THE ACIDIC ENVIRONMENT

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

- Classify common substances as acidic, basic or neutral

Common acids: Vinegar, citric acid, carbonated soft drinksCommon bases: Ammonia, caustic soda, washing soda, shampoo, soapNeutral: Water

- 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

An indicator is a substance that changes colour in solution according to its acidity or alkalinity.

Methyl orange: 3.1 – 4.4: Orange <3.1: Red >4.4 : Yellow

Litmus: 4.5 – 8.2: Purple<4.5: Red>8.2: Blue

Bromothymol Blue: 6.0 – 7.5: Green<6.0: Yellow>7.5: Blue

Phenolphthalein: 8.3 – 10: Pale pink<8.3: Colourless>10: Pink

- Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity

Soil pH testing - Different plants grow better under different pH conditionsTesting home swimming pools – Need to be approximately neutralMonitoring wastes – discharge must be approximately neutralChemical analysis

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

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

Most non-metal oxides are acidic, while most metal oxides are basic. A basic oxide neutralises an acidic solution, and an acidic oxide neutralises a basic solution. Some oxides are amphoteric, with the ability to act as both acids and bases. Other oxides are neutral.

- 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

Generally, as electronegativity decreases to the left of a period and down a group, the oxides of the elements become more basic. Most metal oxides are basic, on the left of the periodic table.

Generally, as electronegativity increases to the right of a period and up a group, the oxides of the elements become more acidic. Most non-metal oxides are acidic, on the right of the periodic table.

- Define Le Chatelier’s principle

Le Chatelier’s principle states that “when a stress is applied to a system at equilibrium, the system will react in such a way to counteract the change and establish a new equilibrium”.

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

Change in concentration

If the concentration of any of the reactants is changed, the equilibrium will shift in a direction to oppose the change. If the concentration of a reactant on the LHS is increased, the equilibrium will shift to the right to consume it. If the concentration of a reactant on the RHS is decreased, the equilibrium will shift to the left to produce more of it.

However, the addition of a pure solid or liquid does not affect the concentration and thus will not cause a change in equilibrium.

Temperature

If the temperature of an equilibrium is changed, depending on the enthalpy change of the reaction the equilibrium will shift to oppose that change. For an exothermic equilibrium, an increase in temperature causes a shift to the left.

For an endothermic equilibrium, an increase in temperature causes a shift to the right.

Total pressure

For an equilibrium that has one or more gaseous substance, an increase or decrease in pressure will cause a shift in the equilibrium. If pressure is increased, the equilibrium will shift to the side with lesser gas molecules. If pressure is decreased, the equilibrium will shift to the side with more gas molecules.

- Describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier’s principle

CO2 (g) ⇌ CO2 (aq) ΔH<0

CO2 (aq) + H2O(l) ⇌ H2CO3 (aq)

H2CO3 (aq) ⇌ H+ (aq) + HCO3

- (aq)

HCO3- (aq) ⇌ H+

(aq) + CO32-

(aq)

Concentration: If water is removed from the equilibrium (eg via salting), the equilibrium will shift to produce more CO2 (g). Also, when the cap of a carbonated drink is opened, CO2 (g) escapes so the equilibrium shifts to produce more CO2 (g) to replace it.

Temperature: Since the dissolution process is exothermic, an increase in temperature will favour the left hand side, producing more CO2 (g).

Pressure: When there is a decrease in pressure (eg bottle cap opened), the left side of the equilibrium will be favoured due to Le Chatelier’s Principle.

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

Sulfur dioxide

SO2 is produced by the combustion of fossil fuels with sulfur impurities, volcanoes and bushfires, the action of bacteria on decomposing organic matter, and smelting of iron or other sulfide ores (e.g. ZnS)

Oxides of nitrogen

NOx is produced in high temperature environments such as car engines, through bacterial action or in high energy sparks such as lightning.

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

Sulfur dioxide

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

Roasting of chalcopyrite (copper ore):

2CuFeS2 (s) + 3O2 (g) 2CuS(s) +2FeO(s) + 2SO2 (g)

Roasting of iron sulfide:

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

Oxides of Nitrogen

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

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

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

- Increased combustion of fossil fuels and use of smelters since Industrial Revolution- Use of motor cars and generators increased – NOx increased- Photochemical smog in London and Los Angeles led to emission controls

- Explain the formation and effects of acid rain

Normal rain is slightly acidic due to the dissolved CO2, with a pH of about 5-6.5. Acid rain is rain that has pH lower than 5. When dissolved in water, SO2 and NO2 both lower the pH of rain.

