI, Germany called for more explosives
Identify that ammonia can be synthesised from its component gases, nitrogen and hydrogen
Hydrogen + atmospheric nitrogen ↔ ammonia i.e. 3H2 (g) + N2 (g) ↔ 2NH3 (g) Identify the reaction of hydrogen with nitrogen as exothermic
∆H = -92 kJ mol-1 (exothermic)
Describe that synthesis of ammonia occurs as a reversible reaction that will reach equilibrium
The synthesis of ammonia is a reversible reaction, meaning that ammonia is formed at the same time as it is being decomposed. Equilibrium is reached when the synthesis (forward) reaction proceeds at the same rate as the decomposition (reverse) reaction.
Explain why the rate of reaction is increased by higher temperatures
Increasing reactants’ temperatures → increase in their kinetic energy → brings them closer to their activation energy →higher chance of successful collisions → increased reaction rate
Explain why the yield of product in the Haber process is reduced at higher temperatures using Le Chatelier’s principle
By Le Chatelier’s principle, if the temperature is increased, the equilibrium will shift to the side that uses up the heat. Since the synthesis (forward) reaction of ammonia is exothermic (heat is produced), the equilibrium will shift to the left to oppose the change→ decomposition of ammonia → reduced yield of ammonia.
Explain why the Haber process is based on a delicate balancing act involving reaction energy, reaction rate and equilibrium
Increased temperatures → reduced yield but also an increased reaction rate a balanced, compromised temperature is needed to maximise yield and have a moderate reaction rate
Explain that the use of a catalyst will lower the reaction temperature required and identify the catalyst(s) used in the Haber process
The addition of an iron catalyst, such as magnetite Fe3O4, lowers the activation energy, which enables a
faster reaction rate. Thus, the reaction proceeds at a moderate rate at lower temperatures. The catalyst has no effect on temperature so does not affect the equilibrium.
Analyse the impact of increased pressure on the system involved in the Haber process
A high pressure (250 × atmospheric pressure) increases both the reaction rate and the yield. Higher pressure → more moles of gas in the closed container → increased chance of successful collision → faster reaction rate
Higher pressure → shifts equilibrium to the right (as least gaseous molecules are produced on this side) → more ammonia (increased yield)
30 Explain why monitoring of the reaction vessel used in the Haber process is crucial and discuss the
monitoring required
Temperatures need to be monitored as excessively high temperatures permanently damage the catalyst and low temperatures compromise the optimum yield and reaction rate
Pressure needs to be monitored as high pressure could be explosive
Incoming gases lower the reaction’s efficiency and in the case of oxygen, could lead to explosions if it reacts with hydrogen gas
Gather and process information from secondary sources to describe the conditions under which Haber developed the industrial synthesis of ammonia and evaluate its significance at that time in world history
Haber (and Bosch) discovered that a temperature of 400°C, high pressure of 20MPa and an iron catalyst were the ideal conditions needed to optimise the synthesis of ammonia from hydrogen and nitrogen.
Significance:
The synthesis of ammonia was necessary for the production of explosives for Germany in WWI after the British cut off Chilean guano (bird dropping) supplies
Haber’s contributions helped Germany’s war efforts, and even prolonged the war
The synthesis of ammonia facilitated the manufacture of fertilisers for food production for growing populations
31
9.4.3 Manufactured products, including food, drugs and household chemicals, are
analysed to determine or ensure their chemical composition
Deduce the ions present in a sample from the results of tests
Perform a first hand investigation to carry out a range of tests, including flame tests, to identify the following ions: Phosphate Sulfate Carbonate Chloride Barium Calcium Lead Copper Iron Unknown Cations Ba2+, Ca2+, Pb2+, Cu2+, Fe3+, Fe2+ Add Cl- by adding dilute 1 mol L-1 HCl White precipitate forms No precipitate forms Cation: Pb2+ Add H2SO4 No precipitate
forms White precipitate
forms Cations: Cu2+, Fe3+, Fe2+ Cations: Ba2+, Ca2+ Add NaOH Conduct flame test Brown precipitate forms → Fe3+ Grey green precipitate forms → Fe2+ Blue precipitate forms → Cu2+ Apple green flame → Ba2+ Orange red flame → Ca2+
32 Gather, process and present information to describe and explain evidence for the need to monitor levels of one of the above ions in substances used in society
Phosphate
Low concentrations are essential in waterways for normal aquatic plant growth
High phosphate concentrations → algal bloom → algae grows to cover lake/river surface → kills fish + uses up all the phosphate → then, algae decays anaerobically → low oxygen levels in water → death of marine life in the water
Monitoring the amount of phosphate entering and already present in waterways, scientists can guard against algal blooms
Unknown Anions CO32-, SO42-, PO43-, Cl-
Add 2 mol L-1 HNO3 Bubbles seen No bubbles Silver chloride precipitate forms → Cl-
Anion: CO32- Acidify solution with HNO3
and add Ba(NO3)2
Anions: PO43-, Cl- No precipitate forms White/ pale blue precipitate forms Anion: SO42-
Add NH3 → makes solution basic, then add aqueous
Ba(NO3)2
White precipitate forms → PO43-
Acidify by adding HNO3, then add aqueous AgNO3
