Chemistry:

The Acidic Environment1 Indicators were identified with the observation that the colour of some flowers depends on soil composition

1.1 Classify common substances as acidic, basic or neutral:Acids: General Properties:

o They taste souro pH < 7o H+ donoro Acids are neutralised by baseso When in solution, they can conduct electricity (electrolyte)o They are corrosiveo They affect the colour of certain natural and synthetic dyes (indicators)o For LITMUS: blue acid red (litmus is a dye made from LICHENS)

Strong acids: HCl H2SO4 HNO3 HBr Weak acids: Acetic (CH3COOH), citric (C6H8O7), formic (HCOOH), HF

Bases: General Properties:

o They taste bittero pH > 7o H+ acceptoro Bases are neutralised by acidso Are electrolytes (conduct electricity in solution)o They affect the colour of indicatorso For LITMUS: red base blue o They feel slippery (bases react with oils on our skin, forming soaps)o Are mainly insoluble in water (aqueous bases are called alkalis)

Strong bases: NaOH, X+OH- (where X is a metal) Weak bases: NH3 (ammonia), CO3

2-

Page | 1

Table of examples of acidic, basic and neutral substances:Chemical/substance Acidic/neutral/basic

Acetic (ethanoic acid) AcidicHydrochloric acid AcidicLemon and citric juices AcidicStomach juices AcidicVinegar AcidicMagnesium sulphate Slightly basicSodium hydrogen carbonate Basic/acidic [amphiprotic]Sodium hydroxide BasicDrain cleaner BasicWashing soda BasicPure water NeutralEthanol NeutralSugar solution NeutralSodium chloride solution (salt) Neutral

1.2 Identify that indicators such as litmus, phenolphthalein, methyl orange and bromothol blue can be used to determine the acidic or basic nature of material over a range, and that range is identified by a change in indicator colour:Indicator: An indicator is a substance which in solution changes colour depending on whether the solution is acidic or alkalineCommonly used indicators: Litmus:

o red for pH 0 – 5.0o from pH 5.0 – 8.0 it goes through purple. o turns blue for solutions higher than pH 8.0

Phenolphthalein: o colourless between pH 0 – 8.3 o from pH 8.3 – 10.0 it goes through pink. o turns crimson for solutions higher than pH 10.0

Methyl Orange: o red for pH 0 – 3.1 o from pH 3.1 – 4.4 it goes through orangeo turns yellow for pH > 4.4

Bromothymol Blue:o yellow for pH 0 – 6.0 o from pH 6.0 – 7.6 it goes through greeno turns blue for solutions higher than pH 7.6

Universal Indicator: is a convenient mixed indicator which is also commonly used

Page | 2

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

o Some plants prefer soils that are acidic, some crops prefer basic soilso When testing the pH of soil, the colour can hide the indicator colour change. To prevent

this, a neutral white powder, such as barium sulphate, can be added to the top layer of damp soil before adding the indicator. Excess soil acidity can be fixed by adding basic substances such as lime.

Checking the pH of water in swimming pools: o pH in pool needs to be kept around 7.4. If the water is too acidic, it will cause irritations to

the eyes and skin. If the water is too basic, green algal scum will grow in the pool. Monitoring pH of Chemical wastes

o Wastes produced from laboratories or photographic film centres tend to be highly acidic.o pH of wastes must be neutralised before they can be safely disposed.o Indicators are used to measure the pH, and substances added to neutralise it

1.5 Identify data and choose resources to gather information about the colour changes of a range of indicators:

Indicator Colour in base Colour in acidLitmus Red (below pH = 5) Blue (above pH 7.6)Phenolphthalein Colourless (below pH = 8.3) Red (above pH = 10)Bromothymol Blue Yellow (below pH = 6.0) Blue (above pH = 7.6)Methyl Orange Red (below pH = 3.1) Yellow (above pH = 4.4)

Page | 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 concentrations of these acidic oxides have been increasing since the Industrial Revolution

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

Acidic Oxides:o Reacts with water to form an acid / reacts with bases to form saltso Are the oxides of non-metalso When they are in solution they act as acids: e.g.