SO2 can react with water to produce sulfurous acid which then reacts with atmospheric oxygen to form sulfuric acid.

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

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

SO2 can alternatively react with atmospheric oxygen to form SO3, which then reacts with water to form H2SO4.

2SO2 (g) + O2 (g) 2SO3 (g)

SO3 (g) + H2O(l) H2SO4 (aq)

NO2 can react with water to produce nitrous acid and nitric acid.

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

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

Effects of acid rain

- Increased acidity of water bodies such as lakes, with adverse effects on aquatic ecology affecting fish population.

- Decreasing the pH of soil affects the growth of plants, including crops and forests. Some plants can only grow in specific pH conditions, and various mineral nutrients in the soil are dissolved by acid rain.

- Acid rain causes corrosion and damage to marble and limestone building and structures and metal structures. Marble and limestone is mainly CaCO3, which reacts with acid rain.

CaCO3 (s) + H2SO4 (aq) CaSO4 (s) + H2O(l) + CO2 (g)

Acids corrode metals, attacking steel structures for example.

Fe(s) + H2SO4 (aq) FeSO4 (s) + H2 (g)

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

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

Acids are substances that donate protons. An acid dissolved in water will donate protons to water, producing a H3O+ hydronium ion and the conjugate base of the acid. Strong acids ionise fully while weak acids partially ionise.

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

Acetic (ethanoic) acid (CH3COOH) is a weak monoprotic acid. It is the major component of white vinegar.

CH3COOH(aq) + H2O(l) ⇋ CH3COO-(aq) + H3O+

(aq)

Citric (2-hydroxypropane-1,2,3-tricarboxylic) acid (C6H8O7) is a weak triprotic acid, found commonly in citric fruits.

Hydrochloric acid (HCl) is a strong monoprotic acid, found in the stomach as gastric juice for digestion.

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

(aq)

Sulfuric acid (H2SO4) is a common diprotic acid, ionising completely in its first stage and partially in its second stage (about 35%).

H2SO4 (aq) + H2O(l) H3O+(aq) + HSO4

-(aq)

HSO4-(aq) + H2O(l) ⇋ H3O+

(aq) + SO42-

(aq)

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

The pH scale is an indicator of H+ ion concentration, which can be used to compare the strength and concentration of acids and bases. Solutions with pH <7 are acidic while solutions with pH >7 are basic. The pH scale is based upon the self-ionisation of water with ionisation constant Kw = 10-14.

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

A strong acid is an acid that is completely ionised.A weak acid is an acid that is only partially ionised.A concentrated acid is an acid solution that has a high concentration of acid particles.A dilute acid is an acid solution that has a low concentration of acid particles.

For example, HCl is always a strong acid even at very low concentrations, which CH3COOH is always a weak acid even at high concentrations.

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

pH=−log10¿

pH+pOH=14

The pH scale is a logarithmic scale, so a difference of one pH is equivalent to a difference of 10 times in H+ or OH- concentration.

- 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 strongest acid between citric, acetic and hydrochloric acid is hydrochloric acid, followed by citric acid and then acetic acid. Hydrochloric acid is a strong acid, ionising fully, while citric acid and acetic acid are weak acids which partially ionise. Citric acid is a stronger acid than acetic acid because it has a higher degree of ionisation.

- 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 such as HCl will completely ionise.

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

(aq)

However, a weak acid will only partially ionise. An equilibrium will be established with both products and reactants present in the final solution. The degree of ionisation determines how strong the weak acid is.

CH3COOH(aq) + H2O(l) ⇋ CH3COO-(aq) + H3O+

(aq)

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

Weak acids are sometimes added to food as a:

- preservative – inhibition of growth of microbes using pH as a control- antioxidant – prevention of spoiling by oxidation- flavouring agent – to improve the flavour of foods/drinks

Some common acids in foods are acetic acid (vinegar), ascorbic acid (vitamin C), citric acid (citrus fruits), lactic acid (milk), propanoic acid (preservative in bread) and benzoic acid (prunes, plums etc).

