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

Point 1.6 – Solve problems by applying information about the colour changes of indicators to classify some household substances as acidic, neutral or basic.

The colour changes of different indicators between certain pH ranges, can allow you to solve the problem of whether an unknown substance is acidic, neutral or basic. By knowing these pH ranges where colour changes occur we can predict the colour of the indicator. However for universal indicator there will be numerous colour changes. It will work across all pH ranges. For very acidic solutions it will show red then as it gets less acidic turn a yellow colour. When the solution is neutral, it will turn green and when the solution becomes increasingly basic, it will turn a blue/violet colour. Universal indicator can be used on substances of unknown pH and you will get quite accurate results, whereas with the other indicators listed above, you have to know whether it is an acidic, basic or neutral substance. A problem with using indicators to test the pH of a green detergent is that the colour of the detergent would mask any colour change of the indicator. This problem can be overcome by diluting the detergent or using a pH meter or pH probe.

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

Zn : H2

1 : 1

Point 2.9 – 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.

Example – Calculating volumes of gases given masses of some substances

Question: Calculate the volume of hydrogen evolved when 6.54g of zinc dissolves in excess sulfuric acid at 25ºC and 100kPa.

The first step is to right a balanced equation:

Zn (s) + H2SO4 (aq) ZnSO4 (aq) + H2 (g) The next step is to find the number of moles of zinc:

Number of moles = m/M= 6.54/65.41 (Molar mass of zinc)

= 0.10 moles of zinc.

The next step is to look at the mole ratio between zinc and hydrogen:

Since the mole ratio is 1:1, hydrogen has the same number of moles as zinc thus hydrogen has 0.10 moles.

K2CO3 : CO2

1 : 1

The next step is to use the equation: n = V/Vm

Rearrange the equation to make V (volume) the subject: V = (n) x (Vm).

Now substitute into formula: (NOTE: 1 mole of gas at 100kPa has a volume of 22.7 L at 273 K (0oC) or 24.8 L at 298 K (25oC).

V = (0.10) x (24.79) = 2.479 L = 2.48 L of hydrogen produced (3 significant figures)

Example – Calculating the masses of substances given gaseous volumes

Question: A sample of potassium carbonate was dissolved in excess 1.00mol L-1

hydrochloric acid and the evolved carbon dioxide collected and dried. Its volume at 0ºC/100kPa was 450mL. Calculate the mass of potassium carbonate used.

The first step is to right a balanced equation:

K2CO3 (s) + 2HCl (aq) 2KCl (aq) + CO2 (aq) + H2O (l)

The next step is to find the number of moles of carbon dioxide. Number of moles = V/Vm

= 0.45/22.71 = 0.0198 moles of carbon dioxide.

The next step is to look at the mole ratio between potassium carbonate and carbon dioxide.

Since the mole ratio is 1:1, potassium carbonate has the same number of moles as carbon dioxide thus potassium carbonate has 0.0198 moles.

The next step is to find the mass of potassium carbonate used by using the following formula: n = m/MRearrange the formula to make m (mass) the subject: m = (n) x (M)Now substitute:

m = (0.0198) x (138.21) [Molar mass of potassium carbonate] = 2.74 g of potassium carbonate used (3 significant figures)

Point 2.10 – Explain the formation and effects of acid rain.

Acid rain is formed when acidic oxides (such as carbon dioxide, sulfur dioxide and oxides of nitrogen among others) dissolve in water droplets in the atmosphere and precipitate out as acid rain. The acidic oxides form by the combustion of compounds in fossil fuel which contains carbon, sulfur and nitrogen. These acidic oxides form acids when they dissolve in rainwater:

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

[Sulfur dioxide + water sulfurous acid]

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

[Nitrogen dioxide + water nitric oxide + nitric acid]

Acid rain has many severe effects on society and the surrounding environment:

Increases weathering and erosion – many rocks crumble due to acid rain and this may mean that marble statues as well as buildings are attacked and crumble.

Increased attack on metals and alloys – Acid rain attacks structural metals (e.g. steel) in buildings and bridges.

Destruction of forests and vegetation – Plant leaves and roots are damaged by acid rain. Effect on waterways – The acid rain which is washed into the waterways will directly

affect aquatic organisms in that they will not be able to reproduce and it might even kill them. Fish eggs will not hatch.