Carbon dioxide (CO2): CO2 (g) + H2O (l) H2CO3 (aq)

Carbon dioxide + water carbonic acid Sulfur dioxide (SO2):

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

Sulfur dioxide + water sulfurous acid Nitrogen dioxide (NO2):

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

Nitrogen dioxide + water Nitrous acid (weak) + nitric acid (strong) These are all gaseous at room temperature and are all acidic

2.2 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:

Metals are located on the LHS of the periodic table. Generally, metal oxides are basic. Metals closer to the bottom left-hand corner are more basic, because the metals become more metallic in their properties.

Non-metals are located on the RHS of the periodic table. Generally, non-metal oxides are acidic (except for Group 8 – the noble gases – which do not form oxides). The oxides of non-metals become more acidic toward the top right-hand corner of the table.

The oxides of some elements between these two groups are amphoteric (can act as either an acid or a base). These tend to be the oxides of the semi-metals.

As you move across the periods, the oxides move from being basic to acidic (in different degrees)

Page | 4

2.3 Define Le Chatelier’s Principle: When a system at equilibrium is changed, the equilibrium re-establishes to minimize the

effect of that change This means that either the forward or reverse reaction proceeds at a faster rate until

equilibrium is re-established A system is in equilibrium when the forward and reverse reactions are occurring at the same

rate For equilibrium to be established the system must be closed; i.e. no loss of reactants or

products occurring

2.4 Identify factors which can affect the equilibrium in a reversible reaction: Pressure (only when gases are involved), temperature and concentration can affect the

equilibrium in a reversible reaction Note: Catalysts increase the rate of reaction, as equilibrium is reached faster, but do not

affect the point of equilibrium

Change in Concentration:o Take for example, the copper complex ion equilibrium:

EG: Cu(H2O)42+

(aq) + 4Cl־(aq) CuCl4־2

(aq) + 4H2O (l)

(BLUE) (GREEN)o If Cl־ ions are increased the equilibrium will shift to the RIGHT, as more CuCl4

is־2formed hence the system will become more green.

o If more water is added the equilibrium will shift to the LEFT, as more Cu(H2O)42+ is

formed, and the system will become bluer.o If water is removed the equilibrium will shift to the RIGHT, so that the water lost will

be replaced by the forward reaction.

Change in Gas Pressure (volume):o Take, for example, the equilibrium between dinitrogen-tetroxide and nitrogen dioxide:

EG: N2O4 (g) 2NO2 (g)

(1 mole) (2 moles) o If the TOTAL pressure is increased, the equilibrium will shift to the left, to decrease

pressure, as there are fewer moles produce on the left.o If the TOTAL pressure is decreased, the equilibrium will shift to the right, to increase

the pressure, as more moles are produced on the right.

Change in Temperature: o Take for example, the decomposition of calcium carbonate (within a closed system; that is,

nothing is allowed to escape). It is an endothermic reaction; the change in heat is positive:EG: CaCO3 (s) CaO (s) + CO2 (g) ΔH = 178 kJ/mol

o If the temperature is INCREASED, the forward reaction will increase, with equilibrium lying more on the right, as the endothermic forward cooling opposes the imposed heating.

o If the temperature is DECREASED, the reverse reaction will increase, with equilibrium lying more on the left, as the exothermic reverse heating opposes the imposed cooling.

Page | 5

2.5 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 carbon dioxide, water and carbonic acid system:

The solution of carbon dioxide gas in water: CO2(g) CO2(aq) [Equation 1]The reaction of dissolved carbon dioxide with water: H2O(l) + CO2(aq) H2CO3(aq) [Equation 2]Carbonic acid behaves as an acid by releasing a hydrogen ion: H2CO3(aq) H+

(aq) + HCO3-

(aq) [Equation 3]

Consider the following changes: An increase in CO2 gas pressure:

o Increases concentration of dissolved CO2

o Therefore, Eq2 shifts the equilibrium to the right, increasing concentration of carbonic acid ( H2CO3(aq) )

o Eq3 shifts equilibrium to the right, increasing concentration of H+ and lowers the pH (more acidic)

