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
Classification
Acid
Neutral
Base
Properties
Sour-tasting
Corrosive to active
metals
Dissolves limestone
Unreactive towards
fats
Not much taste
Unreactive towards
most metals
Unreactive towards
limestone
Unreactive towards
fats
Bitter-tasting (soapy)
Unreactive towards
most metals
Unreactive towards
limestone
Reacts with fats
Colour Change (litmus)
Colour changes from
blue to red
Substances
Vinegar
Soda water
Soft drinks
Lemon juice
Battery acid
No colour change.
Stays blue/purple
Water
Glucose (sugar)
Sodium chloride (table
salt)
Colour changes from
red to blue
Baking soda solution
Lime water
Ammonia solutions
Bleach
Detergent
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;
Identify data and choose resources to gather information about the colour changes of a range of
indicators
Acid-base indicators are substances that change colour depending on the pH
The working/effective range of an indicator is the pH range in which a colour change is observed.
 Litmus paper(working range: 5-8) can be used to provide a general classification for a
substance to identify whether it is an acid or base. It is best used in solutions that are
moderately/strongly basic or acidic.
 Methyl orange (working range 3.1-4.4)
 Bromothymol blue (working range 6.0-7.6)
 Phenolphthalein (working range 8.3-10.0)
pH
Methyl orange
Litmus
14 base
13
12
11
10
9
8
7 neutral
6
5
4
3
2
1
0 acid
Yellow
Yellow
Yellow
Yellow
Yellow
Yellow
Yellow
Yellow
Yellow
Orange
Orange-red
Orange-red
Red
Red
Red
Blue
Blue
Blue
Blue
Blue
Blue
Blue-purple
Blue-purple
Blue-purple
Bluish red
Red
Red
Red
Red
Red
Bromothymol
Blue
Blue
Blue
Blue
Blue
Blue
Blue
Green-blue
Green
Yellow-green
Yellow
Yellow
Yellow
Yellow
Yellow
Yellow
Phenolphthalein
Crimson (red)
Crimson (red)
Crimson (red)
Crimson (red)
Deep pink
Pink
Light pink
Colourless
Colourless
Colourless
Colourless
Colourless
Colourless
Colourless
Colourless
Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity
Indicators can be used for various purposes in everyday life.
 Testing the acidity of water
- In swimming pools, the acidity levels must be monitored in order to effectively kill
microbes. Samples of pool water can be tested using a pool test kit which consists of
phenol red indicator. If the pH is below 6.8, the indicator turns yellow; if the indicator is
above 8.4, the indicator turns red purple. If the indicator is pink or orange, the acidity of
the pool water is appropriate.
 Testing the acidity of soil
- Soils consisting of different pH are required for different plants/crops. Soil pH can be
measured using electronic instruments or universal indicator. The soil is first mixed with
water in a tube and then indicator is added. Sometimes the colour of the soil can hide
the indicator colour change; to counter this, a neutral white powder (e.g. barium
sulphate) can be added to the top layer of the damp soil before adding indicator. This
then allows the colour of the indicator to be clearly visible.
2. While we usually think of the air around us as neutral, the atmosphere naturally ontains
acidic oxides of carbon, nitrogen and sulphur. The concentrations of these acidic oxides
have been increasing since the Industrial Revolution
Identify oxides of non-metals which act as acids and describe the conditions under which they act
as acids:
Acidic oxides
Acidic oxides:
 React with water to form acids.
 React with bases to form salts.
Non-metal oxides behave as acids. Some non-metal oxides that act as acids are: SO2, NO2 and P2O5
(phosphorus pentoxide).
 CO2 (g) + H2O (l)
 SO2 (g) + H2O (l)
 2NO2 (g) + H2O (l)
H2CO3 (aq)
H2SO3 (aq)
HNO3 (aq) + HNO2 (aq) (nitrous acid)
HNO2 (aq) + O2 (g)
HNO3 (aq)
 P2O5 (g) + H2O (l)
2H3PO4 (aq) (phosphoric acid)
*Note: Exceptions include N2O, CO and NO; these are neutral oxides
Basic oxides
Basic oxides:
 React with water to form bases
 React with acids to form salts
Metal oxides behave as BASES. Some metal oxides that act as bases are: K2O, Na2O and MgO. In
solution, they tend to form basic hydroxides
 K2O (s) + H2O (l)
2KOH (aq) (potassium hydroxide)
 Na2O (s) + H2O (l)
2NaOH (aq) (sodium hydroxide)
 MgO (s) + H2O (l)
Mg(OH)2 (aq) (magnesium hydroxide)
Amphoteric oxides
Amphoteric oxides are oxides that can act as both acids and bases (e.g. Al2O3)
Their behaviour depends on the reaction they are put in.