Damage to soils – Acid rain can cause chemical reactions in the soil which in turn affect the vegetation.

Harm to humans – The sulfate particles present in the acid rain may cause bronchitis and asthma in humans.

Point 2.11 - Identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25˚C and 100kPa.

A soft drink can be decarbonated using the salting out method.

I will have to plan this experiment by trying to make this experiment as valid as possible. To do this I will have to figure out a way to successfully remove the carbon dioxide from a quantity of soft drink in a way that allows me to make accurate and reliable measurements. Data needed in this experiment include: weight of soda bottle, weight of soda bottle after decarbonation and the volume of the degassed drink. Calculations I will need include: loss of mass (due to CO2) and the conversion of grams of CO2 lost to moles of CO2. Method for the salting out method:

Obtain a fresh bottle of soda water and weigh using an electronic balance. Set up a control, same size and type of container containing same volume of water as

the volume of the soft drink and set aside for entire experiment. Weigh 5.00 g of salt and add into a 500mL beaker also weighed. Remove cap (of soda can) and add all the soda water slowly into the beaker. Don’t add it

too fast as liquid will escape in froth. Start timing with stopwatch and weigh beaker + contents every minute until you get 5

constant readings. Record results in a table Repeat experiment 10 times for reliable results.

Accuracy – More accurate weighing instruments e.g. to more decimal places. Validity – Using pure salt, whether all of the carbon dioxide has left the bottle. Use of a control. Reliability- Repeating experiment, comparing results with other groups and averaging results.

*** Find the volume of carbon dioxide which escaped by using the n = V/Vm formula.

Point 2.12 – 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.

Secondary sources from which you can obtain the information for this dot point include: internet web sites, science journals, reference books and school material.

Industrial production of sulfur dioxide: Combustion of fossil fuels in power plants

4FeS2 (s) + 11O2 (g) 2Fe2O3 (s) + 8SO2 (g)

Smelting of sulfide ores:

Industrial production of nitrogen oxides:

Lightning’s high voltage discharge makes nitrogen combine with oxygen to form nitrogen oxides.

High temperatures in internal combustion engines produce nitric acid:

N2 (g) + O2 (g) 2NO (g) Nitrogen monoxide produced by burning biomass.

Positives: sulfur dioxide used in the food industry as a food preservative in the food industry and oxides of nitrogen are used for fertilisers. The following are negatives of releasing these oxides:

When dissolved in rain, forms acid rain which can have severe consequences on the living organisms, buildings and society.

Oxides of nitrogen contribute to the formation of photochemical smog which is a serious problem in big cities especially in USA and Europe.

If these oxides are inhaled by humans, could cause severe illnesses such as asthma and bronchitis.

These oxides polluting the air could mean that animals such as birds might not be able to breathe freely and consequently they may die.

For sulfur dioxide and the oxides of nitrogen to be a huge concern in our daily lives, there needs to be evidence on their impact on the environment and society. If this is proven, it has indeed a huge impact on us today.

Overall the emission of sulfur dioxide and the oxides of nitrogen are of huge concern as they have a huge impact on the environment and society- especially when industry produces numerous tonnes a day of the oxides. Evidence has been gathered and it has proved the negative impact of these acidic oxides. Therefore more research needs to be undertaken to try and minimise this ever increasing problem in our lives today.

3. Acids occur in many foods, drinks and even within our stomachs. Point 3.6 – 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.

Let’s say we have 3 beakers. In one there is 0.1 mol L-1 of citric acid mixed with water, in another there is 0.1 mol L-1 of acetic acid mixed with water and in the last beaker there is 0.1 mol L -1 of hydrochloric acid mixed in water.

Degree of ionisation = [H+] / [HX] x 100 %

Where [H+] is the concentration of hydrogen ions and; [HX] = concentration of whole solution

Let’s say that HCl has a pH of 1.0, CH3COOH has a pH of 3.0 and citric acid 2.0. Using the 2nd boxed formula above we can figure out the [H+] of each of the acids. HCl = 0.1 mol L-1, CH3COOH = 0.01 mol L-1 and citric acid = 0.001 mol L-1. To find the degree of ionisation of an acid, we use the following equation:

Doing these calculations shows that the degrees of ionisation are: HCl – 100% CH3COOH – 1% ***These are not true values, just hypothetical Citric acid – 10%

From these results we can see that HCl is the easiest to ionise and therefore is willing to give away hydrogen ions (proton donor). The other 2 acids do not readily give away these hydrogen ions. Strong acids are completely ionised while weak ones are not. Therefore we can conclude that HCl is the strongest out of these acids, followed by citric acid and then acetic acid. (Citric acid has a lower pH than acetic acid because it is triprotic which means it releases H+ ions in 3 steps to acetic acid’s one step.