Raising the temperature of the system:o Increases the amount of heat energy available for decomposition of a substanceo Thus, carbonic acid ( H2CO3(aq) ) decomposes and the Eq2 equilibrium shifts to the lefto The higher concentration of CO2(aq) causes Eq1 to shift to the left and more CO2(g) is

released Adding alkali to the solution:

o Hydroxide ions from the alkali will react with the H+ in Eq3o Thus, lowering concentration of H+, causing Eq3 equilibrium to shift to the righto This lowers concentration of carbonic acid ( H2CO3(aq) ) and causes Eq2 equilibrium to

shift to the righto This lowers the concentration of dissolved CO2 resulting in the equilibrium in Eq1 to

shift to the right, thus increasing the solubility of CO2

2.6 Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen:Sulfur Dioxide (SO2):

NATURAL: Volcanic gases, bushfires, decomposition of organic matter and sulfur-rich geothermal hot springs releasing gases.

INDUSTRIAL: Processing and burning of fossil fuels and extracting metals from sulfur-rich ores, such as galena (PbS).

Nitrogen Oxides (NO, NO2): NATURAL: The reaction of nitrogen and oxygen in the atmosphere due to high temperatures

of lightning. INDUSTRIAL: Combustion of fossil fuels, both in cars and in power stations. The nitrogen in

the air reacts with oxygen in the hot engines. Power stations release large volumes of NO2 into the atmosphere.

=

Page | 6

2.7 Describe using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen:Sulfur Dioxide:

When organic matter decomposes it produces hydrogen sulfide (H2S), which then oxidises (reacts with oxygen) to produce sulfur dioxide:

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

The burning of sulfur-rich coal and other fossil fuels directly combines sulfur with oxygen:S (s) + O2 (g) SO2 (g)

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

2PbS (s) + 3O2 (g) 2PbO (s) + 2SO2

Oxides of Nitrogen: Nitric oxide is produced either when lightning, with its high temperatures combines nitrogen

and oxygen:N2 (g) + O2 (g) 2NO (g)

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

Nitrogen dioxide is formed when nitric oxide reacts with oxygen in the air:2NO (g) + O2 (g) 2NO2 (g)

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

Oxides of sulfur and nitrogen have been increasing in the atmosphere; these oxides occur in relatively low concentrations, such as 0.01ppm (parts per million)

Analysis of gas found in ice-core samples excavated from Antarctica shows that levels of N2O in the atmosphere has increased by about 10%

It can also be stated that the increased burning of fossil fuels after the Industrial Revolution did indeed lead to a rise in oxides of sulfur; evidence for this is that the air quality of major industrial cities, such as London, deteriorated greatly

NO2 leads to the formation of photochemical smog, a direct indicator of excessive levels of nitrogen oxides in the atmosphere

Acid rain forms when atmospheric water reacts with these compounds; hence an increase in acid rain points to an increase in these compounds

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 0oC and 100kPa or 25oC and 100kPa:Definitions:

Mole: the quantity of a substance that contains 6.022 x 1023 particles (Avogadro’s number) Molar mass: the formula mass of a substance in grams (molecular weight) Molar volume: the volume, in litres, occupied by one mole of any gas at a particular

temperature and pressure.The molar volume of any gas is always the same for any particular temperature and pressure.

o At 100kPa pressure, one mole of any gas will occupy 22.71L, at 0oC, and 24.79L at 25oC

Page | 7

Calculations of molar volumes require balanced equations and use of mole relationships.o Molar formulas:

2.10 Explain the formation and effects of acid rain: Acid rain is any rain that has a pH < 5

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

o SO2 (g) + H2O (l) H2SO3 (aq) Sulfurous acid then reacts with oxygen; this is catalysed by air particles:

o 2H2SO3 (aq) + O2 (g) 2H2SO4 (aq) Nitrogen dioxide also reacts with rain, making nitric and nitrous acids:

o 2NO2 (g) + H2O (l) HNO3 (aq) + HNO2 (aq) Nitrous acid then reacts with oxygen, again catalysed by air particles:

o 2HNO2 (aq) + O2 (g) 2HNO3 (aq) Thus, in industrialised areas, rain can contain relatively high levels of strong acids - nitric and

sulfuric acids

Impacts of acid rain on society and the environment: Urban and structural damage to metals and alloys such as structural steel, erodes marble

and limestone statues and buildings (contains carbonates which acids readily react with) Can cause stunted plant growth, defoliation and removes essential nutrients Acid rain lowers pH of lakes and streams, killing sensitive aquatic life Insoluble compounds in the soil, i.e. aluminium sulphate, become soluble in acidic water,

releasing toxic aluminium ions into the soil and streams sulphate particles have contributed to the rising incidence of respiratory diseases in humans

