CHEM2: 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

• An acid is a substance which in solution produces hydrogen ions, H+, or more strictly hydronium ions, H3O+. • A base is a substance which either contains the oxide O2- or the hydroxide ion, OH--. That, or a substance that

when in solution produces the hydroxide ion. • A neutral solution is one in which the hydrogen ion concentration and hydroxide ion concentration is equal.

Properties of: Acids:

• Sour taste • Sting or burn the skin • Conduct electricity in solution • Turns blue litmus red.

Bases: • Bitter taste • Slippery or soapy feel • Conduct electricity in solution • Turns red litmus blue.

Common Substances: Acids:

• Vinegar is ethanoic (acetic) acid • Fruits such as lemons/grapefruit contain citric acid and ascorbic acid (vitamin

C) • Acetic, citric, ascorbic and benzoic acid are used as food additives

o Increase nutritional value, enhance flavour o Reduce pH so that microorganisms that cause spoilage can no longer

reproduce • Hydrochloric acid: in stomach acid, used by builders to clean bricks, and for

pH maintenance in swimming pools • Sulfuric acid: used to make fertilisers, in manufacture of soaps/detergents,

catalyst for many reactions • Nitric acid: used in production of pesticides and explosives • Carbonic acid: found in carbonated beverages • Lactic acid: found in milk and muscle tissue

Bases:

• Sodium hydroxide: used in manufacture of soap, and in drain cleaners • Ammonia: found in urine, used to make fertilisers including ammonium sulphate, and in household cleaners • Magnesium hydroxide Mg(OH)2, Aluminium hydroxide Al(OH)3: used in antacid preparations for relief of upset

stomachs and gastric reflux

SOLVE PROBLEMS BY APPLYING INFORMATION ABOUT THE COLOUR CHANGES OF INDICATORS TO CLASSIFY SOME HOUSEHOLD SUBSTANCES AS ACIDIC, NEUTRAL OR BASIC Acids: vinegar, fruits, milk, additives in other foods, soft drinks Bases: soap, detergent, drain cleaner, bleach, baking soda Prac Method:

1) Take about 1 mL of household substance and add a few drops of universal indicator 2) Compare the colour produced to the pH colour chart

Prac Conclusion: • Bleach, Morning Fresh, shampoo, lemon, hand wash, ammonia, dishwasher rinse, industrial soap, vinegar, milk • Household substances for human consumption are generally acidic. Products which come into contact with

human skin are mildly acidic or basic to avoid skin damage. Cleaning substances are generally very basic. 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

• Many acid-base indicators are made from pigments extracted from leaves/flowers of plants • Litmus is extracted from various species of lichens • Indicators change colour over different acidity ranges. Therefore different indicators can be used in different

situations. • Universal indicator is a mixture of several indicators and

undergoes a series of colour changes over a large range of pH values

Indicator Colour change pH range

Methyl orange Red-yellow 3.1-4.4

Bromothymol blue Yellow-blue 6.2-7.6

Litmus Red-purple 5.0-8.0

Phenolphthalein Colourless-red 8.3-10.0

IDENTIFY AND DESCRIBE SOME EVERYDAY USES OF INDICATORS INCLUDING THE TESTING OF SOIL ACIDITY/BASICITY

• Maintenance of acidity levels (pH 7.2-7.6) in swimming pools o Sample of pool water is collected in a vial, and phenol red is added

Acidity levels need to be controlled to avoid skin/eye irritation • Used by farmers to test acidity of soil

o Soil is moistened with water and indicator added. A white neutral powder, barium sulfate, is placed on the soil so the colour change can be observed. If soil is acidic, chemicals such as lime (calcium oxide) are added to make soil more effective for

growing particular crops • Monitoring industrial wastes • Monitoring wastes from photographic processing • To indicate completion of acid-base titrations

PERFORM A FIRST-HAND INVESTIGATION TO PREPARE AND TEST A NATURAL INDICATOR

• Anthocyanin: obtained by boiling red cabbage leaves and pouring off red/purple liquid o Red in acidic solutions, green in mildly basic solutions, yellow in very basic solutions

IDENTIFY DATA AND CHOOSE RESOURCES TO GATHER INFORMATION ABOUT THE COLOUR CHANGES OF A RANGE OF INDICATORS

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

• When non-metal oxides react with water, acids are produced o An acidic oxide is one that either:

Reacts with water to form an acid, or Reacts with bases to form salts

• Some oxides of non-metals act as acids when in solution o Nitrogen dioxide Nitric & nitrous acid: 2 NO2(g) + H2O(l) HNO3(aq) + HNO2(aq)

o Sulfur trioxide Sulfuric acid: SO3(g) + H2O(l) H2SO4(aq)

o Phosphoric acid: P4H10(s) + 6 H2O(l) 4 H3PO4(aq) ???

o Perchloric acid: Cl2O7(l) + H2O(l) 2 HClO4(aq)

Except for CO, NO and N2O which are neutral • Other oxides of non-metals can act as acids in acid-base reactions

o SO2(g) + 2 NaOH(aq) Na2SO3(aq) + H2O(l) Sulfur dioxide + sodium hydroxide sodium sulfite + water

o SiO2(s) + 2 NaOH(aq) Na2SiO3(aq) + H2O(l) Silicon dioxide + sodium hydroxide sodium silicate + water

ANALYSE THE POSITION OF THESE NON-METALS IN THE PERIODIC TABLE AND OUTLINE THE RELATIONSHIP BETWEEN POSITION OF ELEMENTS IN THE PERIODIC TABLE AND ACIDITY/BASICITY OF OXIDES

• Oxides of elements are increasingly acidic

going from left to right across a period in the periodic table

• Upper-right hand corner acidic oxide o Sulfuric acid and Perchloric acid are

strongly acidic; silicon dioxide is weakly acidic

• Lower-left hand corner basic oxide o Sodium oxide, magnesium oxide are bases [ionic compounds]

Na2O + 2 HCl 2 NaCl + H2O DEFINE LE CHATELIER’S PRINCIPLE “If a chemical system at equilibrium is subjected to a change in conditions, the system will adjust to re-establish equilibrium in such a way as to partially counteract the imposed change”

• Systems at equilibrium have constant concentrations of reactants and products, and equal/opposing reaction rates

• If changes are made to the conditions, the system may no longer be at equilibrium, and the system will tend to re-establish equilibrium by favouring one side of the equation.