Strong acids are better conductors of electricity than weak acids as they release more ions which allows current to flow.

Point 3.8 – Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals.

Substances to test: HCl, H2SO4, NaCl, Na2CO3, NaHCO3, NH4OH, lemonade, orange juice, washing powder, rainwater, alcohol etc.

Put each of the solutions into a beaker Measure the pH of each using a pH meter or probe. Repeat experiment 10 times for reliable results.

Using a pH meter or probe is a non-destructive way of testing whether a chemical solution is acidic, basic or neutral. Using indicator solution or indicator paper (e.g. universal indicator) is a destructive way of testing, as the indicator will contaminate the portion of solution tested. You must clean the pH meter or probe with distilled water before use as any contamination may alter the pH values. Also before the actual experiment, test other substances of known pH (called buffers) to see if the pH meter or probe is functioning properly.

Accuracy – pH meter or probe is more accurate than an indicator as it provides an actual pH reading rather than a range in which a pH falls.

Reliability – Repeat experiment 10 times, compare results with other groups and average out results. Validity – Make sure the pH meter or probe is thoroughly cleaned with distilled water before use so that no contamination can alter the results.

Point 3.9 – Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids.

Before doing this experiment plan it in a way it is going to be valid. So therefore use the same concentration for the strong acid and a weak acid.

0.1 M solution of hydrochloric acid 0.1 M solution of acetic acid Add about 5 drops of each Put them in a test tube each. Use pH probe and data logger to measure pH of each acid. Repeat experiment 10 times for reliable results.

The lower the pH of the substance, the more acidic it is. Therefore it completely ionises in water and is a strong acid. Perform this experiment in a way that is safe. Wear safety goggles because acids and bases are corrosive and irritates eyes and skin. If spilt on skin, wash under cold water from tap.

Point 3.10 – Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids.

You can gather data from secondary sources such as the internet, scientific journals, books and school material. Process the data by picking out the main points and comparing the reliability of the data to other sources.

If an acid is a strong acid, the equation will usually be written with an arrow, , from left to right showing that ionisation of the acid molecules is almost complete. If an acid is a weak acid,

the equation will usually be written with the reversible arrows, , that show that significant amounts of reactants (un-ionised molecules) as well as products ( H+ and an acid anion) are present in equilibrium.

Ionisation of HCl: HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)

Ionisation of H2SO4: H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq)

Ionisation of acetic acid:

Point 3.11 – Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids.

Molecular model kits can be used to simulate the ionisation of strong and weak acids. Ionisation of strong acid – Using the model kit construct 4 HCl molecules by attaching balls representing hydrogen and chlorine with plastic bonds and construct 4 molecules of water. Now from each HCl molecule, remove the hydrogen atom and attach it to each water molecule so now there will be 4 hydronium ions and 4 chloride ions left over which signals complete ionisation. Ionisation of weak acid – Using the model kit, construct 4 CH3COOH (acetic acid) molecules and 4 water molecules. Remove only one hydrogen atom from one molecule of acetic acid and attach it to one water molecule. Now there should be 3 unionised acetic acid molecules and one

hydronium ion. This represents incomplete ionisation. (***Note: in organic acids, the H that ionises comes from a –COOH group e.g. acetic acid).

Point 3.12 – Gather and process information from secondary sources to explain the use of acids as food additives.

Secondary sources from which the information can be obtained from are: websites such as food standards Australia, Google, reference books, scientific journals and school material. Many acids are used as food additives:

Citric acid and tartaric acid are often added to jams to give a tart a sharp taste. The acidity and high sugar content inhibits the growth of microbes.

Leavening agents such as phosphoric acid and fumaric acid combine with NaHCO3 to produce CO2. They are present in many packaged biscuits and cake mixes.