2.11 Analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and the oxides of nitrogen and evaluate reasons for concern about their release in the environment:

After the industrial revolution in the early 1800’s there was a great increase in emissions of sulfur dioxide in growing industrial cities, leading to the degradation of air quality

The industrial sources of sulfur dioxide include the burning of coal in power stations and the smelting of sulphide ores

The sources nitrogen oxides include car engines and other high temperature combustion environments

Eventually, late in the 20th century regulations were passed which controlled these emissions

There is environmental concerns about the concentrations of these gases because nitrogen dioxide and sulfur dioxide irritates the respiratory tract and causes breathing discomfort

Page | 8

Other concerns include those as a result of acid rain The main problem with nitrogen dioxide is that it leads to the formation of ozone in

photochemical smog which has harmful effects

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

3.1 Define acids as proton donors and describe the ionisation of acids in water: An acid is a substance that releases H+ ions Another name for the H+ ion is a ‘proton’; this is because a positive hydrogen atom (H+) is

just a hydrogen nucleus (no electrons) and hence is just a proton When acids react with other substances, the H+ ion is transferred to another species; that is

why acids can be defined as proton-donorsIn water, acids ionise (separate into its ions):

E.g: Pure hydrogen chloride added to water:HCl (g) H+

(aq) + Cl־(aq)

The complete reaction - ionisation:HCl (g) + H2O (l) H3O+

(aq) + Cl־(aq)

So in reality, acids ionise in water to form hydronium ions (H3O+) Acids that can only donate one H+ are called monoprotic, i.e. HCl Acids that can donate two H+ are called diprotic. i.e. sulfuric acid (H2SO4) Those that can donate three H+ are called triprotic. i.e. phosphoric acid (H3PO4)

Page | 9

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

Name Formula Found in/used forAcetic acid, systematic name: ethanoic acid

CH3COOH Vinegar

Citric acid (2-hydroxypropane -1,2,3-tricarboxylic acid)

C6H8O7 Citrus fruitsWidely used as a preservative

Hydrochloric acid HCl Stomach acidUsed to clean bricks

Sulfuric acid H2SO4 Most industrially produced acid, acid rain, car batteries, fertilizers

3.3 Describe the use of the pH scale in comparing acids and bases: The pH scale is used to determine the acidity or basicity of a substance.

o It is numbered from 0 to 14. A pH of 7 is attributed to neutral substances (mainly, PURE water). A pH < 7 refers to acidic substances A pH > 7 refers to basic substances

Hence, the pH scale allows us to compare acids and bases, and their strengths.

3.4 Describe acids and their solutions with the appropriate use of the terms strong, concentrated and dilute:

A strong acid is one which completely ionises in water (100%) o Strong acids include hydrochloric acid, nitric acid, hydrobromic acid and sulfuric acid

A weak acid is an acid which only partially ionises in water o Weak acids include acetic (ethanoic) acid, citric acid and carbonic acid (H2CO3)

A concentrated acid is one which has a high number of moles per litre A dilute acid is one which has a low number of moles per litre

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

The way pH is calculated is through an equation related to H+ concentration This equation is:

pH = -log10 [H+] A change of one in the pH scale means a ten fold change in the concentration of hydrogen

ions because pH is based on a logarithmic scale (base 10)o EG: pH of 4; [H+] = 10-4 while a pH of 3; [H+] = 10-3

Therefore; 10-3 / 10-4 = 10 Note: In calculations; don’t forget to consider whether or not the solution is

diprotic/triprotic and adjust calculations accordingly

Page | 10

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:

The degree of ionisation of an acid refers to the percentage of H+ ions that have been released

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

Recall that strong acids completely ionise in solution, while weak acids only partially ionise in solution