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

Natural acids: Hydrochloric acid in the stomach, acetic acid in vinegar, citric acid in citrus fruits, ascorbic acid in fruits and vegetables

Natural bases: Ammonia in urine, carbonates such as CaCO3 (limestone)

4. Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refined

- Outline the historical development of ideas about acids including those of:

Lavoisier

Antoine Lavoisier hypothesised in 1776 that acids are substances that contain oxygen. However, this theory was soon disproved, as many bases were found to contain oxygen and many acids without oxygen were found.

Davy

Sir Humphry Davy in 1815 showed that hydrochloric acid did not contain oxygen and proposed that acids are substances that contain hydrogen. In 1838 Justus von Liebig further refined this definition to substances which contained replaceable hydrogen, that release H2 gas when reacting with metals. Although this provided a way to recognise acids, it did not explain why some hydrogen compounds were not acidic.

Arrhenius

Svante Arrhenius proposed in 1884 that acids are substances which ionise in water to produce H+ ions. He also proposed that bases were substance which dissociated to form OH- ions in water. This explained the strength of acids in terms of their ionisation, but did not explain why substances like carbonates and metal oxides act as bases even though they do not dissociate to form OH-. It also did not take into account the solvent in which the substance was dissolved in, and could not define amphoteric substances which can act as both acids and bases.

- Outline the Brönsted-Lowry theory of acids and bases

In 1923 Johannes Brönsted and Thomas Lowry independently proposed a new theory of acids and bases within months of each other. They proposed that acids are proton donors and bases are proton acceptors. This new definition took the solvent into account, focusing on the role of water as an ionising solvent, and accounted for the acidity or basic nature of many more substances such as metal oxides.

- Describe the relationship between an acid and its conjugate base and a base and its conjugate acid

In Brönsted Lowry theory, when an acid donates a proton it forms a conjugate base, and when a base accepts a proton it forms a conjugate acid.

The conjugate of an acid or base has the potential to donate or receive a proton to reform the original acid or base.

The conjugate of a strong acid will be a weak base. The conjugate of a strong base will be a weak acid. The conjugate of a weak acid will be a weak base. The stronger the acid, the weaker its conjugate base will be and vice versa.

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

A salt is an ionic compound that has a cation other than H+ and an anion other than OH- or O2-. In other words, they are formed by the neutralisation of an acid and a base.

An acidic salt is formed by the neutralisation of a strong acid and a weak base, since the cation is a weak acid. For example, NH4

+ is the conjugate acid of the weak base NH3.A basic salt is formed by the neutralisation of a strong base and a weak acid, since the anion is a weak base. For example, CH3CHOO- is the conjugate base of the weak acid CH3COOH.A neutral salt is formed by the neutralisation of a strong base and strong acid, or a weak base and weak acid.

- Identify conjugate acid/base pairs

The conjugate base of an acid is the acid with one less proton, and the conjugate acid of a base is the base with one more proton.

e.g. HNO3/NO3-, H2O/OH-, H2O/H3O+, CH3COOH/CH3COO-

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

Amphoteric substances are substances that can react as both acids and bases. Some examples of amphoteric substances are Al2O3 and ZnO.

Al2O3 (s) + 6HCl(aq) 2AlCl3 (aq) + 3H2O(l)

Al2O3 (s) + 2NaOH(aq) 2NaAlO2 (aq) + H2O(l)

Amphiprotic substances are a specific type of amphoteric substances, those which act as both B/L acids and B/L bases. In other words, they act as both proton donors and proton acceptors. Common amphiprotic substances include H2O, HCO3

- and HPO42-. They must be able to both accept and donate a proton.

HCO3- as a B/L acid:

HCO3-(aq) + OH-

(aq) CO32-

(aq) + H2O(l)

HCO3- as a B/L base:

HCO3-(aq) + H+

(aq) H2CO3 (aq)

- Identify neutralisation as a proton transfer reaction which is exothermic

Neutralisation is a reaction in which protons are transferred from the acid to the base. The resulting products have a lower enthalpy that the reactants, releasing heat. Neutralisation is thus an exothermic process.

- Describe the correct technique for conducting titrations and preparation of standard solutions

A standard solution is a solution with an accurately known concentration. A primary standard is one made from a highly pure solid (e.g. anhydrous Na2CO3 or NaHCO3). The primary standard must be:

- Available in highly pure form- Accurately known chemical composition- Of high molecular weight to reduce the error in weighing

- Soluble in water- Stable and unaffected by air when weighing

Preparation of primary standard:

The required mass of primary standard is weighed out directly into a clean dry beaker. It is then dissolved in a small amount of water using a wash bottle. This is added to a volumetric flask cleaned with distilled water with a wash bottle, making sure to transfer all of the solution to the flask. The flask is then filled with distilled water to the mark, and the flask stoppered and inverted several times to mix the contents thoroughly.

