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

Acid Base/AlkaliProduces hydronium or H+ ions in solution Contains O²ˉ OHˉ or produces the hydroxide ion in solution

- have sour taste- corrosive- conduct electricity- turns blue litmus red (BAR)

- soapy feel- bitter taste- good conductors of electricity- red litmus blue

Strong Acids:- HCl- HBr (hydrobromic acid)- HI (hydroiodic acid)- H2SO4

- HNO3

- HClO3

- HClO4

Strong Bases:- NaOH (sodium hydroxide)- KOH- LiOH- RbOH- CsOH- Ca(OH)2

- Sr(OH)2

- Ba(OH)2

Weak Acids:- CH3COOH (acetic acid)- C6H8O6 (ascorbic acid)- 2-hydroxypropane-1,2,3-tricarboxylic acid

(Citric acid)- HCN (hydrocyanic)- Phosphoric acid (H3PO4)- Hypochlorous acid (HOCl)- HNO2 (nitrous acid)- HF- H2S- CCl3COOH (trichloroacetic)- CH3NH2

Weak bases:- NH3 (ammonia)- NH4OH (ammonium hydroxide)- N(CH3)3 (Trimethyl ammonia)- C5H5N (Pyridine)

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

Acidic Neutral Basic

- lactic acid (milk)- car battery acid- carbonated soft drinks

(carbonic acid)- vinegar - fruit juices- ascorbic acid (vitamin C)

- water- lactose solution- salt water- glucose solution- alcohol mixed with H2O

- ammonia solution- oven cleaners- washing soda solution- limewater- baking soda solution- NaOH (caustic soda)

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: substance which changes colour in solution depending on how acidic or basic the solution is.

Methyl orange: 3.1 – 4.4

Bromothymol blue: 6 - 7.6

Litmus: 5.5 – 8

Phenolphthalein: 8.2 - 10

identify and describe some everyday uses of indicators including the testing of soil acidity/basicityIndicators are cheap and convenient for:

- testing pH of soils: some plants need more acidic soils then others to grow--> used in agricultureo if soil is too acidic adjustments are made by adding calcium oxide (lime)o soil sample is obtained and diluted with pure watero solution is filteredo an indicator is added

- testing home swimming pools : pool needs to be approximately neutral adding chemicals for sanitation purposes changes its pH balance

o if pH is too low, skin irritant, eye irritanto range pH: 7.2 – 7.6o use phenol red collect water sample and add indicator

- monitoring wastes in labs: wastes are discharged off into sewerage ways --> sewerage ways lead out to the ocean therefore wastes pH should be measured and neutralised

o domestic waste and industrial waste water is tested to ensure that it isn’t too acidic so that it wont corrode sinks, drains and sewerage pipes.

- Hydrangea grown in basic soils are pink In acidic soils are blue

perform a first-hand investigation to prepare and test a natural indicatorAim: to produce an indicator using a natural substanceMethod:

1. A few red cabbage leaves were crushed using a mortar and pestle and then boiled over a blue flame in a 500mL beaker with 200mL of water.

2. When the colour of the cabbage leaves became pale indicator the transfer of its substance to the water, it was taken off the Bunsen flame and left to cool (if solution is too pale, the cabbage will need to be boiled for longer to concentrate the solution)

3. Two 2mL samples of NaOH was measured and placed into separate test tubes. This step was repeated for all other substances for testing while solids were added to a solvent (i.e. water).

4. Three drops of red cabbage indictor was added to the solution in one of its the test tubes and the colour change (if any) observed was recorded.

5. The second solution in test tube was tested for its pH using an universal indicator. Three drops of universal indicator solution was dropped using a pipette into the test tube while the colour change was compared to the indicator chart.

6. The pH of the substance was then matched up to the colour change from addition of the red cabbage indicator. This comparison was used to generalise a trend in colour change of red cabbage indicator in acidic and basic substances.

Results:Substance Red Cabbage Indicator Colour Universal indicator Colour Acidic/Basic/Neutral

Trends: red cabbage turns - green/blue for basic substances - purple for neutral- pink for acidic substances.

Discussion:We can compare our results with standard results. If there is an inconsistency between the two results this could be explained by the different solutions used to be tested. While the experiment carried out used common household acids and bases, the acids and bases used to obtain the colour changes above were more accurate (pH was exactly the number written instead of belonging to a pH range).

Risk assessment:- use heat mats + gloves + tongs (to remove wire gauze) and safety goggles for testing- Apron should be worn (red cabbage can stain) - Read labels of safety precautions heed safety warnings (don’t test CAUSTIC or POISON marked substances)- Operator awareness + monitoring required- Rinse after contact with substances

- Use a fume hold or have good ventilation

Conclusion: an indicator can be produced using a natural substance while its own colour changes occur for different pH ranges.