The only elements that combine to form amphoteric oxides are beryllium, aluminium, zinc, tin and
lead.
Analyse the position of these non-metals in the Period Table and outline the relationship between
position of elements in the Periodic Table and acidity/basicity of oxides
Non-metallic oxides are generally acidic. Acidity generally increases across a period and increases
down a group.
Metallic oxides are generally basic (sources of O2+). The basicity of metal oxides generally decreases
across a period and increases down a group
Transition metal oxides display variable acid-base behaviour and are often amphoteric. Oxides of
metalloids are often amphoteric
Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen
Sulfur dioxide:
Natural sources include: volcanoes, bushfires and bacterial decomposition
Industrial sources: burning of coal in electricity production, smelting of metals
Oxides of nitrogen:
Natural sources include: lightning storms
Industrial sources: internal combustion within cars and other heavy machinery
Describe, using equations, examples of chemical reactions which release sulphur dioxide and
chemical reactions which release oxides of nitrogen
Sulfur Dioxide:
 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:
 When nitrogen and oxygen react at high temperature (e.g. in engines of motor vehicles),
nitric oxide is formed:
N2 (g) + O2 (g)
2NO (g)
 Nitric acid is neutral; however it reacts with oxygen in the air to form nitrogen dioxide, which
is acidic.
2NO (g) + O2 (g)
NO2 (g)
SO2 (sulfur dioxide)
NO (nitric oxide)
NO2 (nitrogen dioxide)
Properties
Colourless
Colourless
Reddish-brown
Pungent colour
No smell
Choking odour
Soluble in water**
Insoluble
Soluble in water **
Uses
Food preservative
Nitric acid
Bleaching
Fertilisers
Fumigant
Explosives
Issues:
Respiratory irritant
Synthesising NO2
Respiratory irritant (3(1ppm) – asthmatics
5ppm). At higher
are particularly
concentrations it can
susceptible
cause tissue damage
Assess the evidence which indicates increases in atmospheric concentration of oxides of sulphur
and nitrogen
Since the Industrial Revolution of the 19th Century, vast quantities of nitrogen and sulphur oxides
have been released in the air. The main sources of these oxides are from the smelting of metals,
burning of coal/oil and internal combustion within engines.
Oxides of sulphur and nitrogen dioxide were difficult to measure due to its relatively low
concentrations in the atmosphere. Chemical instruments able to measure such low concentrations
were only commercially available since the 1970s. Most of the evidence for the increases in sulphur
oxides was in heavily industrial areas, where the increased effects of acid rain (which is formed when
the oxides react with rainwater) were observed.
Explain the formation and effects of acid rain
‘Acid rain’ generally describes rain with moderately low pH of around 4-5. It occurs when the
atmosphere is polluted with acidic oxides (e.g. sulfur dioxide and nitrogen dioxide) which have high
solubility in water. The main causes of acid precipitation are related to industrial emissions.
The following reactions account for the formation of acid rain:
Nitric oxides
 2NO2 (g) + H2O (l)
 HNO2 (aq) + O2 (g)
Sulfur oxides
 SO2 (g) + H2O (l)
 2H2SO3 (aq) + O2 (g)
HNO3 (aq) + HNO2 (aq) (nitric, nitrous acid)
HNO3 (aq)
H2SO3 (aq) (sulfurous acid)
2H2SO4 (aq)
(OR)
 2SO2 (g) + O2 (g)
 SO3 (g) + H2O (l)
2SO3 (g)
H2SO4 (aq)
Effects of acid rain
 Has a corrosive effect on limestone and marble. For example, the calcium carbonate in the
marble is attacked by sulfuric acid in acid rain to form insoluble calcium sulfate:
 Can attack metallic structures of iron or steel. The iron is oxidised by the acid and becomes
chemically weathered:
 Acidifies soils, which can inhibit the growth of certain plants.
 Aquatic organisms are also affected by acid rain as it changes the pH of many lakes. As the
lake water becomes more acidic, the presence of hydronium ions interferes with the carbon
dioxide/carbonate equilibrium. Plants that require CO2 and crustaceans (with shells that
require carbonate) consequently suffer.
Analyse information from secondary sources to summarise the industrial origins of sulphur dioxide
and oxides of nitrogen and evaluate reasons for concern about their release into the environment
Oxides of sulphur and nitrogen tend to cause acid rain (with the exception of NO, which is a neutral
oxide, but it reacts with oxygen to form NO2). Also, all oxides of sulphur and nitrogen are respiratory
irritants. Both of these effects (acid rain/respiratory irritant) can damage ecosystems if the
concentrations of the oxides are not moderated
 The acid rain caused by dissolution of SO2 and SO3 in rainwater (H2SO4) lowers the pH of
lakes and soils, interfering with the balance of nutrients (e.g. CO32+). The acidity can also
affect the soil and inhibit growth of certain plants.