Acetic acid in vinegar helps to preserve foods such as pickles and chutney by inhibiting the growth of bacteria and microbes.

Sulfur dioxide is used to prevent dried fruits such as apricots and white wine from being attacked by moulds or other microbes.

Ascorbic acid (vitamin C) is an antioxidant and is added to food to prevent spoilage. Phosphoric acid is also added to soft drinks to increase the sharpness of the drinks.

Point 3.13 – Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition.

Data includes the types of substances and their chemical composition. I will gather the data from internet sites, school material, reference books and scientific journals. I will process the data by picking out the main points (making sure the data is consistent throughout the sources) and presenting the data in a table.

Naturally occurring acid/base Acid present Chemical CompositionCitric juice Citric acid, ascorbic juice Organic compoundsVinegar Acetic acid CH3COOH and waterStomach acid Hydrochloric acid HClSour milk, yoghurt Lactic acid CH3CHOHCOOHAnt stings Methanoic acid HCOOH

Point 3.14 – Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations.

Calculating pH for acids:

Secondary data maybe from questions with [H+] concentrations given or the pH given. I will process the necessary data which will include values and calculations to be solved.

OR

Question: Calculate the pH of 0.1 mol L-1 hydrochloric acid.

pH = -log10 [H+] Where [H+] is the concentration of hydrogen ions in mol L-1.

Since HCl is fully ionised in water, the hydrogen ion concentration [H+] is 0.1 mol L-1. Using a calculator, substitute:

pH = -log10 [0.1] = 1 Therefore the pH is 1.

Question: Calculate the [H+] of carbonic acid which as a pH of 2.1.

For this problem, we use the equation:

Using calculator, substitute: [H+] = 10-2.1 x 2 (because H2)

= 0.016 mol L -1 is the hydrogen ion concentration

To calculate the pH of bases we use the following formula:

Question: Calculate the pH of a 0.001 mol L-1 sodium hydroxide (NaOH) solution

Since NaOH is completely ionised in water, the OH- (hydroxide ion) concentration is 0.001 mol L-1. Now we can calculate the [H+] by rearranging the formula:

[H+] = 1.0 x 10-14 / 0.001 = 1 x 10-11 mol L-1. Now we can calculate the pH:

pH = -log10 [1 x 10-11]

= 11 Therefore the pH is 11.

Question: Calculate the [H+] of sodium hydroxide which as a pH of 12.5

Substitute into formula:

[H+] = 10-12.5

= 3.16 x 10 -13 mol L -1

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.

Kw = [H+] [OH-] = 1.0 x 10-14

Where Kw is the water ionisation constant.

Point 4.1 – Outline the historical development of ideas about acids including those of: Lavoisier, Davy and Arrhenius.

Antoine Lavoisier (1776) observed that non-metallic oxides produced acidic solutions. He defined acids as substances containing oxygen e.g. H2SO4.

Humphry Davy (1810) disproved of Lavoisier’s theory by showing that HCl didn’t contain any oxygen. He proposed that acids were substances which contained hydrogen that could be partially or totally replaced by a metal.

H2SO4 (aq) + Mg MgSO4 (aq) + H2 (g)

Svante Arrhenius (1884) found that an acid released hydrogen ions H+ when dissolved in water and that they were electrolytes. Therefore he proposed that acids were substances which produce H+ ions in solution. He also realised that strong and weak acids differed according to their degree of ionisation.

H2SO4 (aq) 2H+ (aq) + SO4 2-

Lavoisier and Davy's definitions were based on observable properties while Arrhenius, put forward definitions based on concepts about particles too small to be directly observed.

Point 4.4 – Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature.

(I’m going to shorten Bronsted-Lowry to B-L). In general, a strong B-L acid (e.g. HCl) has a weak conjugate base while a strong B-L base has a weak conjugate acid. Acidic salts are formed when a weak base is neutralised by a strong acid. For example:

NH4NO3 (aq) + H2O (l) NH3 (g) + H3O+ (aq) + NO3- (aq)

In the above example, ammonium salt is acidic as it releases a hydrogen atom and water accepts it to form a hydronium ion. The nitrate ion is a very weak base which does not neutralise the hydronium ions and as a result the [H+] increases which makes it an acidic solution.