2. While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulphur. 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- the higher the oxidation state of a non-metal, the more acidic the oxide - non-metal oxides dissolve in water to form acids except for the neutral oxides

Acidic Oxides – NON METALS

Basic Oxides - METALS Amphoteric Oxides Neutral oxides

Covalent compounds Ionic compounds ZnO, PbO, Al2O3 CO, N2O, NO

Reacts with water to form an acid

Reacts with water to form a base

Behaves as both a BL acid and Bl base

Formed from non-metals: C and N

Reacts with bases to form salts vice versa

Reacts with acids to form salts but not with other alkalis

Reacts with both acids and bases to form a salt -> i.e. neutralises

Doesn’t react with both acids and bases

Covalent compounds Ionic compounds Al, Zn, Be, Sn, Pb form amphoteric oxides

C and N form neutral oxides

Neutralisation reaction:Acid + base salt + H2O

Eg. HCl(aq) + NaOH(aq) NaCl(aq) + H2OThe net ionic will be:

H+(aq) + OH-

(aq) H2O(l)

Neutralisation reactions generally produce water as well as salts but there are exceptions:Eg. H2SO4 (aq) + 2NH3 (aq) (NH4)2SO4 (aq)

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

Elements on the left of the periodic table are more metallic and less electronegative. These oxides are more basic

decreasing electronegativity more metallic more basic

Elements on the right of the periodic table are more non-metallic) hence has a higher electronegativity. Therefore their oxides are more acidic.

Increasing electronegativity More non-metallic

More acidic

define Le Chatelier’s principle

Le Chatelier’s Principle“When a system of equilibrium is disturbed, it will shift in such a direction to counteract the change.”

This is used to predict how a system of equilibrium will respond to a disturbance.

Chemical equilibrium: state of reversible reaction where the rate of forward reaction is equal to the rate of the backward reaction.

identify factors which can affect the equilibrium in a reversible reaction

To obtain a chemical equilibrium:- Reaction must be reversible- Reaction is performed in a closed container- Sufficient time must be allowed to reach a chemical equilibrium

Position of equilibrium: the dominant direction of the reversible reaction

Factor Achieved byPressure or concentration Adding more molecules

Temperature changing the temperature by adding or removing heat this affects only exothermic and endothermic equilibrium systems

Volume changing the size of the system

Adding a catalyst: - Has no effect on the final position of a chemical equilibrium- The same amounts of reactants and products are obtained for the same reaction at the same temp condition- Helps to reduce the time taken for a reversible reaction to reach the state of equilibrium by lowering the activation

energy for both forward and backward reactions.

Possible changes to a chemical equilibrium:- Concentration:

o The reaction shifts to the opposite side of the change in concentrationo Eg H2 + I2 2HI(g); if h2 increases, the reaction shifts to the right hand side to favour the production of HI and

counteract the change in equilibriumo When there is an addition of a pure solid and/or pure liquid there is no shift in of the chemical equilibrium

- Pressure/volumeo Pressure is proportional to the number of moles of gaseous moleculeso If there is an increase in pressure in a vessel, there is a decrease in volume to counteract the change o If pressure decreases, volume increases

- Temperatureo Depends on whether the reaction is exothermic or endothermico Decreasing temperature of an exothermic reaction causes the equilibrium to shift to the right to release heato Increasing temperature of an endothermic reaction causes the equilibrium to shift to the left to absorb heat

- Catalysto Catalysts have no effect on the final position of a chemical equilibrium rates of forward/backward reactions

are increased equalllyo Lowers activation energy of the forward/backward reactionso Reduces time take for the reversible reaction to reach equilibrium

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 solubility of carbon dioxide gas in water can be fully described using four equilibrium equations: - CO2(g) CO2(aq) EXOTHERMIC

- H2O(l) + CO2(g) H2CO3(aq)

- H2CO3(aq) H+(aq) + HCO3-(aq)

- HCO3-(aq) H+(aq) + CO3

2-(aq)

There is a chemical equilibrium between the reversible reactions of carbon dioxide gas into aqueous and aqueous into gas.

Fizzing in a drink- When lid is opened

o Pressure decreases (CO2 (g) escapes into the atmosphere)o Equilibrium shifts to LHS to favour production of CO2 (g)

- Place in a refrigeratoro Fizzing of drink is lesso Temperature is decreased equilibrium shifts to the RHS (since exothermic)

o i.e. favours production of aqueous CO2 (CO2(g) dissolves into CO2(aq))

Le Chatelier's principle predicts that: - addition of acid (increased concentration of H+) shifts equilibrium to the left - addition of base (OH- reacts with and reduces concentration of H+) shifts equilibrium to the right - addition of a soluble carbonate (increased concentration of CO3

2-) shifts equilibrium to the left

Example: Use Lechatelier’s Principle to relate the increase in burning fossil fuels to a possible increase in the acidity of the oceans

The increasing in burning fossil fuels leads to an increase of concentration of SO2 gas in the atmosphere. This leads to the formation of acidic rain when SO2 dissolves in the water vapour in the atmosphere. The formation of acid rain due to SO2:

H2O (g) + SO2(g) H+(aq)+ HSO3

-(aq)

The acid rain falling into the ocean leads to an increase it its acidity.In terms of Lechatelier’s principle, the increasing SO2(g) causes equilibrium to shift to the right., producing H+ which increases acidity hence acidity of oceans increases.

identify natural and industrial sources of sulfur dioxide and oxides of nitrogenNatural Sources Industrial Sources

Sulfur dioxide

- geothermal hot springs- volcanic activity

- smelting of sulphur ores in the extraction of metals process

- burning of fossil fuels such as coal which contains sulphur

Nitric Oxide

- lightning activity generates high temperatures to allow oxygen and nitrogen to combine

burning of fossil fuels in transport vehicles (cars, trucks etc) and power stations ( for generation of electricity)Nitrogen

DioxideNitric oxide combines with oxygen in the presence of lightning

Dinitrogen Oxide / nitrous oxide

Bacteria in nitrogenous material in soils

use of nitrogenous fertiliser in agriculture which provides raw material for the bacteria in the soil to feed off and produce dinitrogen oxide

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

EffectsSulphur dioxide

- irritates the respiratory system- causes breathing difficulties even at low concentrations- unpleasant odour- causes formation of acid rain

Nitrogen Dioxide

- irritates respiratory tracts- causes breathing discomfort at concentrations of 3 to 5 ppm (higher concentrations does

extensive tissue damage)- sunlight acts on nitrogen dioxide in the presence of hydrocarbons and oxygen to form ozone

NO2 + sunlight NO + OO + O2 O3

- ozone aggravates respiratory problems eg asthma- contributes to global warming rise in sea level adverse changes to the biosphere- pollutes atmosphere and increases the chances of the formation of photochemical smog

visually unattractive an a health hazard- heath effects for the aged with respiratory weaknesses- health effects money spent on medication and visits to doctors/ surgery for severe cases

Analysis of secondary sourcesSecondary sources available:

- internet search engines - newspaper- scientific magazines / journals- books- CD-ROMs

Evaluate:- Usefulness- accessibility- validity- reliability- accuracy

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

Process ReactionSulfur dioxide

Smelting of sulphide ores 2 ZnS (s) + 3O2 (g) 2ZnO (s ) + 2SO2(g)

Combustion of coal containing sulphur S + O2 (g) SO2 (g)

Nitric Oxide

Nitrogen combining with oxygen in the atmosphere

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

Nitrogen Dioxide

Nitric oxide combining with oxygen in the atmosphere

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

assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogenDirect evidence for increases in atmospheric oxides is more difficult to measure:

- Sulphur dioxide is naturally released into the atmosphere by volcanic activity and geothermal hot springs. Since volcanic activity varies year to year, it is difficult gain an accurate measure of the increase in atmospheric sulphur dioxide due to industrial sources.

- Instruments to measure very low concentrations have only been commercially available since the 1970s so there isn’t accurate data before then

Indirect evidence for increases of atmospheric oxides:- A decrease in pH in lakes and rivers- Corrosion of buildings and statues- Formation of photochemical smog- Damage to pine forests- Antipollution measures have been employed to reduce industrialised output of SO2

o The government have input measures to reduce the sulphur content in fuelso Catalytic converters are made compulsory for new cars to reduce the oxides of nitrogen released

Evidence of atmospheric oxides is the effects of acid rain are more difficult to observe directly To have a valid detection of acid rain due to the oxides of sulphur and nitrogen, long term data must be collected and compared. It is difficult to measure the increase in these concentrations using direct evidence however indirect evidence is the best alternative more practical

calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0˚C and 100kPa or 25˚C and 100kPa- equal numbers of molecules of different gases occupy the same volume (at same temp and pressure)- to calculate volumes

1. find number of moles produced2. find number of moles reacted3. find volume by multiplying number of moles with the volumes of gases at the corresponding pressure

explain the formation and effects of acid rainOxide Acid formed ReactionCarbon dioxide Carbonic acid CO2(g)+ H2O(l) H2CO3(aq)

Sulfur dioxide Sulphurous acid SO2 (g) + H2O(l)) H2SO3(aq) --> Sulfuric acid 2H2SO3(aq) + O2(g) H2SO4(aq)

Nitrogen dioxide Nitrous and nitric acid

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

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

HNO3(aq)+ H2O(l) H3O+ + NO3-

Emissions of pollutants such as SO2 (sulphur dioxide) and the oxides of nitrogen lead to the formation of acid rain. In the atmosphere, when water vapour condenses to form rain drops, the pollutants are dissolved in the process which forms acid rain. The oxides having dissolve in water turn into acids.