• Used in industry to improve the yield of processes by changing conditions under which the reaction is carried out

IDENTIFY FACTORS WHICH CAN AFFECT THE EQUILIBRIUM IN A REVERSIBLE REACTION Changing Concentration:

• Le Chatelier’s principle predicts that increasing concentration of one of the substances involved will cause the system to favour the direction that will decrease the concentration of that substance, and vice versa

• 𝑁2𝑂4(𝑔) ↔ 2𝑁𝑂2(𝑔)

o E.g. if [NO2] was increased, system would re-establish equilibrium by partially decreasing NO2

concentration. The reverse reaction would be favoured. o Final NO2 concentration would be less than after addition of NO2, but more than at initial equilibrium

• Note that changing the quantity of solid or liquid species has no effect on position of equilibrium, because concentrations of these solids and liquids present remain constant

Changing volume or pressure of a gaseous system: • Increase in external pressure will lead to decrease in volume, with resulting increase in concentrations of all

components o This results in the system no longer being at equilibrium only if there are different numbers of moles of

gaseous reactant and product molecules in the balanced chemical equation. • 𝑁2𝑂4(𝑔) ↔ 2𝑁𝑂2(𝑔)

o E.g. if volume of the system is halved by doubling external pressure, total concentration of particles in the system is increased, system will adjust to re-establish equilibrium in such a way as to partially decrease the concentration of gaseous particles. The reverse reaction would be favoured, leading to a decrease in total number of particles

o When equilibrium is re-established, there will be more N2O4 and less NO2 than after the change was made, but concentrations of both are greater than at initial equilibrium

• Note that when an equilibrium system contains equal numbers of gaseous molecules on both sides of the equation, changing volume doesn’t affect equilibrium

Changing temperature:

• 𝑁2𝑂4(𝑔) + 57 𝑘𝐽 ↔ 2𝑁𝑂2(𝑔)

o The forward reaction is endothermic and absorbs heat; reverse reaction releases heat

o E.g. if temperature of the system is increased, system will adjust to re-establish equilibrium in such a way as to decrease temperature. The forward reaction is favoured, as it absorbs heat

o Results in an increase in [NO2] and decrease in [N2O4] o Since N2O4 is a colourless gas while NO2 is a dark brown gas, the

mixture is dark brown at higher temperatures, and a lighter colour when cooler

Adding a catalyst:

• Catalyst lowers activation energy for both forward and reverse reactions, and increases rate for both reactions equally

• Doesn’t affect equilibrium, but allows system to reach equilibrium in less time DESCRIBE THE SOLUBILITY OF CARBON DIOXIDE IN WATER UNDER VARIOUS CONDITIONS AS AN EQUILIBRIUM PROCESS AND EXPLAIN IN TERMS OF LE CHATELIER’S PRINCIPLE

• Carbon dioxide establishes an equilibrium between gaseous and dissolved CO2

𝐶𝑂2(𝑔) ↔ 𝐶𝑂2(𝑎𝑞)

• Dissolved CO2 reacts with water to form weakly acidic solution of carbonic acid

𝐶𝑂2(𝑎𝑞) + 𝐻2𝑂(𝑙) ↔ 𝐻2𝐶𝑂3(𝑎𝑞)

• Carbonic acid establishes equilibria involving hydrogencarbonate and carbonate ions

𝐻2𝐶𝑂3(𝑎𝑞) ↔ 𝐻+(𝑎𝑞) + 𝐻𝐶𝑂3−(𝑎𝑞)

𝐻𝐶𝑂3−(𝑎𝑞) ↔ 𝐻+(𝑎𝑞) + 𝐶𝑂32−(𝑎𝑞)

• Rain water has absorbed carbon dioxide from the atmosphere, and hence is slightly acidic

General net reaction*: 𝐶𝑂2(𝑔) + 𝐻2𝑂(𝑙) ↔ 𝐻2𝐶𝑂3(𝑎𝑞)

Effect of carbon dioxide pressure on solubility: • In a closed system such a bottle of soft drink, gaseous CO2 is in equilibrium with dissolved CO2, which is in

equilibrium with H2CO3 𝐶𝑂2(𝑔) ↔ 𝐶𝑂2(𝑎𝑞)

𝐶𝑂2(𝑎𝑞) + 𝐻2𝑂(𝑙) ↔ 𝐻2𝐶𝑂3(𝑎𝑞)

• Gas above solution in bottle has increased concentration of carbon dioxide (4-5 times normal atmospheric pressure). This favours the forward reaction by Le Chatelier’s Principle, and thereby increases solubility of the gas, and acidity of solution

• When bottle of soft drink is open, pressure of CO2 gas decreases and the equilibrium adjusts to favour the reaction that increases CO2 pressure. As a result, bubbles of CO2 gas come out of solution and the drink goes ‘flat’

Effect of temperature on solubility: Dissolving of CO2 in water is exothermic

𝐶𝑂2(𝑔) ↔ 𝐶𝑂2(𝑎𝑞) + heat

Therefore, increasing temperature favours the endothermic reverse reaction, leading to release of CO2 from solution

o Heating a soft drink will accelerate rate at which solution becomes ‘flat’ Effect of acidity (concentration of H + ions) on carbon dioxide solubility: Carbonic acid establishes equilibria involving hydrogencarbonate and carbonate ions

𝐻2𝐶𝑂3(𝑎𝑞) ↔ 𝐻+(𝑎𝑞) + 𝐻𝐶𝑂3−(𝑎𝑞)

𝐻𝐶𝑂3−(𝑎𝑞) ↔ 𝐻+(𝑎𝑞) + 𝐶𝑂32−(𝑎𝑞)

If an acid is added e.g. vinegar, the additional 𝐻+(𝑎𝑞) drives both equilibria to the left

Increased 𝐻2𝐶𝑂3(𝑎𝑞) causes release of CO2 from solution 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 IDENTIFY NATURAL AND INDUSTRIAL SOURCES OF SULFUR DIOXIDE AND OXIDES OF NITROGEN a n d DESCRIBE, USING EQUATIONS, EXAMPLES OF CHEMICAL REACTIONS WHICH RELEASE SULFUR DIOXIDE AND CHEMICAL REACTIONS WHICH RELEASE OXIDES OF NITROGEN Sulfur dioxide: Natural

• Sulfur is present in organic matter as part of proteins. Bacteria decompose organic matter to produce hydrogen sulphide (H2S), which oxidises to form sulfur dioxide

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

• Volcanic gases released during eruptions and from geysers • Combustion of organic matter, e.g. bushfires

Sulfur dioxide: Industrial • Burning of fossil fuels (coal and petroleum) in power plants and motor vehicles

o When coal is burnt, iron sulphide (FeS2) impurities oxidise to form sulfur dioxide 4 FeS2(s) + 11 O2(g) 2 Fe2O3(s) + 8 SO2(g)

• Smelting of sulphide ores in industrial plants releases sulfur dioxide into atmosphere

o E.g. roasting of zinc sulphide in extraction of zinc 2 ZnS(s) + 3 O2(g) 2 ZnO(s) + 2 SO2(g)

• Incineration of garbage • Petroleum refineries • Industries using sulfur dioxide for production of sulfuric acid, or

other products Nitrogen oxides: Natural

• NO (nitric oxide) and N2O (nitrous oxide) produced by soil bacteria

• NO produced in high-temperature environments, such as during lightning strikes N2(g) + O2(g) 2 NO(g)

• NO2 (nitrogen dioxide) produced by action of sunlight on NO 2 NO(g) + O2(g) 2 NO2(g)

Nitrogen oxides: Industrial • NO produced from burning of biomass • NO and NO2 produced in combustion of fuel in motor vehicles and power stations

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

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

• N2O manufactured as a fuel for racing cars and for use as sedative/analgesic 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 Reasons for Concern about their Release into the Environment

• Concentrations of oxides of sulfur and nitrogen has increased as a result of human activity, particularly since onset of Industrial Revolution

• Sulfur oxides are irritating gases which present a major health risk, especially for people with respiratory disorders

o London smogs of 1950s were caused by smoke and sulfur dioxide, and caused many deaths o Emission of sulfur dioxide has increased considerably since beginning of Industrial Revolution, due to

burning of fossil fuels (especially coal) and smelting of sulfide ores o E.g. copper mining in Queenstown, Tasmania involved roasting copper sulfide ores, and released sulfur

dioxide fumes which killed off plant growth and produced acid rain, which also destroyed vegetation. • Nitrogen oxides are important components in formation of

photochemical smog and acid rain

Sulfur dioxide, sulfur trioxide and nitrogen dioxide are acidic oxides which combine with water to produce airborne acid droplets, i.e. acid rain Reducing emissions