Sodium chloride solution is a neutral salt, because Na+ and Cl- (ions from the strong base NaOH and the strong acid HCl) do not undergo hydrolysis and are neutralised (i.e. property of the acid and base destroyed). The pH of the water is not changed.

Basic salts are formed when a weak acid is neutralised by a strong base. For example:

CO32- + H2O HCO3

- + OH-

The sodium carbonate salt solution is very basic as it accepts a hydrogen ion from water (acting as an acid) and it forms strong basic hydroxide ions which are not neutralised by the weak HCO3

- thus the [H+] decreases which makes it a basic solution.

Other salts:

Name of salt Acidic/neutral/basicAmmonium sulfate AcidicLithium chloride NeutralCalcium acetate Basic

Calcium carbonate BasicSodium sulfate Neutral

Point 4.5 – Identify conjugate acid/base pairs.

According to the Bronsted-Lowry theory, an acid has a conjugate base and a base has a conjugate acid.

Question: Hydrochloric acid dissolves in water to form a solution with a pH less than 7. Write a balanced equation for this reaction and use the Bronsted-Lowry theory to identify the conjugate acid-base pairs.

Equation: HCl (aq) + H2O (l) Cl- (aq) + H3O+

(aq) Acid and base in the forward reaction are called Bronsted-Lowry acid and base. HCl gives away a hydrogen ion (proton) and forms chloride ions therefore it is an acid. Water receives the hydrogen ion from HCl therefore it is a base.

Reverse reaction: Cl- (aq) + H3O+

(aq) HCl (aq) + H2O (l)

We can see in the reverse reaction, Cl- ion is the base and the H3O+ ion is the acid. Thus the conjugate acid-base pairs are:

HCl (B-L acid) / Cl- (conjugate base) H3O+ (conjugate acid)/ H2O (B-L base)

Other conjugate acids and bases:

Conjugate acid Conjugate baseH3O+ H2OHCl Cl-

CH3COOH CH3COO-

H2SO4 HSO4-

HSO4- SO4

2-

H2O OH-

Point 4.6 – Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions.

An amphiprotic substance is a molecule or ion which can act as an acid (proton donor) and as a base (proton acceptor) in an acid-base reaction. Examples of amphiprotic substances include: H2O, HCO3

- and HSO4-.

Water is an amphiprotic substance. The 1st equation shows water as an acid and the 2nd equation shows water as a base:

H2O + OH- OH- + H2O

In the above reaction water is acting as an acid as it donates a hydrogen ion (proton) to the hydroxide ion.

H2O + H3O+ H3O+ + H2O

In the above reaction water is acting as a base as it accepts a hydrogen ion from the hydronium ion.

The hydrogen carbonate (bicarbonate) ion is also an amphiprotic substance. The 1st equation shows hydrogen carbonate as a base and the 2nd equation shows water as an acid:

HCO3-+ H3O+ H2CO3 + H2O

In the above reaction the hydrogen carbonate ion is acting as a base as it accepts a hydrogen ion from the hydronium ion.

HCO3- + OH- CO3

2- + H2O

In the above reaction the hydrogen carbonate ion is acting as an acid as it donates a hydrogen ion to the hydroxide ion.

Point 4.10 – Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions.

Gather data about history of acid-base reactions such as the works of Arrhenius and Bronsted-Lowry. Secondary sources which I can use to obtain the data: websites such as HSC Online, reference texts, scientific journals and school material. Once I gather my information I will need to check the validity of the information by comparing the consistency of the information across a variety of different sources. Then I can process the data by picking out the main points and then collating them into a table. Compare Lavoisier, Davy, Arrhenius and Bronsted-Lowry theories.

For more information on this dot point go to Dot point 4.1

Point 4.11 - Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions.

Examples of salt solutions to test: ammonium chloride, potassium carbonate, sodium chloride, hydrogen sulfate etc.

Add 5 drops of each salt into a test tube and measure pH using universal indicator and pH probe + data logger.

Record results in a table as well as colour change Make sure you rinse pH probe with distilled water after measuring pH of each salt to

reduce chance of contamination. Repeat experiment 10 times for reliable results.

Accuracy: pH meter more accurate than indicator solution or universal indicator.

Reliability: Repeat experiment 10 times, compare results with other groups and average out results (if possible).