Hydrochloric acid is a strong acid, meaning that it fully ionizes in water (100% degree of ionisation)

Citric and acetic acid are weak acids, meaning that they only partially ionize in water

3.7 Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions:Strong acids:

Completely ionise [H+] ions released joins with water molecules to form hydronium [H3O+] ions No equilibrium in the ionisation reaction for strong acids, only the forward reaction occurs

and the reaction goes to completionWeak acids:

Only partially ionise Equilibrium reaction; note that the equation lies mostly on the left as it is partial ionisation The molecular form of the acid is in equilibrium with its ions The weaker the acid, the smaller % of acid molecules ionized

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

Acids are added to food for 2 reasons: as preservatives, and to add flavourPreservatives:

Ethanoic acid (in the form of vinegar) is used as a preservative, e.g. in pickles Propanoic acid is often used as a preservative in bread Sulfur dioxide is added to food as a preservative, as it forms sulfurous acid, which kills

bacteria in food – i.e. in Devon and bologna Citric acid is a natural preservative, often added to jams and conserves

Flavour: Carbonic acid is added to soft drinks to add ‘fizz’ Phosphoric acid is also added to soft drinks to add ‘tartness’ of flavour Ethanoic acid, as vinegar, is also used as flavouring

Page | 11

3.9 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:

Common Substance Acid present Chemical compositionCitric juice, i.e. oranges and lemons Citric acid Organic compoundsVinegar Acetic acid Organic compound and waterStomach acid Hydrochloric acid HClSour milk, yoghurt Lactic acid CH3CHOHCOOH

Ant stings Methanoic acid HCOOH

Naturally occurring bases:Common substance Base present Natural/ synthetic Chemical composition

Limewater Calcium hydroxide Synthetic Ca(OH)2

Bicarbonate of soda Sodium hydrogen carbonate

Natural NaHCO3

Lime Calcium oxide Synthetic CaODrain and oven cleaner Sodium hydroxide Synthetic NaOHHousehold cleaners (with ammonia)

Ammonium hydroxide Synthetic NH4OH

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.

4.1 Outline the historical development of ideas about acids including those of Lavoisier, Davy and Arrhenius:Lavoisier; stated that ACIDS contain OXYGEN: Proposed that acids were corrosive substances that all contained oxygen. He hypothesised that since many of the common acids contain oxygen (such as acetic acid,

CH3COOH, carbonic acid, H2CO3, sulfuric acid, H2SO4), all acids must contain oxygen. He thought that oxygen was the source of the acidity

Davy; stated that ACIDS contain HYDROGEN: He showed that hydrochloric acid (HCl) did not contain oxygen, disproving Lavoisier’s hypothesis Many other non-oxygen containing acids had been discovered, such as hydrofluoric acid (HF),

hydrobromic acid (HBr) and hydrocyanic acid (HCN). Thus he hypothesised that all acids contain hydrogen.Arrhenius; stated that ACIDS produced H + IONS when in WATER: He proposed the idea that acids disassociate into their ions when they are dissolved in water. Thus, he hypothesised that acids release a H+ when in an aqueous solution. He also said that

bases release OH־ ions in aqueous solutions.

Page | 12

4.2 Outline the Bronsted-Lowry theory of acids and bases: In the Bronsted Lowry theory, acids are defined as proton donors and bases are defined as

proton acceptors An acid donate a proton (H+) to a base A base will accept a proton (H+) from an acid

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

An acids conjugate base is the base it produces when it gives up its hydrogen iono A conjugate base is the original acid with a proton removed

A bases conjugate acid is the acid it produces when it gains a hydrogen iono A conjugate acid is the original base with a proton added

EG: Reaction between hydrochloric acid and water:HCl (aq) + H2O (l) H3O+

(aq) + Cl־ (aq)

o The original acid is HCl; its conjugate base is Cl־ (acid minus proton).o The original base is H2O; its conjugate acid is H3O+ (base plus proton).