Acid rain causes:- increased acidity in lakes and rivers : affects fish populations (when pH<5 aquatic animals die) and the river

ecosystem- damage to plants and hinders their growth

o eg pine forests affects agricultural and pine industry- erosion of marble and limestone structures requires funds for repairing- severe damage to vegetation--> protective waxes are lost on leaves causing damage to plants- increase in pH of soil inhibits plants from absorbing sufficient nutrients such as calcium or potassium- corrosion of metals- limestone buildings and monuments will crumble when attacked by acid rain:CaCO3 (s) + H2SO4 (aq) CaSO4(s) + CO2 (g) + H2O (l)

Rain which has not been subjected to pollutants is still slightly acidic. This is due to the presence of carbon dioxide in the atmosphere which dissolves in water to form carbonic acid.

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

Example Explain why acid rain has not been as severe in Australia as Europe/USA.Australia has a relatively small population meaning there are less cars and appliances which emit gases which cause acid rain. Also because Australia is surrounded by oceans, it is protected from the atmospheric pollution of other countries.Major cities of Australia are also on the coast which helps disperse pollution produced by the industry. Our coal and oil supplies also have relatively low sulphur content

Demonstrate the effect of more carbon dioxide on the pH of oceansCarbon dioxide dissolves in water to form carbonic acid which ionises to produce H+.

CO2(g) CO2(aq) EXOTHERMIC

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

H2CO3(aq) H+(aq) + HCO3-(aq)

A greater amount of carbon dioxide in the ocean increases the production of carbonic acid because according to Le Chatelier’s principle, when an equilibrium system is disturbed, the system will react in such a way to counteract the change. Therefore an increase in CO2 concentration drives the system to the right, increasing hydrogen ion concentration and lowering the pH.Also, with an increase of carbon dioxide emissions more acid rain formed from carbon dioxide is produced. This falls into the oceans and ionises to produce hydronium ions.

Outline problems that could arise from these oceanic changesThe decrease of pH in the ocean water would have detrimental effects on aquatic organisms as the lower pH would prevent the fish eggs from hatching and may affect the growth of some marine organisms which are sensitive to changes in acidity. Furthermore, the acidic conditions result in reduced reproduction which leads to a drop in fish numbers hence affecting the food chain. It also increases the rate of corrosion of ships. Shells of invertebrates and corals contain CaCO3 which dissolves under acidic conditions as shown in the following equation:

CaCO3 + H+ Ca 2+ + HCO3-

Describe a process which might offset these effects by removing carbon dioxide from the oceans.- Global warming: since CO2(g) CO2 (aq) is exothermic, by increasing the temperature, the reaction will move to the

left by Le Chatellier’s principle. This will remove the aqueous carbon dioxide form the ocean to form carbon dioxide gas.

- Photosynthesis of plants reduces concentration of dissolved CO2 hence raising the pH of the ocean6CO2 (g) + 6H2O(l) C6H12O6 (aq)+ + 6O2 (g)

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

Aim: To determine the amount of CO2 gas dissolved in a soft drink (carbonated drink)Method:

1. A 500 mL beaker was weight using a balance. Then a 300mL bottle of coke was weighed before opening and then was completely emptied into the 500mL beaker. This was then weighed again. The measurements were recorded.

2. The beaker holding the carbonated drink was heated using a low heating flame. It was set up using a tripod above a Bunsen burner with the beaker sifting on top of a wire gauze on the tripod.

3. The beaker was gently warmed using a heating flame, preferably an electronic heater. Before the solution began to boil the flame/heater was removed to prevent evaporation of the carbonated drink.

4. The beaker was allowed to cool and then it was reweighed on the balance. This measurement was recorded and used to calculate the loss in CO2 (g) from solution by subtracting it from the initial mass. Using this measurement the number of moles of CO2 released was calculated by substituting values into the equation n=m/M. This quantity was converted into litres at 250C and 101.3 kPa using the molar volume.

Risk Assessment: The experiment involves the heating of a carbonated drink to speed up the release of gas from its aqeous form.

- To prevent spilling planning and careful monitoring while solution is heated.

- Wear goggles when observing beaker it might fix and “spit”- Use heat mats and tongs and gloves to handle hot instruments.

Calculations required to determine the volume of gas released Loss of mass due to escape of carbon dioxide gas. (final mass – initial mass) Conversion of grams of CO2 lost to moles of CO2. Number of CO2 released Volume of CO2

Discussion:- Differences between standard values and experimental results due to evaporating- Methods to improve reliability repetition - Methods to improve accuracy use electronic balance- Some remnants of coke could be remaining in its packaging bottle affecting the results- Discuss reason for heating solution

o A gradual increase in heat prevents the carbonated drink to foam or spray out of the beakero By Le Chatelier’s principle, the reaction will favor the release of CO2 (g) in order to absorb more heat (i.e.

decrease the temperature).o This speeds up the experiment

- Using a hot plate would provide a very gentle heat and help heat the solution without evaporation

OTHER Methods:- warming method: uses the drop in water solubility of gas on heating- salting method: uses the stronger attraction between soluble salt ions with water than dissolved gas molecules have

for water.- Shaking the bottle of carbonated drink and open lif to reduce pressure and in turn release carbon dioxide + repetition.