Titration:

Both the burette and pipette must be washed with distilled water and then the solution that will be used in them, while the conical and volumetric flasks are to be washed with distilled water only. The burette is filled with the solution of known concentration, called the titrant. The solution to be analysed is placed in a flask underneath the burette using the pipette, adding one or two drops of suitable indicator. The titrant is slowly run from the burette into the flask until the colour of the indicator changes indicating the end point of the titration. The volume on the burette is read as accurately as possible, and this is repeated several times and an average taken. The concentration of the unknown is then calculated using this information.

The equivalence point is the point at which all the moles of H+ of the acid have reacted with an equivalent number of moles of OH- of the base. The indicator chosen needs to be one that will show the end point, the point when the indicator changes its colour and the titration stopped, as close to the equivalence point as possible. Equivalence can occur at high or low pH, depending on the acids and bases used. A strong acid and a weak base will have an equivalence point at low pH, and a weak acid and strong base will have a high equivalence point. However, the neutralisation of a strong acid and strong base will have a neutral equivalence point.

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

A buffer is a solution which resists changes in pH when a small amount of acid or base is added. A buffer solution contains a comparable amount of a weak Brönsted-Lowry acid and its conjugate base (e.g. CH3COOH/CH3COO-). An example of a natural buffer is H2CO3/HCO3

- in blood, keeping the pH range of blood between 6.4 and 7.4.

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

(aq)

When an acid is added to the solution, the concentration of H3O+ ions increases. In accordance with Le Chatelier’s principle, the equilibrium is shifted towards the left, and the added H3O+ ions are consumed by reacting more HCO3

-. The effect on the pH is thus minimised.

When a base is added to the solution, the concentration of OH- ions increases, reacting with H3O+ and removing it from equilibrium. As a result, the concentration of H3O+ ions decreases and the equilibrium is shifted to right, reacting more H2CO3 with H2O and replacing the removed H3O+. The effect of the base on the pH of the solution is thus minimised.

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

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

The alkanol functional group is –OH, while the alkanoic acid functional group is –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

To name alkanols, the usual naming conventions apply but with the suffix –ol, while the same applies with alkanoic acids but with the suffix –anoic acid.

To name the ester, the number of carbon atoms from the alkanol is named as the alkyl group (e.g. butanol becomes butyl), and the number of carbon atoms from the alkanoic acid is named as the alkanoate chain (e.g. propanoic acid becomes propanoate).

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

The intermolecular forces displayed in both alkanols and alkanoic acids are dispersion forces and hydrogen bonding, due to the OH group in both. This means that their melting and boiling points are higher than their parent alkanes. Alkanoic acids have greater molecular weights than their equivalent alkanols, leading to higher dispersion forces in alkanoic acids. Furthermore, alkanoic acids are able to form a greater degree of hydrogen bonding than alkanols, hence alkanoic acids have higher boiling and melting points than their equivalent alkanols.

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

alkanol(aq) + alkanoic acid(aq) ⇋ alkyl alkanoate(aq) + H2O(aq)

- Describe the purpose of using acid in esterification for catalysis

Concentrated sulfuric acid is used in esterification, acting as a catalyst to speed up the reaction. It also acts as a dehydrating agent which absorbs water, shifting the equilibrium to the right and thus increasing the yield of the ester. The acid not only increases rate of reaction, but also increases the yield.

- Explain the need for refluxing during esterification

To speed up the reaction, the reaction mixture is heated to a higher temperature (usually near the boiling point of the alkanol). Refluxing is the process of heating a mixture of liquids in a flask with an attached condenser to prevent the loss of volatile reactants and products, allowing the reaction to be carried out a higher temperature. Boiling chips are used to promote even boiling, and a hot water bath is used as an indirect heat source rather than a naked flame to avoid the combustion of the flammable reactants or products.

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

ethyl ethanoate – solvent for paint and lacquerspropyl ethanoate – volatile solvent, wood lacquerpentyl propanoate – solvent, appliance coatings, enamels, lacquers

ethyl methanoate – artificial rum flavouringpentyl enthanoate – pear essence/flavourmethyl butanoate – apple essence/flavouriso-pentyl ethanoate – banana essence/flavouroctyl ethanoate – orange essence/flavour