The relationship between acid/bases and their conjugates:Acid/Base Conjugate

STRONG acid (e.g. HCl) Extremely WEAK base (e.g. Cl-)STRONG base (e.g. OH-) Extremely WEAK acid (e.g. H2O)WEAK acid (e.g. CH3COOH) WEAK base (e.g. CH3COO-)WEAK base (e.g. NO2

-) WEAK acid (e.g. HNO2)Extremely WEAK acid (e.g. H2O) STRONG base (e.g. OH-)Extremely WEAK base (e.g. H2O) STRONG acid (e.g. H3O+)

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

A strong acid and a strong base will make a neutral salt A weak acid and a weak base will make a neutral salt A strong acid and a weak base will make a slightly acidic salt A weak acid and a strong base will make a slightly basic salt

These are the general rules for neutralisation reactions, but the acidity or alkalinity of a salt must be proven by reacting the salt’s ions with water:

Weak acid and strong base:E.g. CH3COOH + NaOH CH3COONa + H2O

o Reacting ions with water – proving basicity: CH3COO- + H2O CH3COOH + OH - Strong acid and weak base:

E.g. HNO3 + NH3 NH4NO3

o Reacting ions with water – proving acidity: NH4+ + H2O NH3 + H3O +

Strong acid and strong base:E.g. KOH + HCl KCl + H2O

o Neither of the ions in the product reacts with water; hence it is neutral, pH=7 Weak acid and weak base: similar to above.

Page | 13

4.5 Identify conjugate acid base pairs: Conjugate base is the acid without a proton; conjugate acid is the base with a proton.

o e.g. HCl (aq) + H2O (l) H3O+ (aq) + Cl־ (aq)

Acid/base pairs are HCl/Cl־ and H2O/H3O+

o e.g. NH3 (aq) + H2O (l) NH4+

(aq) + OH־ (aq)

Acid/base pairs are H2O/OH־ and NH3/NH4+

o e.g. What is the conjugate acid of NH2-?

It is the base plus a proton; hence it is NH3

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

An amphiprotic substance is one that can act as both an acid and a base by giving and receiving protons

Behaviour depends on the environment they are placed in:o E.g. The hydrogen carbonate ion HCO 3

- is amphiprotic

o HCO3 ־

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

o HCO3 ־

(aq) + OH ־ (aq) CO3

־2 (aq) + H2O (l)

Other amphiprotic species include the hydrogen sulfite ion (HSO3and water(H2O) (־

Note : Amphoteric refers to any substance that can act as an acid or a base, including amphiprotic substances which do so by accepting or donating a proton. All amphiprotic substances are amphoteric. However, some substances are amphoteric without being amphiprotic – they can act like an acid and a base, but can’t accept or donate protons. Examples include: Zn, Al

4.7 Identify neutralisation as a proton transfer reaction which is exothermic: Neutralisation reactions are reactions between acids and bases As acids are proton-donors & bases are proton-acceptors, neutralisation reactions between

acids and bases are clearly proton-transfer reactions All neutralisation reactions are exothermic; they all release heat energy

4.8 Describe the correct technique for conducting titrations and preparation of standard solutions:Standard Solution:

A solution of accurately known concentration is called a standard solution For a chemical to be suitable to prepare as a standard solution, it must:

o Be a water soluble solid and have high purityo Have an accurately known formulao Be stable in air; i.e. it does not lose/gain water or react with oxygen or CO2 in air

The solution is prepared by:o Accurately weighing a calculated amount of solid o Dissolving it in watero Transferring the entire dissolved solid to a volumetric flask o Adding water to the flask to prepare a fixed volume of solution

Page | 14

A standard solution can be reacted with a solution of unknown concentration using the titration technique

One reactant in solution is slowly added to another reactant in solution until an end point is reached. The end point of the titration is usually indicated by a change in colour of a small amount of indicator solution added to the mixture of reactants.