NB:- Do not spill small amounts of spilt carbonated drink results in a large percentage error- To improve accuracy use electronic balance- Remnants of coke could be remaining in its packaging bottle affecting the results- To improve validity and reliability cover the beaker while it is cooling (to prevent evaporation)

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

define acids as proton donors and describe the ionisation of acids in waterAcids are formed by metal oxides which ionise in water by donating a proton to the water molecule, forming the metal aqueous and a hydronium ion. When acid molecule is placed in water it can IONISE; releasing a proton (hydrogen ion) and forming a negative ion. Free H+ ions do not exist in aqueous solution, therefore it is donated to the water molecule and joins to form

hydronium

NB: A hydrogen atom, H, consists of one proton and one electron. A hydrogen ion, H+ is formed when an H loses its electron, leaving just a proton. A proton and a hydrogen ion are thus the same and can be represented by H+.

- When an acid molecule is placed in water, it ionizes by releasing a proton and forming a negative ion. - The proton, H+, can attach to a water molecule, H2O, forming what is called a hydrated hydrogen ion or

hydronium ion, H3O+

gather and process information from secondary sources to write ionic equations to represent the ionisation of acidsIonisation of an acid

identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic), hydrochloric and sulfuric acid- Sulfuric acid is called a diprotic acid because each molecule can release up to two protons.

- Phosphoric acid is called a triprotic acid because each molecule can release up to three protons.

- Acetic acid occurs naturally in the decomposition of biological material, eg by oxidation of ethanol it is found in wine alcohol. Most acetic acid used by humanity is manufactured industrially.

- Citric acid occurs naturally in fruits but most added to food as preservatives is manufactured.

Hydrochloric and sulfuric acid are produced industrially on a large scale. Most sulfuric acid is manufactured, but it can also occur naturally. Sulfuric acid is produced by glands in the lining of our stomachs helps enzymes to break complex food molecules For example, most sulfur dioxide released into the earth’s atmosphere is oxidized and dissolved in water to form the

sulfuric acid in acid rain. If the acid rain results from volcanic eruption it could be regarded as natural, but if acid rain results from smelting of sulfide ores, it could be regarded as manufactured.

describe the use of the pH scale in comparing acids and basesThe more pH scale ranges from 1 to 14 according to the concentration of hydrogen ions present in the substance. Acids have a pH ranging between 0 and 7 ; a neutral solution has pH=7 ; a basic/alkaline has pH>7

describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute strength of an acid refers to the degree of ionisation it has

a weak acid doesn’t fully ionise when dissolved in water (backward reaction is more successful than forward reaction--> ions react easily to reform the acid and the water)

a strong acid has a degree of ionisation of 100%

when an acid partly ionises it means it only partly converts to ions while the rest of the acid remains as molecules concentration refers to the amount of acid per unit volume

identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold change in [H+]

pH is a measure of the concentration of hydrogen ions in a solution. Strong acids like hydrochloric acid are the sort of acids you normally use in the lab have a pH around 0 to 1.

NB: The lower the pH, the higher the concentration of hydrogen ions in the solution.

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

- Hydrochloric acid completely ionises in water, therefore is a very strong acid.- Citric acid however is a weak acid as it only partially ionises in water to form fewer H+ ions.

o Tripotic up to 3 hydrogen ions are producedo Stronger than acetic acid

- Acetic acid is also a weak acid as it only partially ionises in water. Because of this the dissociated acid exists in equilibrium with its aqueous molecules.

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

A strong acid has a very high degree of ionisation meaning their molecules completely ionise in water. Because of this, as all the molecules of the acid are ionised, this reaction is not a reversible reaction since the forward reaction is completely dominant and there are no remaining reactants. (IONISATION REACTION GOES TO COMPLETION) independent of concentration

A weak acid has a low degree of ionisation meaning their molecules partially ionise in water. In this situation only a part of the acid is ionised, meaning the reaction can be reversible. (DOESNT GO TO COMPLETION EQUIILIBRIUM) depends on concentration

Electrical conductivity: ions moving in a liquid can conduct electricity. Electrical conductivity of an acid can be determined by applying a fixed DC voltage between the two electrodes and measuring the current flow. The more ions in solution the higher the electrical conductivity.

plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acidsExample of a strong acid problem: What is the pH of a 0.010 M HCl solution? -->Hint: HCl is a strong acid. The H3O+ concentration will equal the HCl concentration: ([H3O+] = 0.010 M) The pH is determined by taking the negative log of the [H3O+] concentration:

pH = -log[H3O+] = -log(0.010) = 2.00

Example strong base problem: What is the pH of a 0.010 M NaOH solution? -->Hint: NaOH is a strong base. The [OH-] concentration will equal the concentration of NaOH: [OH-] = 0.010 M pH is determined by calculating the pOH.