Validity: Have equal concentrations of salt solutions as well as same amount. Use a control (maybe).

Point 4.12 - 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.

Titration: Anhydrous sodium carbonate against hydrochloric acid. Purpose – to find the concentration of HCl.

Prepare 0.1 M, 250mL solution of anhydrous sodium carbonate.

Na2CO3 (aq) + 2HCl (aq) 2NaCl (aq) + H2O (l)

Number of moles of Na2CO3= (0.1) x (0.25) = 0.025moles

Mass of Na2CO3 – m = (n) x (M) = (0.025) x (105.99) = 2.65 g

Dissolve 2.65 g of Na2CO3 with distilled water up to 250mL in a volumetric flask. This is the standard solution. Concentration = 0.1M

Now rinse the pipette several times with distilled water and then with the some of the unknown concentration of HCl.

Now rinse the burette with distilled water and then with some of the standard solution. Fill the burette with the standard solution (Na2CO3) until bottom of meniscus is on

0.00mL. Using the pipette, take a known volume of unknown concentration solution (HCl) and put

it in a conical flask (e.g. 25mL of HCl). Conical flask also needs to be cleaned with distilled water prior this experiment.

Add a few drops of bromothymol blue indicator to the conical flask to measure the equivalence point (the point when exact quantities of acid and base have reacted). This will be shown when the end-point (the point at which the indicator changes colour) is achieved.

Now carefully add the standard solution from the burette into the conical flask. When the indicator changes colour, stop adding the solution from the burette and record the value on the burette. This is the first titre.

The 1st titration is a rough one so the next 3 titrations have to very accurate.

When 25.00mL of hydrochloric acid solution of unknown concentration was titrated (using above method) against a standard sodium carbonate solution of concentration 0.1M, 37.9mL of sodium carbonate was needed to completely neutralise HCl. Calculate the concentration of the hydrochloric acid.

Number of moles of Na2CO3 reacted = (37.9/1000) x 0.1 = 3.79 x 10-3 moles

Mole to mole ratio of Na2CO3: HCl = 1:2 = 3.79 x 10-3: 7.58 x 10-3

Therefore 7.58 x 10-3 moles of HCl react in the reaction.

Concentration = (7.58 x 10-3) / (25/1000) = 0.3032 = 0.303 mol L -1 HCl

For strong acid/strong base titrations you can use the following indicators: bromothymol blue, litmus.

For strong acid/weak base titrations use methyl orange indicator because it changes colour in the acidic region (pH 3.1 – 4.4)

For strong base/weak acid titration use phenolphthalein indicator as it changes colour in the basic region.

Point 4.13 - Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies.

Domestic household substance: vinegar

Here is a brief summary of the experiment:

We titrated vinegar (acetic acid) of unknown concentration against a sodium hydroxide standard solution.

Rinsed 50mL burette with distilled water then NaOH standard solution. Rinsed 2mL pipette with distilled water then vinegar. Rinsed conical flask with distilled water then added 3 drops of phenolphthalein

indicator. Used pipette to draw 2mL of vinegar (bottom of meniscus on calibration line) and

put it into conical flask. Filled burette with standard solution until bottom of meniscus is on 0.00mL. Slowly released standard solution from burette until solution in conical flask

changed colour (continuously swirl conical flask). This is rough titre Repeat procedure 3 more times to get concordant results (within 0.2mL of each

other). Repeat entire experiment 10 times to obtain reliable results.

Computer based technologies could include the use of a pH probe + data logger connected to a computer. The pH probe could be present inside the conical flask rather than the indicator drops, thus able to measure the continuous changes in pH. The computer could give us a titration graph which would clearly indicate the equivalence point of the reaction.

Point 4.14 – 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.

Secondary sources include: comprehension passages, flowcharts, internet sites, books, scientific journals etc.

Often in industries and laboratories, there are large acid and/or base spills. An effective method needs to be utilised to minimise danger. Acids and bases are corrosive and can cause irritations to eyes and skin thus any spill must be removed immediately.