 Sulfur and nitrogen oxides produced from car exhausts pose a health hazard on population
areas due to irritation of the lungs, particularly asthmatics.
 NO2 forms photochemical smog. Sunlight reacts with nitrogen dioxide, hydrocarbons and
oxygen to form ozone. Ozone is harmful in the troposphere at concentrations as low as
0.1ppm
-
NO2
O· + O2
NO + O·
O3
Define Le Chatelier’s principle
The French chemist Henri Le Chatelier suggested that:
“When a system at equilibrium is disturbed, the equilibrium position will shift in the direction which
tends to minimise, or counteract, the effect of disturbance”
Identify factors which can affect the equilibrium in a reversible reaction
 Concentration – if the concentration of a solute reactant is increased, the equilibrium
position shifts to use up the added reactants by producing more products. (Note
concentration is actually referring to C=n/V)
 Pressure – if the pressure of an equilibrium position is increased, then the equilibrium
position shifts to reduce the pressure (the side with less moles of gas
 Volume – corresponds to pressure. If volume is increased, pressure is reduced; if volume
reduced then pressure is increased
 Temperature – if the temperature of an endothermic equilibrium system is increased, the
equilibrium position shifts to use up the heat by producing more products. If the
temperature of an exothermic equilibrium system is increased, the equilibrium position
shifts to use up heat by producing more reactants
* For more detail, refer to notes
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 in water:
1. CO2 (g)
CO2 (aq)
2. CO2 (aq) + H2O (l)
[Δv = -ve]
H2CO3 (aq)
3. H2CO3 (aq)
H+ (aq) + HCO3- (aq)
4. HCO3- (aq)
H+ (aq) + CO32- (aq)
Like all dissolution reactions, this dissolution reaction is exothermic
 Pressure – if we increase the gas pressure, equilibrium 1 shifts to the right, thus dissolving
more carbon dioxide. This causes equilibrium 2 to shift to the right (due to increase in
concentration of reactant) - causing an increase in the production of carbon acid. The
increase in concentration of carbonic acid shifts equilibrium 3 and 4 to the right. The
solution becomes more acidic (increased concentration of H+ ions). If we open a bottle of
soft drink, the pressure decreases and the opposite effect occurs. Bubbles of CO2 gas can be
seen coming out of the solution, and the soft drink becomes ‘flat’ as gaseous carbon dioxide
escapes
 Temperature – Since equilibrium 1 is exothermic, if we warm the soft drink, the equilibrium
shifts towards the left (since heat is thought of as the ‘product’ of equilibrium 1). This leads
to the release of CO2 from solution.
 pH – If an additional source of acidity was added, the additional H+ would drive equilibrium 3
and 4 towards the left, and eventually lead to the release of CO2 from solution. Conversely,
the addition of a base would decrease the concentration of H+ (due to the neutralisation
reaction it creates), thus shifting the equilibrium towards the right – dissolving more CO2 and
producing more carbonate ions.
Calculate volumes of gases given masses of some substances in reactions, and calculate masses of
substances given gaseous volumes, in reactions involving gases at standard temp/pressure
This is based on Avagadro’s deduction that equal volumes of gases, under the same conditions
(standard) would always have the same amount of particles.
3. Acids occur in many foods, drinks and even within our stomachs
Strong acids: Hydrochloric, sulfuric, nitric
Strong bases: Hydroxides of group I, II metals (NaOH)
Weak acids: Acetic, citric, carbonic, hydrogen fluoride, sulfurous, nitrous, phosphoric
Weak bases: Ammonia
Define acids as proton donors and describe the ionisation of acids in water
An acid can be defined as a substance, that in solution, produces hydrogen ions (H+), or protons.
However, in aqueous solution, free H+ does not actually exist; as the hydrogen ion actually combines
with a water molecule to form a hydronium ion (H3O+)
An ionisation reaction essentially involves an acid being ionised, for example:
HNO3 (l) + H2O (l)
H3O+(aq) + NO3-(aq)
Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic),
hydrochloric and sulfuric acid
Strong acids
Hydrochloric
acid
(HCl)
Sulfuric acid
(H2SO4)
Nitric acid
(HNO3)
Weak acids
Acetic acid
(CH3COOH)
Hydrochloric acid (HCl) is a strong acid made by passing hydrogen chloride gas into
water. The final concentrated solution is about 12 mol/L.
Sulfuric acid is another strong acid; in concentrated solution it is about 18 mol/L.