1. take the negative log of the [OH-] concentration2. convert the pH to pOH.

To calculate pOH: P[OH] = -log[OH-] = -log(0.010) = 2.00

The pH can now be determined by using the formula pH + pOH = 14: pH = 14.00 - 2.00 = 12.00

To find degree of ionisation:- concentration of hydrogen ions/ concentration of acid

gather and process information from secondary sources to explain the use of acids as food additivesAcetic acid and citric acid are used as food additives are they are both weak acids.Acids have the ability to prevent the growth of bacteria and foreign microorganisms. Used as a preservativeDue to its sour taste, acids can also be used for flavouring.

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

Naturally occurring acids- hydrochloric acid: produced by glands in stomachs to allow for the efficient operation of enzymes that break

complex food molecules into easily transportable small molecules that are absorbed into the bloodstream when they pass into the intestine

- acetic acid : acid is present in vinegar which is commonly made from wine by the oxidation of ethanol- citric acid: found in citrus fruits

- vitamin C (abscorbic acid)Naturally occurring bases

- ammonia : NH3

o present in the stale urine of humans and other animals manufactured in huge quantities for the production of fertilisers, nitric acid and other chemicals

- amine: compounds with some alkyl groups - metallic oxides- carbonates

Verbs to know

Account for State reasons for, report on, give an account of, narrate a series of events or transactions

Analyse Identify components and the relationships among them, draw out and relate implications

Apply Use, utilise or employ in a particular situation

Appreciate Make a judgement about the value of something

Assess Make a judgement of value, quality, outcomes, results or size

Calculate Determine from given facts, figures or information

Clarify Make clear or plain

Classify Arrange into classes, groups or categories

Compare Show how things are similar or different

Construct Make, build or put together arguments

Contrast Show how things are different or opposite

Critically (analyse/evaluate)

Add a degree of level of accuracy, depth, knowledge and understanding, logic, questioning, reflection and quality to an analysis or evaluation

Deduce Draw conclusions

Define State the meaning of and identify essential qualities

Demonstrate Show by example

Describe Provide characteristics and features

Discuss Identify issues and provide points for and against

Distinguish Recognise or note/indicate as being distinct or different from, note difference between things

Evaluate Make a judgement based on criteria

Examine Inquire into

Explain Relate cause and effect, make the relationship between things evident, provide why and/or how

Extract Choose relevant and/or appropriate details

Extrapolate Infer from what is known

Identify Recognise and name

Interpret Draw meaning from

Investigate Plan, inquire into and draw conclusions about

Justify Support and argument or conclusion

Outline Sketch in general terms; indicate the main features

Predict Suggest what may happen based on available information

Propose Put forward (a point of view or suggestion, argument etc) for consideration or action

Recall Present remembered ideas, facts or experiences

Recommend Provide reason in favour

Recount Retell a series of events

Summarise Express concisely the relevant details

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:o Lavoisiero Davyo Arrhenius

Changing Concepts of Acids

1776 – Antoine LavOiser: non-metal compounds containing oxygen produced acids when they dissolved in water. “Acids are substances which contain oxygen” DISPROVED many acids do not contain oxygen and many substances containing oxygen are basic

1815 – Sir Humphrey Davey: active metals which reacted with acids produced hydrogen gas. “Acids are substances which contain replaceable hydrogen” FLAWED only provided us with a way to recognise acids but didn’t explain the properties of acids. Also did not explain why many hydrogen compounds were not acidic.

1884 – Savant ArrH+enius : solutions of acids conducted electricity hence these acids are electrolytes and therefore contains ions. “Acids are substances which produce hydrogen ions in water”. Then he showed that an aqeous solution of a base conducted electricity. Hence the ions responsible for the conductivity of an alkaline solution were hydroxide ions. “ bases are substances which ionise water to produce hydroxide ions” important in explaining differences between strong and weak acids and the development of the pH scale

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

Lavoisier: was wrong but stimulated research into the composition of acids and created awareness of the need to define an acid.

Davy: defined an acid interms of its properties and reactions helped classify substances/did not interpret propertiesArrhenius: increased our understanding significantly interpreted acid properties in terms of the hydrogen ions they

produced- explained weak/strong acids in terms of the extent to which the ionisation reaction proceeded- increased our understanding considerably big step forward towards developing the concept of an acid- restricted definition of a base: a substance that produces hydroxide ions in solution bases can be insoluble

(do not produce hydroxide ions in solution)Bronsted-Lowry: increased our understanding further by showing acidity depends on both the structure of the substance as well as its properties relative to another reactant/its solvent.

- Showed neutralisation does not need to involve ionisation to H+ and OH- ions- Can proceed directly by proton transfer

outline the Brönsted-Lowry theory of acids and basesAcid: proton (hydrogen ion) donorBase: proton acceptorNeutralisation: transfer of protons

describe the relationship between an acid and its conjugate base and a base and its conjugate acidAn acid gives up a proton to form its conjugate base.