• Reducing sulfur dioxide emissions o Reducing sulfur content of fuels (OZ coal) o Using a scrubber to remove sulfur dioxide

Scrubber passes gaseous emissions through a slurry of lime (CaO), which neutralises the sulfur dioxide to form calcium sulfite

• SO2(g) + CaO(s) CaSO3(s)

Calcium sulfite oxidised by oxygen to form gypsum • 2CaSO3(s) + O2(g) 2CaSO4(s)

• Reducing nitric oxide emissions o Catalytic converters catalyse decomposition of NO back to N2 and O2,

and also decompose carbon monoxide into carbon dioxide o Rhodium or platinum catalyst

• 2NO(g) + 2CO(g) N2(g) + 2CO2(g)

EXPLAIN THE FORMATION AND EFFECTS OF ACID RAIN Formation:

• Unpolluted rain is slightly acidic (6-6.5 pH) due to carbon dioxide dissolved in water to form carbonic acid CO2(aq) + H2O(l) H2CO3(aq)

• Acid rain: rain with pH < 5 o Caused by rain dissolving sulfur oxides and nitrogen dioxide, so the rain becomes dilute solution of

various acids • Sulfur dioxide is readily oxidised in air to form sulfur trioxide

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

• When dissolved in rain, sulfur and nitrogen oxides produce acids and hence acid rain o Nitric acid: 4 NO2(g) + 2 H2O(l) + O2(g) 4 HNO3(aq)

o Nitric and nitrous acid: 2 NO2(g) + H2O(l) HNO3(aq) + HNO2(aq)

o Sulfuric acid from sulfur trioxide: SO3(g) + H2O(l) H2SO4(aq)

o Sulfurous acid from sulfur dioxide: SO2(g) + H2O(l) H2SO3(aq)

Effects: • Causes water in lakes to become too acidic to support fish and other aquatic life, which have a limited pH

range for survival o At lower pH’s, fish eggs can’t hatch and adult fish may die o Acidity leads to reduced calcium uptake in fish, causing reduced skeletal growth o Reduction in fish size/population affects animals that prey on the fish, and ultimately affects whole

ecosystem • Causes damage to plant foliage, including crops and forests

o Partly due to changes in soil pH resulting in reduced productivity o Acids dissolve and remove nutrients important for plant growth

2 H+(aq) + Mg2+

(clay) 2 H+(clay) + Mg2+

(aq)

o Acids dissolve minerals, releasing toxic heavy metals • Causes damage to metal/stone buildings, structures, and statues

o Sulfuric acid in rain causes destruction of limestone and marble buildings/statues • CaCO3(s) + H2SO4(aq) Ca2+

(aq) + SO42-

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

As water evaporates in rock crevices, calcium sulfate (gypsum) can crystallise out, causing rocks to crumble

Gypsum forms a surface to which dirt/soot particles can readily bind o Iron in buildings can react with hydrogen ions to produce hydrogen gas

• Fe(s) + 2H+(aq) Fe2+

(aq) + H2(g) • Bronchitis and asthma in humans are worsened by acid rain

ASSESS THE EVIDENCE WHICH INDICATES INCREASES IN ATMOSPHERIC CONCENTRATION OF OXIDES OF SULFUR AND NITROGEN History:

• After the industrial revolution of the early 1800’s there was a great increase in the emissions of sulfur dioxide. This continued to rise until the 1950’s and 1960’s, when deaths were recorded as a result of the air pollution.

• London smogs of 1952 caused by coal soot and sulfur dioxide from burning coal resulted in death of 4000 people. Stringent regulations were brought in after.

_________________

The Evidence: • Increases in atmospheric concentrations of these chemicals have cause acid rain, and resulting damage to

ecosystems and structures • Extensive evidence for an increase of over 25% in atmospheric carbon dioxide levels over the last two hundred

years. o Quantitative analysis of trapped air bubbles in Antarctic ice, measurement of carbon isotopes in old

trees • However, finding evidence for increases in atmospheric sulfur oxides and nitrogen oxides is more difficult for the

following reasons: o Sulfur dioxide and nitrogen dioxide form ions that dissolve in water making them difficult to measure. o Levels for SO2 and NOx are only about 0.001 ppm in populated parts of the Earth. o The chemical instruments able to measure very low concentrations have only been commercially

available since the 1970s. 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

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 a n d GATHER AND PROCESS INFORMATION FROM SECONDARY SOURCES TO WRITE IONIC EQUATIONS TO REPRESENT THE IONISATION OF ACIDS

• Ionise: To dissociate atoms or molecules into electrically charged species (ions) • The ionisation of an acid in water can be summarised as the dissociation of substance into ions, which form

hydrogen ions in water. Simplistic model: HCl(g) → H+

(aq) + Cl-(aq)

• Bronsted-Lowry Theory: An acid is a proton donor; a base is a proton acceptor HCl(g) + H2O(l) → H3O+

(aq) + Cl-(aq)

o HCl is donating a proton and is therefore acting as an acid. H2O is accepting a proton and therefore acting as a base

o When an acid molecule is in contact with water it can ionise, donating a proton to a water molecule to make a hydronium ion H+

(aq) + H2O(l) → H3O+(aq)

• Sulfuric acid is called a diprotic acid because each molecule can release up to two protons.

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

IDENTIFY ACIDS INCLUDING ACETIC (ETHANOIC), CITRIC (2-HYDROXYPROPANE-1,2,3-TRICARBOXYLIC), HYDROCHLORIC AND SULFURIC ACID a n d IDENTIFY DATA, GATHER AND PROCESS INFORMATION FROM SECONDARY SOURCES TO IDENTIFY EXAMPLES OF NATURALLY OCCURRING ACIDS AND BASES AND THEIR CHEMICAL COMPOSITION Acetic acid (ethanoic acid): CH3COOH

• Occurs naturally in the decomposition of biological material, such as when wine naturally ferments.

• Most acetic acid used by humans is manufactured industrially. o Vinegar is 4% acetic acid

Citric acid (2-hydroxypropane-1,2,3-tricarboxylic acid): C6H8O7

• Occurs naturally in citrus fruits • Formed during cellular respiration in body cells • Added to food as preservative

Hydrochloric acid: HCl • Stomach acid

o High acidity activates enzymes in stomach catalysing protein breakdown • Produced industrially on a large scale

Sulfuric acid: H2SO4 • Can be produced from sulfur dioxide emissions of volcanic eruptions • Can be produced from smelting of sulfide ores

o Most sulfur dioxide released into the earth’s atmosphere is oxidised and dissolved in water to form the sulfuric acid in acid rain.