Titrations: A titration is a procedure used to find experimentally the concentration in moles per litre

(molarity) of a solution Preparing for titrations involves appropriately rinsing the glassware to be used:

o Volumetric flask: - Rinsed thoroughly with distilled water multiple times

o Conical flask:- Rinsed thoroughly with distilled water multiple times

o Burettes:- Rinsed thoroughly with distilled water multiple times – with tap opened to allow water to flow out and thoroughly rinse the tip – with tap closed to wash the sides of the burette and poured out the top.- It is then finally rinsed by the same procedure with the solution going into it

o Pipette:- Rinsed thoroughly with distilled water multiple times- Then rinsed with the solution going into it

Once appropriate rinsing has been performed, the glasswares can be filledo For this example, presume the acid is added to the burette and the base to the conical

flasko The acid is poured into the burette and measured so that the bottom of the meniscus is

on the zero marko Using the pipette and the pipette filler bulb, a fixed volume of the base (say 25 mL) is

drawn from the volumetric flask and deposited into the conical flask o A few drops of the chosen indicator are added to the conical flask

Performing the titration: o The conical flask is placed on a white tile (to make the solution’s colour clear) under the

burette, which is held by a clamp on a retort stando The tap is slowly opened, and the conical flask is continuously swirled in the process of

drops addedo When colour changes become significant and longer lasting, begin adding drops slowly

and stop when the indicator changes colour definitely – at the end pointo Titrations are performed multiple times to achieve accurate results

Page | 15

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

A buffer is a solution that resists rapid changes to pH when an acid or base is added Buffer solutions must contain components which will remove any hydrogen ions or

hydroxide ions that you might add to it – or else the pH will change Buffer solutions contain equal concentrations of a weak acid and its conjugate base An example of a buffer in a natural system is hydrogen carbonate and carbonate ions

present in human blood which help maintain its pH at about 7.4

4.10 Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills:

It is important to immediately neutralise any chemical spills involving strong acids and bases, as they are corrosive and can be extremely dangerous.

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

When neutralising an acid or a base the following procedure is followed:o The most preferred agents of neutralisation has the properties of being stable, easily

transported, solid (powdered), cheap and amphiprotic (so it can act as a WEAK acid or a WEAK base).

o This is the safest material, as it can neutralise both acids and bases; even if an excess is used, it is very weak, and so does not pose any safety risks.

o The neutralised product is then absorbed using paper towels and disposed. The most common substance used to neutralise spills in laboratories is powdered sodium

hydrogen carbonate; this is because the hydrogen carbonate ion (HCO3-) is an amphiprotic

species, and it is cheap and readily available substance. Strong acids and bases must never be used to neutralise spills; if an excess is used, the spill

will become dangerous again.

Page | 16

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

5.1 Describe the differences between the alkanol and alkanoic acid functional groups in carbon compounds:Alkanol:

Functional group: hydroxyl group (–OH): Structure: an oxygen and a hydrogen molecule covalently bonded. General formula: CnH2n+1OH Polarity: polar due to OH functional group Soluble in water Alkanols are alcohols derived from alkanes (functional group with single C–C bonds) E.g. Ethanol (C2H5OH), Butanol (C4H9OH)

Alkanoic Acid: Functional group: carboxyl group (–COOH). Structure: an oxygen is double-bonded to a central carbon, and an OH group is

single-bonded to the same carbon General formula: Cn-1H2n+1COOH Polarity: more polar than alkanols due to the presence of the COOH functional group Soluble in water Alkanoic acids are carboxylic acids derived from alkanes Higher melting and boiling points than similar alkanols E.g. Methanoic acid (HCOOH), Propanoic acid (C2H5COOH)

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 C1to C8:Forming Esters: Alkanol + Alkanoic acid EsterNaming Alkanoic Acids:

Count the number of carbons; taking the name of the parent alkane with the same number of carbons, drop the ‘e’ and add on ‘-oic acid’

There are 2 exceptions to the rule; the IUPAC-preferred name for the alkanoic acids for 1-C and 2-C are not methanoic or ethanoic acids, but rather formic acid and acetic acid

Naming Alkanols: Similarly, count the number of carbons, take the parent alkane name, drop the ‘e’ and add

on an ‘-ol’ In this case, there are no exceptions. The IUPAC names coincide with all the systematic

names; methanol and ethanol are considered correctNaming Esters:

An ester is always in the following order: alkanol then alkanoic acid If given the alkanol and the alkanoic acid, name the ester:

o Firstly, take the alkanol and replace ‘-anol’ with ‘-yl’o Secondly, take the alkanoic acid and replace ‘-oic acid’ with ‘-oate’o Finally, place the 2 words together; alkanol then alkanoic acid

Page | 17

If given the ester structural formula, name the ester:o Firstly, you have to identify how many carbons were in the alkanol and the alkanoic

acid; split the ester along the C-O-C bondo The side with the C=O bond is the acid, as only the carboxyl group has a double bond

o E.g. propyl-butanoate

o Continue with naming by same procedure above

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

For the same number of carbons in a straight carbon-chain, the highest boiling points and melting points belong to the alkanoic acids, then the alkanols, and then the parent alkanes

The intermolecular forces between a collection of molecules determines what physical state they will exist in (solid, liquid or gas) for a given pressure and temperature.