Acid + water --> H3O+ + conjugate baseNB: product Is called a base as it tends to gain a proton

A base accepts a proton to form its conjugate acidBase + water --> conjugate acid + OH-

NB: product is called a conjugate acid because it tends to give up a proton

identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic natureSalts: compound formed when an acid is neutralised by a base.A salt formed by a strong acid and strong base is neutral, eg NaClA salt formed from a strong acid and a weak base is acidic, eg NH4ClA salt formed from a weak acid and strong base is basic eg CH3COONa

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

Acid Base pH of saltStrong Strong =7Weak Strong >7Strong Weak <7Weak Weak Depends on which is strong

identify conjugate acid/base pairs- Whenever an acid and a base react, they form their conjugates:

HCl + H2O Cl- + H3O+ acid1 base2 conjugate base1 conjugate acid2

- Hydrochloric acid and chloride ion are a conjugate acid-base pair. - Water and hydronium ion are another conjugate acid-base pair

Relative strengths of acid-base conjugates:- conjugate base of a strong acid is an extremely weak base negligible- Conjugate base of a weak acid is a weak base- Conjugate acid of a strong base is an extremely weak acid

Acid conjugate base (acid loses proton/hydrogen ion to become base)Base conjugate acid (base gains proton/hydrogen ion to become acid)

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

amphiprotic: a substance which acts as a both a Bronsted-Lowry acid and base e.g. NH3, H2O, HCO3, HSO4, PH3

Water is an amphiprotic molecule: Water as an acid: H2O H+ + OH- or more fully H2O +H2O H3O+ + OH- Water as a base: H+ + H2O H3O+ or more fully H3O++H2O H2O + H3O+

The hydrogen carbonate (bicarbonate) ion is an amphiprotic ion: As an acid HCO3

- H+ + CO32- or more fully HCO3

-+H2O H3O+ + CO32-

As a base H+ + HCO3- H2CO3 or more fully H3O++ HCO3

- H2O +H2CO3

amphoteric: is able to NEUTRALISE both bases and acids

identify neutralisation as a proton transfer reaction which is exothermic Neutralisation involves the reaction between an acid and a base where the acid their hydrogen and hydroxide ions form water.

- It is an exothermic reaction meaning that heat is liberated as it occurs and that enthalpy change for the reaction is negative.

The net ionic equation of a neutralisation reaction shows it is a proton transfer reaction. That is, a proton from the acid transfers to the hydroxide ion of the base.

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

Neutralisation reactions are widely used for safety in laboratories and factories where acids or bases are used.

Because many acids and alkalis are very corrosive it is important to neutralise any spills of these substances quickly. In addition sewage authorities put strict limits on the pH of factory and laboratory effluents discharged to sewers so as not to upset bacterial breakdown.

Neutralisation reactions are widely used to ensure that effluents from such places are neither acidic nor alkaline.- Add a weak base (not corrosive) e.g. CaCO3 suitable for all types of spills- Neutralises spill to reduce the risk of corrosive action of acid- A substance containing an amphiprotic ion, such as the hydrogen carbonate ion in NaHCO3, is quite suitable for

neutralising chemical spills.

o If the chemical spill contains an acid, H+ + HCO3- H2O + CO2

If the spill contains a base, HCO3- + OH- CO3

2- + H2O.

o NaHCO3 is suitable for neutralising chemical spills of acids, bases and unknown acidity or basicity. it is also a weak base and has a weak charge attracts protons less vigorously

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

A buffer contains both a weak acid and its conjugate base. It is a solution which resists a change in pH when a small amount of strong acid or base is added.

This is because the weak acid and its conjugate base can react with both acid and base HA(aq) + OH-(aq) -> A-(aq) + H2O(l) A-(aq) + H+(aq) -> HA(aq)

NB: The conjugate base is added as a salt of the conjugate base: for example, to make an acetic acid buffer you would add acetic acid and sodium acetate to water.

A natural buffer system would be hydrogen carbonate in our blood.

Hydrogen carbonates ions help maintain the pH of blood to be about 7.4:

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

(aq)

When an acid is added, [H3O+] increases, this shifts the equilibrium to the left. This causes more HCO3+ ions to react with the

added H3O+ to form H2CO3. The solution thereby minimises the change in pH due to the added acid.When a base is added, [OH-] increases. These OH- ions react with the H3O+ ions to neutralise, which in turn increases the production of H2O (l). By Le Chatelier’s principle, the equilibrium then shifts to the right where the H2CO3 ionises more to form its HCO3

- and H3O+ ions. This in turn minimises the pH change due to the added base.

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

titration: volumetric analytical technique used to find the concentration of a solution by reacting it with a solution of an accurately known concentration.