GATHER AND PROCESS INFORMATION FROM SECONDARY SOURCES TO EXPLAIN THE USE OF ACIDS AS FOOD ADDITIVES

• Acids are used as food additives o Increases nutritional value o Enhance flavour of the food o Prevent spoilage by reducing pH to a level where microorganisms that cause spoilage can no longer

reproduce ____________________________

• Sulfur dioxide: SO2 o Added to dried fruit and wine to prevent decolourisation

Strongest acid used as food additive: Can cause severe reactions in asthmatics Can be added as gaseous SO2, solution in water as H2SO3, sulfite or bisulfite

• Lactic acid: CH3CHOHCOOH o Added to cheese to balance acidity o Added to carbonated beverages to add tartness

• Acetic acid: CH3COOH o Added to ice cream as a thickening/sterilising agent, and to prevent sugar from crystallising

Main ingredient in vinegar

• Ascorbic acid o Added to read meat to maintain red colour and flavour of red meat by reacting with unwanted oxygen

(antioxidant) • Citric acid:

o Added to carbonated beverages for tart flavouring, and as an antioxidant o Sodium citrate added to fruit drinks, as a buffer to control acidity o Coating freshly cut fruit such as apples with lemon juice (citric acid) inactivates natural enzymes that

cause ‘browning’ DESCRIBE THE USE OF THE PH SCALE IN COMPARING ACIDS AND BASES

• Auto-ionisation of water 2H2O(l) ↔ H3O+

(aq) + OH-(aq)

o Ionisation constant at 25⁰C for water is given by: Kw = [H+] [OH-] = 1.0 x 10-14

o In pure water, [H+] = [OH-] = 1.0 x 10-7 mol L-1

o In acidic solutions, [H+] is greater than 1.0 x 10-7 mol L-1

and [OH-] is less than 1.0 x 10-7 mol L-1

• The pH scale allows us to measure acidity of various samples by comparing concentrations of H+ and OH- ions in solutions

o pH scale allows us to indicate acidity without using large indices o A value less than 7 indicates a solution is acidic; a value greater than 7 indicates a solution if

basic/alkaline o pH stands for ‘potential Hydrogen’

IDENTIFY PH AS -LOG10 [H+] AND EXPLAIN THAT A CHANGE IN PH OF 1 MEANS A TEN-FOLD CHANGE IN [H+] a n d PROCESS INFORMATION FROM SECONDARY SOURCES TO CALCULATE PH OF STRONG ACIDS GIVEN APPROPRIATE HYDROGEN ION CONCENTRATIONS

• pH = - log10[H+]

o This equation can be used to calculate pH of strong acids given hydrogen ion concentrations

• [H+] = 10-pH

• Therefore a change of one pH unit represents a tenfold change in hydrogen ion concentration • In a strong acid solution, each acid molecule is assumed to fully ionise. The concentration of hydrogen

ions and hence pH will depend on whether the acid is monoprotic (releases one proton per molecule), diprotic (releases two protons per molecule) or triprotic (releases three protons per molecule).

• E.g. H2SO4:

H2SO4 2H+ + SO42–

[H+] = 2 x concentration of H2SO4

DESCRIBE THE DIFFERENCE BETWEEN A STRONG AND A WEAK ACID IN TERMS OF AN EQUILIBRIUM BETWEEN THE INTACT MOLECULE AND ITS IONS Strong Acids

• Strong acids are completely ionised to produce hydrogen ions in aqueous solution o A pointed arrow () is used to show this

• E.g. hydrochloric acid HCl(g) H+

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

(aq) + Cl-(aq) o Bronsted-Lowry: HCl is a stronger acid (proton donor) than H3O+, and H2O is a stronger base than Cl-, so

forward reaction is favoured • Hydrochloric, sulfuric and nitric acid are all better proton donors than H3O+ ion, and hence are strong acids

Weak Acids • Weak acids are only partially ionised in water; an equilibrium is established between intact acid molecules and

its ions o A double arrow (↔) is used to show this

• E.g. acetic acid CH3COOH(aq) + H2O(l) ↔ H3O+

(aq) + CH3COO-(aq)

o Equilibrium established between acetic acid molecules, and hydronium and ethanoate ions o Bronsted-Lowry: CH3COOH Is a weaker acid (proton donor) than H3O+, and H2O Is a weaker base than

CH3COO-, • Aquated metal ions can donate a proton from one of their surrounding water molecules (e.g. Al3+ ion)

Strong Bases • Strong bases completely dissociate to produce hydroxide ions in aqueous solution • E.g. potassium hydroxide

KOH(s) K+(aq) + OH-

(aq) o Group I/II metal hydroxides are strong bases o Group II metal hydroxides produce two moles of hydroxide ion for every mole of metal hydroxide, but

have limited solubility • Metallic oxides are also strong bases; oxide ion is stronger base than hydroxide ion

Na2O(s) + H2O(l) 2Na+(aq) + 2OH-

(aq) Weak Bases

• Only a small proportion of molecules or ions react with water to form hydroxide ions in aqueous solution • E.g. ammonia ion

NH3(aq) + H2O(l) ↔ NH4+

(aq + OH-(aq)

o Bronsted-Lowry: NH4+ is a stronger acid than H2O and OH- is a stronger base than NH3, so reaction

occurs to a small extent • Many weak bases are anions such as carbonate, acetate, fluoride and phosphate, which can accept protons

from water to a small extent, producing hydroxide ions Generally, strong acids will react with strong bases to form weaker conjugate acids and weaker conjugate bases. i.e. acids on the ‘strength of acids/bases’ table will react with bases below them. DESCRIBE ACIDS AND THEIR SOLUTIONS WITH THE APPROPRIATE USE OF THE TERMS STRONG, WEAK, CONCENTRATED AND DILUTE

• ‘Concentrated’ and ‘dilute’ describe concentrations of solutions in terms of amount of solute added o A concentrated solution is one in which total concentration of a solute is high, i.e. there is a large

amount of solute in a given amount of solution o A dilute solution is one in which total concentration of a solute is low, i.e. there is a small amount of

solute in a given amount of solution • ‘Strong’ and ‘weak’ describe degree to which substance ionises or dissociates into ions

o A strong acid is one in which virtually all the acid present ionises to form hydrogen ions o A weak acid is one in which only some of the acid molecules present ionise to form hydrogen ions

PLAN AND PERFORM A FIRST-HAND INVESTIGATION TO MEASURE THE PH OF IDENTICAL CONCENTRATIONS OF STRONG AND WEAK ACIDS USE AVAILABLE EVIDENCE TO MODEL THE MOLECULAR NATURE OF ACIDS AND SIMULATE THE IONISATION OF STRONG AND WEAK ACIDS Use molecular model kits. 1) Make 4 HCl molecules, 4 citric acid molecules and eight water molecules. 2) First remove the hydrogen from HCl and attach it to the water molecule to model the hydronium ion. Do this 4 times. This will simulate complete ionisation. 3) Remove one hydrogen from citric acid and attach to water. Do this only once and it will model incomplete ionisation. 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 *all 0.01 mol L-1 Name pH [H+] Hydrochloric acid 2.00 0.01 Citric acid 2.56 2.74 x 10-3 Acetic acid 3.38 4.17 x 10-4

• Order of [H+]: HCl > citric acid > acetic acid • Hydrochloric acid is a strong acid, and almost completely ionises into ions

o Concentration of H+ is equal to concentration of acid • Citric acid is a weaker acid that is only slightly ionised in solution (~1%)

o Concentration of H+ is less than concentration of acid, and so pH is higher o Citric acid is a triprotic acid with three ionisable hydrogen atoms, which contributes to a higher [H+],

resulting in lower pH than acetic acid C6H8O7 + H2O ↔ C6H7O7

- + H3O+

C6H7O7 + H2O ↔ C6H6O72- + H3O+

C6H6O7 + H2O ↔ C6H5O73- + H3O+

• Acetic acid is an even weaker acid which is ionised even less in solution (1%) CH3COOH + H2O ↔ CH3COO- + H3O+