The most effective method is through neutralisation reactions. If there is an acid spill, a weak base is used to neutralise it to form a salt and water:

NaHCO3 (aq) + HCl (aq) NaCl (aq) + CO2 (g) + H2O (l)

If there is a base spill, a weak acid is used to neutralise it:

CH3COOH (aq) + NaOH (aq) NaCH3COO (aq) + H2O (l)

The advantages of neutralisation reactions are:

They are a cost effective way of treating spills especially in industries and laboratories. Quick and easy method of treating potentially dangerous chemical spills. Amphiprotic substances such as NaHCO3 are used because it can neutralise an acid (1)

or a base (2).

(1) NaHCO3 (aq) + H3O+ (aq) Na+

(aq) + H2CO3 (aq) + H2O (l)

(2) NaHCO3 (aq) + OH- (aq) Na+ (aq) + CO32- + H2O (l)

If there is a strong acid or base spill, NaHCO3 can be safely used in excessive amounts as it is relatively harmless.

NaHCO3 is easy to handle because it is a solid and is easy to clean up. It is a stable solid which is safely handled and stored. Because it is a solid, it soaks up the spill.

Disadvantages of neutralisation reactions:

It is an exothermic reaction which means it releases heat. If a substance is added too quickly to neutralise spill, it may cause splattering which may lead to chemicals entering eyes and on skin causing burns and irritants. Safety goggles and lab coats need to be worn.

Not good to use a strong acid/base to neutralise strong base/acid because it is a violent reaction releasing lots of heat.

Neutralisation reactions should not be carried out on the skin as it will only aggravate the problem by causing severe burns. Instead it is better to wash skin with plenty of cold water.

Although it should not be used on the skin, a neutralisation reaction is a very effective way of combating accidents or chemical spills involving acids and/or bases as it is extremely efficient and cost effective due to the use of amphiprotic substances such as NaHCO3. With more research in the future, the danger of the use of neutralisation reactions will only be reduced.

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

6. Point 5.2 – 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 (alkyl alkanoates) are organic compounds prepared by the condensation of an alkanol and an alkanoic acid. Example:

Alkanol + Organic acid Ester + Water

Here is a table showing esters which are produced from alkanols and alkanoic acids:

Alkanol Alkanoic acid EsterMethanol Ethanoic acid Methyl ethanoateEthanol Propanoic acid Ethyl propanoatePropanol Butanoic acid Propyl butanoateButanol Pentanoic acid Butyl pentanoatePentanol Hexanoic acid Pentyl hexanoateHexanol Heptanoic acid Hexyl heptanoateHeptanol Octanoic acid Heptyl octanoateOctanol Methanoic acid Octyl methanoate

Point 5.8 – Identify data, plan, select equipment and perform a first-hand investigation to prepare an ester using reflux.

Data needed to be identified: the boiling points of the alkanol, alkanoic acid, ester and water. Plan the investigation by considering if the boiling points of the reactants and products are sufficiently different for you to easily separate the ester from the rest of the reaction mixture by fractional distillation. Consider the effect of using more than a few drops of concentrated sulfuric acid catalyst and how this might complicate the separation. Consider if you would be better off using differences in densities and water solubility of the components to separate the ester.

Equipment:

20mL alkanol 10mL alkanoic acid 2mL concentrated sulfuric acid Boiling chips Retort stand + clamp + boss head Heating mantle Condenser Round bottomed flask

Allow mixture to cool. Transfer mixture in round bottomed flask to a separating funnel containing water of about

the same volume of the mixture. Shake contents and allow layers to separate. Lower layer contains water soluble acid and alkanol. This can be carefully discarded. Add solid sodium carbonate to remaining mixture to neutralise all acid present. Add water to dissolve any salts. Shake and let layers separate and discard lower layer to leave ester.

Point 5.9 – Process information from secondary sources to identify and describe the uses of esters as flavours and perfumes in processed foods and cosmetics.

This point is in conjunction with dot point 5.7.

Pentyl butanoate is an ester which has an apricot/pear-like odour. Because of its strong, sweet smelling fragrance, it is widely used in the cosmetics industry as a major ingredient in many perfumes.

Ethyl butanoate is an ester which has a pineapple-like smell. Because of this smell, ethyl butanoate is used in the manufacture of artificial rum as its adds an unique pineapple aroma to the rum.

Ethyl ethanoate is an ester which is used in the cosmetics industry as a nail polish remover. It is also used in perfumes where it evaporates quickly leaving a scent of the perfume on the skin. This ester is also used in the food industry where it is found in confectionary and some fruits.