The stronger the intermolecular forces, the more ‘tightly bound’ the molecules are to each other, and hence more energy needs to be forced into the system to overcome these forces (i.e. higher melting or boiling point)

Alkanols: Dispersion forces (weak) Two polar bonds; C–O and O–H bonds Hydrogen bonding occurs between molecules

Alkanoic Acid: Three polar bonds in each molecule; C–O, C=O and O–H bonds Hydrogen bonding occurs between molecules Therefore alkanoic acids have higher melting points and boiling points than alkanols

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

Esterification is the process which forms esters. In the most general sense it is the reaction between an acid and an alkanol.

Esterification is a condensation reaction; i.e. a water molecule is generated Esterification is a reversible reaction in which equilibrium lies much to the left at room

temperature. It is a moderately slow endothermic (absorbs heat) reaction. Generally:

o

E.g. Methanol + acetic (ethanoic) acid methyl acetate + waterCH3OH + CH3COOH CH3COOCH3 + H2O

Page | 18

R–OH + R! –COOH R-COO-R! + H2O alkanol acid ester water

5.5 Describe the purpose of using acid in esterification for catalysis: Concentrated sulfuric acid (H2SO4) is often added during the process of esterification; this

serves two purposes:o Sulfuric acid acts as a catalyst. It speeds up the rate of reaction, allowing the point of

equilibrium to be reached faster.o Sulfuric acid increases the yield/cause of the reaction. It does this by acting as a

dehydrating agent; it absorbs the water, encouraging the forward reaction, shifting equilibrium to the right according to Le Chatelier’s principle.

5.6 Explain the need for refluxing during esterification: Heating the reaction flask has 2 main benefits:

o The higher the temperature, the faster the rate of reaction; equilibrium can be reached much faster than if it was left at room temperature.

o Also, esterification is an endothermic reaction; increasing the heat of the flask encourages the forward reaction, creating more ester.

However, the ester, alkanol and alkanoic acid are all highly volatile substances; any open heating will cause the reactants as well as products to evaporate away.

A refluxing apparatus is basically a condenser placed vertically onto a boiling flask; it cools any vapours that boil off so that they drip back into the flask.

Refluxing allows the mixture to react at high temperatures without fear that the reactants or products will evaporate away.

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

Esters occur commonly in nature in the form of flavouring and scents The scents and flavours of fruits are caused by the presence of esters Animal fats and plant oils are also esters

Production & Uses:Many esters are industrially produced for many reasons/uses: Domestic uses of esters include artificial flavourings for foods, scents for perfumes and as

nail polish remover (ethyl acetate) Short esters such as ethyl acetate are used as industrial solvents, where as larger esters are

used as plasticisers to soften hard plastics (like PVC).

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

Esters are industrially produced to mimic flavours and scents found in nature Domestic food flavourings are often esters dissolved in a solvent such as ethanol Many processed foods are flavoured artificially; e.g. banana-flavoured milk is flavoured with

the ester iso-pentyl acetate. Cosmetics contain esters as scents, such as perfumes, which are comprised of a mixture of

esters in a solvent, or to give soaps, hand-lotions or other cosmetics a pleasant smell. Other cosmetic uses of esters include as solvents for other products, such as nail polish

removers (ethyl acetate)

Page | 19

Practicals to cover:

(1.) Perform a first-hand investigation to prepare and test a natural indicator.

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

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

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

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

(3.) Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids.

(4.) Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions.

(4.) 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.

(4.) Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies.

(5.) Identify data, plan, select equipment and perform a firsthand investigation to prepare an ester using reflux.

Page | 20