SOLVE PROBLEMS AND PERFORM A FIRST-HAND INVESTIGATION TO USE PH METERS/PROBES AND INDICATORS TO DISTINGUISH BETWEEN ACIDIC, BASIC AND NEUTRAL CHEMICALS

4. Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refined OUTLINE THE HISTORICAL DEVELOPMENT OF IDEAS ABOUT ACIDS INCLUDING THOSE OF: - LAVOISIER - DAVY - ARRHENIUS Antoine Lavoisier (1780s): Oxygen

• Non-metallic oxides form acids when dissolved in water o E.g. acetic, carbonic, sulfuric, nitric acid all contained oxygen o Concluded that the presence of oxygen in compounds formed from non-metals caused acidity

Sir Humphrey Davy (1815): Hydrogen • Decomposed hydrochloric acid and found that only hydrogen and chlorine were produced

o Therefore, not all acids contained oxygen • Proposed that acids contain hydrogen, that could be replaced by reaction with a metal

o Other acids with hydrogen rather than oxygen were discovered: HF, HBr, HI, HCN Svante Arrhenius: H+/OH- in solution

• Proposed that acid is a substance that produces hydrogen ions (H+) in aqueous solution, and that a base is a substance that produces hydroxide ions (OH-) in aqueous solution

• Strong acids completely ionise in aqueous solution o HCl(aq) H+

(aq) + Cl-(aq)

o NaOH(aq) Na+(aq) + OH-

(aq)

• Evidence: o Acid solutions are conductive o Acids react with metals to produce hydrogen gas

• Limitations of Arrhenius definition of acids and bases o Many substances that behave as bases, (NH3, Na2CO3) do not contain an OH group o Only applicable to aqueous solutions o Doesn’t address relative strengths of different acids/bases, or explain why some substances can act as

either acids or bases (amphoterism) ________________________ Extension on Arrhenius’ work:

• Monoprotic acids contain only one acidic hydrogen atom per molecule of acid o E.g. hydrochloric, nitric, acetic acid

• Diprotic acids contain two ionisable hydrogen atoms per molecule of acid o E.g. Sulfuric, carbonic acid o Sulfuric acid solution: first proton is completely ionised, but only a small proportion of hydrogensulfate

ions ionise further into hydrogen and sulphate ions H2SO4(l) H+

(aq) + HSO4-(aq)

HSO4-(aq) ↔ H+

(aq) + SO42-

(aq)

o When reacting with a strong base (e.g. NaOH), one mole of sulfuric acid will react with 2 moles of hydroxide ions 2OH-

(aq) + H2SO4(aq) 2H2O(l) + SO42-

(aq)

• Triprotic acids contain three ionisable hydrogen atoms per molecule of acid o E.g. Phosphoric acid (H3PO4)

… HPO42-

(aq) ↔ H+(aq) + PO4

3-(aq)

o One mole of acid will react with three moles of sodium hydroxide • Generally, hydrogen atoms are only acidic if attached to electronegative atoms, most commonly oxygen • A hydrogen ion is simply a proton. But because the proton’s charge is located in such a small volume, attraction

between polar water molecule and proton is greater than for other ions, and proton is more likely to exist as a hydronium ion (H3O+)

• Just as some acids are polyprotic, some bases can supply more than one mole of hydroxide ion per mole of base • Ba(OH)(s) Ba2

+(aq) + 2OH-

(aq) ______________________

GATHER AND PROCESS INFORMATION FROM SECONDARY SOURCES TO TRACE DEVELOPMENTS IN UNDERSTANDING AND DESCRIBING ACID/BASE REACTIONS

• Arrhenius suggested that neutralisation reactions produce water when hydrogen ions and hydroxide ions react o Dissociation of acid: H2SO4(aq) 2H+

(aq) + SO42-

(aq)

o Dissociation of base: Ba(OH)(aq) Ba2+(aq) + 2OH-

(aq)

o Full Equation: H2SO4(aq) + Ba(OH)(aq) BaSO4(aq) + H2O(l)

o Full Ionic Equation: 2H+(aq) + SO4

2-(aq) + Ba2+

(aq) + 2OH-(aq) Ba2+

(aq) + SO42-

(aq) + H2O(l)

o Net reaction is: H+(aq) + OH-

(aq) H2O(l)

• Bronsted-Lowry theory: acid-base reaction involves a transfer of a proton from one species to another OUTLINE THE BRÖNSTED-LOWRY THEORY OF ACIDS AND BASES

• Developed independently by Johannes Bronsted and Thomas Lowry • Bronsted-Lowry theory: acid-base reaction involves a transfer of a proton from an acid to a base

o An acid is a proton donor; a base is a proton acceptor HCl(g) + H2O(l) ↔ H3O+

(aq) + Cl-(aq)

• HCl is donating a proton and is therefore acting as an acid. Cl- is the conjugate base of HCl. H2O is accepting proton and therefore acting as a base. H3O+ is the conjugate acid of H2O

NH3(aq) + H2O(l) ↔ NH4+

(aq) + OH-(aq)

• NH3 accepts a proton, and is acting as a base. H2O donates a proton, and is acting as an acid

• Bronsted-Lowry theory addresses the limitations of Arrhenius’ theory: o Explains basic behaviour of many substances which do not contain an OH group

NH3(aq) + H2O(l) ↔ NH4(aq) + OH-(aq)

o Extends acid-base reactions and acidic/basic properties to solvents other than water, and reactions without a solvent NH3(g) + HCl(g) NH4Cl(s)

o Explains amphoterism; why some substances can act as both acids and bases HCO3

-(aq) + H2O(l) ↔ H2CO3(aq) + OH-

(aq)

HCO3-(aq) + H2O(l) ↔ CO3

2-(aq) + H3O+

(aq)

IDENTIFY AMPHIPROTIC SUBSTANCES AND CONSTRUCT EQUATIONS TO DESCRIBE THEIR BEHAVIOUR IN ACIDIC AND BASIC SOLUTIONS

• Amphiprotic substances can accept or donate protons, thus acting both as an acid and a base according to Bronsted-Lowry theory

o An amphiprotic substance must contain a hydrogen atom, as it can donate a proton • Amphoteric substances can act as both an acid and a base • Since amphiprotic substances can react both as an acid and as a base, all amphiprotic substances are

amphoteric, but not all amphoteric substances are amphiprotic o Some metal oxides are amphoteric but not amphiprotic; they don’t donate a proton, but can participate

in neutralisation reactions, e.g. zinc oxide As a base: ZnO(s) + 2H+

(aq) → Zn2+(aq) + H2O(l)

As an acid: ZnO(s) + H2O(l) + 2OH-(aq) → [Zn(OH)4]2-

(aq)

Identifying amphiprotic substances: • E.g. water (H2O)

o Acid: NH3(aq) + H2O(l) ↔ NH4+

(aq) + OH-(aq)

H2O H+ + OH-

o Base: HCl(g) + H2O(l) ↔ H3O+(aq) + Cl-(aq)

H2O + H+ H3O+

o One acid, one base: H2O(l) + H2O(l) ↔ H3O+(aq) + OH-

(aq)