Standard solution: has an accurately known concentration. There are two types of standard solutions- primary standard solution: made from highly pure solid used to titrate with another solution of unknown

concentration must be available in pure form chemical formula must be accurately known molecular mass should be relatively large to minimise weighing error must be soluble in water mass must not change when exposed to air

- secondary standard solution: the solution titrated by the primary standard solution reused for titration with other solutions.

Indicator: added to conical flask to signal equivalence point by changing colour suitable indicator is used depending on nature of acid-base reaction should change colour exactly or near the equivalence point of titration

End point: when indicator changes colourEquivalence point: when the base and acid has neutralised and the number of moles of the added species is

stoichiometrically equal to the original number of moles of the other species. OR it is when the solutions are mixed in exactly the right proportions according to the equation

Three important apparatus of titrationBurette Is filled with the standard solution. Must be rinsed with deionised water and then the

filling solutionPipette Transfers an accurate volume of the solution of unknown concentration to the conical

flask. Must be washed with deionised water and then the filling solutionIf solution needs to be diluted, it is transferred to the volumetric flask

Volumetric flask

Used to dilute a solution and measure an accurate volume of the primary standard. Rinsed with deionised water before using.

Conical flask Rinsed with de-ionised water filled with indicator and solution of unknown concentration

Steps:Preparing standard solution:

1. weigh the required mass of dried NaHCO3 directly into a clean dry beaker2. dissolve a small amount of distilled water3. add to a clean volumetric flask washed with distilled water4. rinse beaker using distilled water and pour remnants into volumetric flask5. fill the volumetric flask with distilled water to the graduated mark6. stopper the flask and invert 10 times

Conducting Titration:1. fill the burette with the standard solution and adjust solution level to the zero

mark2. add the solution of unknown concentration to the conical flask 3. add one or two drops of indicator to the conical flask4. open the valve of the burette slowly to allow the solution of known

concentration run slowly into the conical flask5. when the indicator changes colour stop the solution by closing the valve of the

burette6. read volume delivered by the burette as accurately as possible7. Calculate the concentration of the unknown solution using molarity calculations.

Titration of:- Strong acid – strong base

o Starts low Ends higho pH is close to 7 at equivalence point

bromothymol blue (6.2 – 7.6) should be used

- Strong acid – weak baseo Starts low ends lowo pH at equivalence point <7o change in pH not as much

methyl orange (3.1 – 4.4) should be used

- Weak acid – strong baseo Starts high ends high

o pH at equivalence point >7 phenolphthalein (8 – 10) should be used

Naturally occurring acids

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 compoundsalkanols: alkanes which have one or more H atoms replaced by OH (hydroxyl group)alkanoic acid: R-COOH

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

Esters are compounds formed when alkanoic acid reacts with alkanols. The general formula of esters: alkyl alkanoate

R – COO – R’ O R – C – O – R’

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

Esters Polar Has a highly electronegative oxygen and asymmetrical shapeIntermolecular forces: dipole- dipole and dispersion

Alkanol Polar short chainsNon-polar long chain

Has alcohol functional group (OH) displays hydrogen bonding

Alkanoic acids

Polar Has presence of OH and extra O (COOH) great extent of hydrogen bonds

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

ESTERIFICATIONAlkanoic acid + alkanol concentrated H2SO4 ester + water

Example:

HDYROLISIS OF ETHEREster + water diluted H2SO4 Alkanoic acid + alkanol

describe the purpose of using acid in esterification for catalysisThe catalyst used in esterification is concentrated H2SO4 (aq). It acts as a catalyst to increase the rate of reaction and by weakening the bonds in water hence lowering the activation energy. The concentrated sulphuric acid absorbs water and forces equilibrium to the right to increase the yield of ester.

explain the need for refluxing during esterificationRefluxing: the process of heating a reaction mixture in a vessel with a cooling condenser attached in order to prevent loss of

any volatile reactant or product

During esterification the components in the reaction mixture are heated where the volatile components then have a chance to escape. To counteract this, a condenser is placed on top of the reaction vessel so that any volatile components pass into the condenser. The condenser can be water or air-cooled and causes the volatile components to condense back to liquid and fall back into the reaction mixture.

Refluxing also improves the safety of the operation, as the volatile components are flammable identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux

- esterification should be carried out in a water bath rather than a naked flame- use an electronic heater- no stopper on reaction vessel- good ventilation use a fume cupboard to avoid inhaling vapours- a catalyst should be used to lower the heat required to be generated

outline some examples of the occurrence, production and uses of esters

- ethyl ethanoate : solver for paint and lacquers- butyl ethanoate: nail lacquers- phenylmethyl ethanoate : jasmine fragrance- ethyl methanoate : rum fragrance

process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics

Esters have :- pleasant fruity odours- occur widely in nature as perfumes and flavouring agents

http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.html

A system in dynamic equilibrium is a particular example of a system in a steady state. In a steady state the rate of inputs is equal to the rate of outputs so that the composition of the system is unchanging in time. For example, a lake is in a steady state when water flows in at the same rate as water flows out