• E.g. hydrogencarbonate ion (HCO3-)

o Acid: HCO3-(aq) + H2O(l) ↔ CO3

2-(aq) + H3O+

(aq)

o Base: HCO3-(aq) + H2O(l) ↔ H2CO3(aq) + OH-

(aq)

IDENTIFY CONJUGATE ACID/BASE PAIRS Whenever an acid and a base react, they form their conjugates: HCl + H2O Cl- + H3O+ acid1 base2 conjugate base1 conjugate acid2

See Figure… *Acid; base

• Hydrochloric acid; chloride ion • Sulfuric acid; hydrogensulfate ion • Hydronium ion; water • Hydrogensulfate ion; sulphate ion • Acetic acid; acetate (ethanoate) ion • Carbonic acid; hydrogencarbonate ion • Ammonium ion; ammonia • Hydrogencarbonate ion; carbonate ion • Water; hydroxide ion • Hydroxide ion; oxide ion

DESCRIBE THE RELATIONSHIP BETWEEN AN ACID AND ITS CONJUGATE BASE AND A BASE AND ITS CONJUGATE ACID

• When an acid donates a proton, it forms its conjugate base • When a base accepts a proton, it forms its conjugate acid • E.g. CH3COOH (acid) and CH3COO- (base)

o CH3COOH(aq) + OH-(aq) ↔ CH3 COO-

(aq) + H2O(l)

o Acid + Base ↔ Conjugate base + Conjugate acid

o CH3COO-(aq) + H3O+

(aq) ↔ R CH3COOH(aq) + H2O(l)

o Base + Acid ↔ Conjugate acid + Conjugate base

• A stronger acid will have a weaker conjugate base o HCl is a strong acid; Cl- ion is the weakest base

• A weaker acid will have a stronger conjugate base IDENTIFY A RANGE OF SALTS WHICH FORM ACIDIC, BASIC OR NEUTRAL SOLUTIONS AND EXPLAIN THEIR ACIDIC, NEUTRAL OR BASIC NATURE Aqueous solutions of salt can be acidic, basic or neutral, depending on the particular ions in the salt

• Neutral: ion do not react with water o Anions from strong acids: Cl-, NO3

-, Br-, I- i.e. conjugate bases of strong acids; e.g. Cl- anions have no tendency to react with water to form

HCl or hydroxide ions o Group I/II cations from strong bases: Li+, Mg2+, Na+, Ca2+, K+, Ba2+

• Basic: ion react with water to form hydroxide ions o Anions from weak acids: F-, S2-, SO4

2-, ClO-, CH3COO-, CO32-, HCO3

-, PO43-, HPO4

2- i.e. conjugate bases of weak acids; e.g. CH3 COO-

(aq) + H2O(l) ↔ CH3COOH(aq) + OH-(aq): only

occurs to a small extent o Some anions from polyprotic acids: HCO3

-, HPO42-

• Acidic: ions react with water to form hydronium ions o Cations from weak bases: NH4

+ NH4

+(aq) + H2O(l) ↔ NH3(aq) + H3O+

(aq) o Some anions from polyprotic acids: HSO4

-, H2PO4-

i.e. anions with hydrogen atoms that can be transferred to water molecules; e.g. HSO4

-(aq + H2O(l)) ↔ H3O+

(aq) + SO42-

(aq) o Cations from aquated metal ions: Al3+, Fe3+

Small, highly charged aquated aluminium ions can give up protons to produce hydronium ions in solution

Neutral Basic Acidic Anions from strong acids Anions from weak acids Cations from weak bases;

*aquated metal ions Cations from strong bases Some anions from

polyprotic acids Some anions from polyprotic acids

• In a salt formed from a strong acid and a weak base, the cation is a weak acid, so pH < 7 • In a salt formed from a weak acid and a strong base, the anion is a weak base, so pH > 7 • In a salt formed from a strong acid and a strong base, neither conjugates react with water, so pH close to 7

o Sodium chloride solution is neutral, because Na+ from the strong base NaOH and Cl- ions from the strong acid HCl, do not react with water to form hydronium or hydroxide ions

• In a salt formed from a weak acid and a weak base, anion and cation react with water to approximately equal extents and therefore cancel each other out, so pH close to 7

o e.g. ammonium acetate NH4CH3COO.

NH4+ + H2O NH3 + H3O+

CH3COO- + H2O CH3COOH + OH-

The resulting reaction, H3O+ + OH- 2H2O, results in a neutral solution.

CHOOSE EQUIPMENT AND PERFORM A FIRST-HAND INVESTIGATION TO IDENTIFY THE PH OF A RANGE OF SALT SOLUTIONS IDENTIFY NEUTRALISATION AS A PROTON TRANSFER REACTION WHICH IS EXOTHERMIC

• Acid-base reactions are known as neutralisation reactions which involve transfer of protons from one species to another

• Acid + metal hydroxide salt + water • Acid + metal oxide salt + water • Acid + carbonate/hydrogencarbonate salt + water + carbon dioxide gas • E.g. HNO3(aq) + NaOH(aq) NaNO3(aq) + H2O(l) + heat

o Sodium and nitrate ions are spectator ions Net ionic equation is: H+

(aq) + OH-(aq) H2O(l)

o This neutralisation reaction is exothermic because it results in the formation of a new covalent bond between hydrogen and oxygen as water molecules are formed

o When the heat of neutralisation is measured for a range of strong acids and strong bases, the amount of

heat released is always about 57 kJ per mole of water formed. This is the heat change for the following

reaction: H+(aq) + OH-

(aq) H2O(l) ∆𝐻 = −57 𝑘𝐽 𝑚𝑜𝑙−1

• E.g. reaction of hydrochloric acid and ammonia solution can be written as: NH3(aq) + H+

(aq) NH4+

(aq)

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

• Many acids/bases are corrosive, and chemical spills of these substances must be cleaned quickly with neutralisation reactions

o E.g. if a small amount of acid has been spilt, it should be covered with excess sodium

hydrogencarbonate

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

Net reaction: H+(aq) + HCO3

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

o *This exothermic reaction will generate a lot of heat, so spills of acids/bases on skin should be washed off with water instead of being neutralised

• Spills of bases can be neutralised with weak acids such as benzoic acid, boric acid or acetic acid o Sodium hydrogencarbonate is a versatile neutralising reagent for both acid and base spills as it contains

the amphiprotic hydrogencarbonate ion

OH-(aq) + HCO3

-(aq) CO3

2- + H2O

• Properties of a good neutralising agent: o Cheap o Stable: does not react with air o Solid o Safe to handle/store o If excess is used, excess presents less danger o Quick-acting neutraliser o Non-toxic

DESCRIBE THE CORRECT TECHNIQUE FOR CONDUCTING TITRATIONS AND PREPARATION OF STANDARD SOLUTIONS A primary standard is a substance of such high purity and stability that it can be used directly for the preparation of standard solutions A standard solution is a solution of accurately known composition and concentration. It can be prepared by weighing out the desired mass, dissolving it in water, and making volume up to an accurately known value For a chemical to be suitable to prepare as a standard solution, it must:

1) be a water soluble solid 2) have high purity - usually Analytical Reagent (A.R.) grade 3) have an accurately known formula 4) be stable in air, i.e. it does not lose or gain water or react with oxygen or carbon dioxide in air. 5) Have a reasonably high relative molecular mass,

• If molecular mass is higher, then any required number of moles of the substance will have a greater mass, so that any weighing errors have a smaller percentage effect on the total mass.

• Basic primary standard: anhydrous sodium carbonate [Na2CO3]

o Note*: Na2CO3 for primary standard, NaHCO3 for neutralisation • Acidic primary standard: hydrated oxalic acid [H2C2O4 . 2 H2O] • If a particular solution must be used as the titrant, but it is not composed of a primary standard, then its

concentration can be found (standardised) by titration against a primary standard. o Sodium hydroxide NaOH is difficult to obtain in very pure form, and hygroscopic (absorbs moisture from

the air) NaOH also reacts with CO2 in the air

o Hydrochloric acid HCl is also not a primary standard; its concentration varies slightly among batches; concentrated HCl fumes and releases HCl fumes

o Concentrated sulfuric acid absorbs water from the air o Hydrated sodium carbonate loses water to the air as it is being weighed o Possible sequence for preparing standard sodium hydroxide solution

Use a standard Na2CO3 solution to standardise a HCl solution Use standard HCl solution to standardise a NaOH solution

How to prepare a standard solution:

1) Weigh on analytical balance exact amount of solid on watch glass 2) Transfer to beaker, dissolve in distilled water 3) Using a funnel, transfer to volumetric flask 4) Wash out beaker with distilled water and tip into flask

5) Using distilled water, make up the volume in the flask until the bottom of the meniscus is level with etched mark on the flask.

6) Stopper the flask and shake to ensure thorough mixing 7) Label solution, giving name and concentration

Titration Equipment

• Volumetric flask: holds an accurately known volume of solution (usually 250 mL) indicated by a line etched into the neck of the flask

o Used for preparation of standard solutions • Pipette: used to accurately deliver a specific volume of solution (usually 25 mL) into the conical flask

o Do not shake out the last drop; the pipette is calibrated to account for this • Burette: used to accurately deliver variable volumes of solution into the conical flask

o Reading is usually estimated to ± 0.05 mL Preparation of Glassware:

• Rinse the pipette with distilled water, and then rinse with the solution that will be placed in there. o If pipette is rinsed with only distilled water, then the water droplets remaining on inside of glassware

will dilute the solution. The titration procedure allows calculation of the concentration of solution transferred from pipette to conical flask, so if the pipette solution is more diluted than the original sample, then the procedure will not accurately determine concentration of the sample.

• Rinse the burette with distilled water, and then rinse with the solution that will be placed in there. o If burette is rinsed with only distilled water then water droplets remaining on inside of glassware will

dilute the solution, and the number of moles of solution added can no longer be accurately determined

• Wash the conical flask with distilled water only o Any water remaining in the flask will not change the number of moles of solution

Titration Procedure:

1) Wash burette with distilled water, and then with some of the titrant solution 2) Set up burette with retort stand 3) Carefully pour titrant solution into burette, with a funnel, until bottom of meniscus is level with the zero mark 4) Wash pipette with distilled water, and then with some of the analyte solution to be titrated 5) Fill a 25mL pipette up to etched mark with analyte solution

a. Note that solution should be poured from original container (e.g. volumetric flask) into a clean, dry beaker, and drawn into pipette from the beaker. This is so that the whole stock of analyte solution is not contaminated by any foreign material in/on the pipette.

6) Wash a conical flask with distilled water 7) Empty the analyte solution from the pipette into a conical flask, and add 3 drops of indicator 8) Position conical flask under burette 9) Using burette, slowly add titrant solution into the conical flask one drop at a time, swirling the flask all the

while, until the required colour change occurs 10) Record amount of titrant solution used up in burette 11) Repeat all previous steps 3 more times, for a total of 1 rough titration, and 3 precise titrations.

Definitions:

• Titration: Procedure used to experimentally find the concentration of a solution. • Indicator: Substance which signals the equivalence point of a titration by changing colours

• Equivalence point: Point when reaction is complete. • End point: Point when indicator changes colour.

• Titrant: Standard solution of known concentration and composition. • Titre: Volume of solution delivered from the burette. • Analyte: any chemical to be analysed • Aliquot: Fixed volume of solution delivered by pipette.

The endpoint is the point in a titration when the indicator changes colour. An indicator is selected that changes colour at the pH of the salt solution formed at the point of neutralisation. This is known as the equivalence point. Titration Curves:

o http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.html • Note that point of inflexion is at pH of the salt solution produced • In a titration curve, independent variable is volume of titrant, and dependent variable is pH of solution

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 QUALITATIVELY DESCRIBE THE EFFECT OF BUFFERS WITH REFERENCE TO A SPECIFIC EXAMPLE IN A NATURAL SYSTEM Effect of Buffers

• Buffer solutions resist changes to pH when small amounts of acids or bases are added • Buffers contain approximately equal amounts of a weak acid and its conjugate base • E.g. ethanoic acid and ethanoate ions

o If a base in the form of hydroxide ions is added, following reaction occurs CH3COOH(aq) + OH-

(aq) H2O(l) + CH3COO-(aq)

o If an acid in the form of hydronium ions is added, following reaction occurs CH3COO-

(aq) + H3O+(aq) H2O(l) + CH3COOH(aq)

o Most of base or acid added to buffer solution is neutralised, and pH remains relatively constant Buffer in a Natural System

• Main buffer system used to control blood pH is carbonic acid/hydrogencarbonate ion buffer system H2CO3(aq) + H2O(l) ↔HCO3

-(aq) + H3O+

(aq) o If a base is added, it is neutralised by carbonic acid o If acid is added, it is neutralised by hydrogen carbonate ions

• Maintains pH of blood within narrow range around 7.4 o Metabolic processes need constant pH because they depend on enzymes, and enzymes only work

efficiently within a narrow pH range. However, many waste products in the body are acids or bases

• Also important in maintaining fairly constant pH of 8.5 in ocean water

PERFORM A FIRST-HAND INVESTIGATION TO DETERMINE THE CONCENTRATION OF A DOMESTIC ACIDIC SUBSTANCE USING COMPUTER-BASED TECHNOLOGIES

5. Esterification is a naturally occurring process which can be performed in the laboratory DESCRIBE THE DIFFERENCES BETWEEN THE ALKANOL AND ALKANOIC ACID FUNCTIONAL GROUPS IN CARBON COMPOUNDS Alkanols: CnH2n+1OH, ROH

• The functional group in alkanols is the hydroxyl group –OH • E.g. methanol [CH3OH], ethanol [CH3CH2OH], 1-propanol [CH3(CH2)2OH]

Alkanoic acid/carboxylic acid: RCOOH • The functional acid in an alkanoic acid is the carboxyl group –COOH

o Consists of a carbon atom, which has a double bond with an oxygen atom and a single bond with a hydroxyl group

• Alkanoic acids are weak acids: RCOOH(aq) ↔ H+(aq) + RCOO-

(aq) o Strength of alkanoic acids decrease with length of hydrocarbon chain

• E.g. methanoic acid [HCOOH], ethanoic acid [CH3COOH], propanoic acid [CH3CH2COOH] EXPLAIN THE DIFFERENCE IN MELTING POINT AND BOILING POINT CAUSED BY STRAIGHT-CHAINED ALKANOIC ACID AND STRAIGHT-CHAINED PRIMARY ALKANOL STRUCTURES

• Hydrogen bonding can occur between neighbouring alkanol molecules as they contain –OH groups

• Alkanoic acids can also form hydrogen bonds between molecules • Alkanoic acids contain a –C=O group as well as an –OH group, and so two

hydrogen bonds can be formed between each alkanoic acid molecule. o As a result, alkanoic acids generally possess higher melting/boiling

points than alkanols o Both alkanoic acids and alkanols generally have higher

melting/boiling points than hydrocarbons of similar molecular mass • As a result of hydrogen bonding, both alkanoic acids and alkanols can

readily dissolve in water, although their solubility decreases with increasing length of the hydrocarbon chain IDENTIFY ESTERIFICATION AS THE REACTION BETWEEN AN ACID AND AN ALKANOL AND DESCRIBE, USING EQUATIONS, EXAMPLES OF ESTERIFICATION a n d 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

• Esterification is the acid-catalysed condensation reaction between an alkanoic acid and an alkanol to form esters, with water as the by-product

• An ester consists of an acid in which at least one –OH hydroxyl group is replaced by an –O-alkyl (alkoxy) group o If an oxygen-18 isotope, O, is used in the alkanol only, it is found in the ester, but not in the water

product. Use of this tracer shows that the O in water comes from the acid.

o Therefore, a hydrogen atom from the alkanol molecule joins with hydroxyl group from the acid to form water

o The reaction is reversible and comparable quantities of alkanol, acid, ester and water are present at

equilibrium

• RCOOH + R’OH 𝐻+�� RCOOR’ + H2O

• Alkanoic acid + Alkanol 𝐻+�� Ester + Water

_______________

• CH3COOH(l) + CH3OH(l) 𝐻2𝑆𝑂4�⎯⎯� CH3COOCH3(l) + H2O(l)

o Common: Acetic acid + methyl alcohol 𝐻2𝑆𝑂4�⎯⎯� methyl acetate + water

o IUPAC: Ethanoic acid + methanol 𝐻2𝑆𝑂4�⎯⎯� methyl ethanoate + water

CH3 is the methyl group; CH3COO- is the ethanoate group

• HCOOH(l) + CH3CH2OH(l) 𝐻2𝑆𝑂4�⎯⎯� HCOOCH2CH3(l) + H2O(l)

o Common: Formic acid + ethyl alcohol 𝐻2𝑆𝑂4�⎯⎯� ethyl formate + water

o IUPAC: Methanoic acid + ethanol 𝐻2𝑆𝑂4�⎯⎯� ethyl methanoate + water

• Ethyl ethanoate (diagram)

o Note*: all liquid states, and don’t forget the H2SO4 catalyst!

***check if numbering is used. Numbering always takes end of alkane closest to CO group as 1., e.g. 1-butyl ethanoate DESCRIBE THE PURPOSE OF USING ACID IN ESTERIFICATION FOR CATALYSIS

• A few drops of concentrated acid needs to be added to a mixture of alkanol and alkanoic acid to catalyse the reaction by lowering activation energy

• Concentrated sulfuric acid is a dehydrating agent; it has a strong affinity for water

• It will shift the equilibrium position to the right by absorbing water, as predicted by Le Chatelier’s principle

o Alcohol + acid ester + water • Using acid in esterification increases the yield of ester as well as speeding up the

reaction EXPLAIN THE NEED FOR REFLUXING DURING ESTERIFICATION

• Esterification requires heat for the reaction to reach equilibrium within an hour, rather than after many days.

o Reactants and ester product may be volatile, and when heated could escape the reaction vessel • Refluxing prevents volatile reactants and products from escaping while being heated • A condenser is placed on top of the reaction vessel so that volatile components pass into the condenser. The

condenser is water or air-cooled, causing the gases to condense and fall back into the reaction mixture • Refluxing also increases safety as volatile components such as alcohols are flammable

______________________ • Indirect heat source such as a water bath is used to prevent flammable reactants or products catching on fire • Boiling chips used to promote even boiling

______________________ • Esters tend to be liquids at room temperature • Esters have boiling points lower than those of alkanoic acids with similar molecular masses because

intermolecular forces in esters are dispersion and dipole-dipole forces rather than hydrogen bonding. In industrial esterification, the ester produced may be removed by fractional distillation to shift the equilibrium right.

OUTLINE SOME EXAMPLES OF THE OCCURRENCE, PRODUCTION AND USES OF ESTERS Occurrence:

• Natural flavours of foods are usually a complex mixture of several esters • Natural perfumes produced by many flowers are also esters • Solid animal fats, and animal/vegetable oils, are natural esters

o These triglycerides are formed from glycerol (1,2,3-propanetriol) and saturated/unsaturated fatty acids o Fats contain mainly saturated fatty acids; oils contain mainly unsaturated fatty acids

Fats include beef tallow, butter, and lard Oils include coconut, corn, olive, and

peanut oil • Waxes are esters formed from long-chain aliphatic

alcohols and long-chain aliphatic alkanoic acids o Natural waxes consist of a mixture of various

esters • Phosphoesters make up backbone of DNA molecule

o Linkage between adjacent deoxyribose sugars is a phosphodiester bond (from phosphoric acid)

Fischer Esterification:

• The carboxylic acid is refluxed in an alcohol. • Using the alcohol as a solvent (i.e. in large excess) shifts

the equilibrium right. This is why all reactants and products are in liquid state

• Sulfuric acid is used as a catalyst • Since esters cannot form hydrogen bonds, they generally have lower boiling points than the reactants. The ester

is removed from the reaction vessel by fractional distillation as it is formed, shifting the equilibrium right Production and Uses:

• Volatile esters have pleasant odours; they are often used to mimic natural odours • Vegetable oils are partially hydrogenated to convert liquid oils into solid margarine • Butter is produced from milk fats • Aspirin: acetylsalicylic acid or 2-ethanoylhydroxybenzoic acid

o Produced by reacting salicyclic acid with acetic anhydride o Used as a mild analgesic or pain-reliever useful in relief of headaches and other pains

o Also prescribed to prevent blood from clotting in patients with cardiovascular diseases • Dialkyl phthalates (e.g. dimethyl phthalate) [Industrial use of ester]

o Dialkyl esters of phthalic acid o Used as plasticisers: When added to plastics such as PVC, phthalates allow polymer chains to slide

against one another, increasing flexibility and durability • Ethyl ethanoate

o Synthesised industrially via the classic Fischer esterification of ethanol and ethanoic acid o Used as nail varnish remover because of its low toxicity and agreeable odour o As a solvent in perfumes, it has an agreeable odour, and evaporates quickly leaving the fragrant

chemicals on the skin • Butyl acetate

o Found naturally in fruits such as apples o Used as synthetic fruit flavouring in candy, ice cream and cheeses

• Benzyl acetate o Produced from benzyl alcohol and acetic acid o Found in oils of jasmine and other flowers o Used in perfumery and cosmetics because it has a sweet aroma like that of the jasmine flower

PROCESS INFORMATION FROM SECONDARY SOURCES TO IDENTIFY AND DESCRIBE THE USES OF ESTERS AS FLAVOURS AND PERFUMES IN PROCESSED FOODS AND COSMETICS IDENTIFY DATA, PLAN, SELECT EQUIPMENT AND PERFORM A FIRST-HAND INVESTIGATION TO PREPARE AN ESTER USING REFLUX