However, unlike hydrochloric acid, it is diprotic, meaning it can donate two protons
per molecule of acid. In contrast, a monoprotic acid is one that can only donate one
proton per molecule. Sulfuric acid ionises in two stages:
*note degree of ionisation decreases at second step
Nitric acid is another strong acid, and when concentrated, is about 16 mol/L. Its
ionisation equation:
Acidic acid, also known as ethanoic acid, is a weak acid – therefore it is only weakly
ionised. Even though it has 4 hydrogen atoms, it is monoprotic. Acetic acid is
formed naturally when microbes ferment sugars in fruit juices. It is also used as the
major component of vinegar.
Citric acid
Citric acid is a weak, tripotic acid (stronger than acetic) found in citrus fruits. Its
C3H5O(COOH)3 chemical name is 2-hydroxypropane-1,2,3-tricarboxylic acid.
It ionises in three steps (though the degree of ionisation decreases at each step)
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. It is numbered from 0-14.
 A pH of 7 is attributed to neutral substances
 A pH <7 refers to acidic substances (strong acids closer to 0)
 A pH >7 refers to basic substances (strong bases closer to 14)
Describe acids and their solutions with the appropriate use of the terms strong, weak,
concentrated and dilute
A concentrated solution is one in which the total concentration of solute species is high
A dilute solution is one which the total concentration of solute species is low
A strong acid is one in which all the acid present in the solution has ionised to hydronium ions
A weak acid is one in which only some of the acid molecules in solution have ionised to form
hydronium ions
Identify pH as –log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+]
The pH of a solution is defined of the negative of the logarithm (base 10 ) of the hydronium ion
concentration:
pH = -log10[H3O+]
A change in pH of one means a 10-fold change in [H+]. For example a solution with a pH of 3 has 10
times the hydronium ion concentration than a solution with pH 4.
Furthermore, there is a similar scale for [OH-]; the relationships between pH and pOH are:
[H+][OH-] = 10-14
i.e. the product of the ion concentrations must equal to 1.0 x 10-14 (known as the water constant- Kw)
Note that water undergoes self-ionisation according to the equation: H2O
H+ + OH-.
+
-7
In pure water the [H ] = [OH ] = 10 .
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
Degree of ionisation is calculated using: concentration of hydronium ions/ concentration of initial
acid, as a percentage.
The strength of an acid is dependent on the ability of the acid to ionise in water. A stronger acid has
a higher degree of ionisation than a weaker acid. Strong acids are completely ionised (100%), whilst
weak acids have degree of ionisations less than 100%.
The following table represents the degree of ionisation when 0.10 mol/L solutions of hydrochloric,
citric, and acetic acid were prepared:
Acid (0.01M)
Hydrochloric
Citric
Acetic
Degree of ionisation (%)
100
27.5
4.2
From the above table, we can conclude that HCl is the strongest acid, followed by citric acid and
then acetic acid. Between citric and acetic acid, it is difficult to compare them as citric acid is triprotic
whilst acetic acid is monoprotic. However, we know that the higher value for citric acid is mainly
from the first step 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 is 100% ionised in solution (i.e. there is no unionised molecules of the acid present).
This means that its ionisation reaction goes to completion. For example:
A weak acid is not completely ionised in solution, and some of the unionised acid molecules still
exist. Thus, its ionisation reaction is written in equilibrium. For example:
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 flavour.
 Preservatives
- Ethanoic acid (in the form of vinegar) is used as a preservative in ‘pickling’.
- 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.
- Citric acid is a natural preservative, often added to jams and conserves.
 Flavourings
- 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.
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: Aqueous HCl; it is produced naturally by the lining of our stomachs. It aids
in the digestion of food.
 Citric acid: C6H8O7; occurs naturally in large quantities in citrus fruits.
 Ethanoic acid: CH3COOH; it is found naturally in vinegar, which is produced by the natural
oxidation of ethanol.
 Lactic acid: C3H6O3; it is formed in the body during strenuous exercise.
Natural bases
 Ammonia: NH3; it is present in the stale urine of animals. It is also formed through the
anaerobic decay of organic matter.
 Metallic oxides: E.g. iron(III) oxide, copper(II) oxide and titanium(IV) oxide. These insoluble
oxides are solid bases found in minerals.
 Calcium carbonate: CaCO3; it is found naturally as limestone.
Process information from secondary sources to calculate pH of strong acids given appropriate
hydrogen ion concentrations
An important thing to note is when calculating the pH of acids that are not monoprotic (i.e. are
diprotic or triprotic). Write out a balanced equation when calculating the pH resulting from a
reaction.
Strong, diprotic acid
Diprotic acids release 2 protons, e.g. H2SO4.
H2SO4
2H+ + SO42For every mole of acid, it releases 2 protons.
E.g. Calculate the pH of 0.1 M sulfuric acid:
 It has 0.1 mol/L; therefore its [H+] = 2 × 0.1 = 0.2 mol/L
 So, pH = -log10(0.2) = 0.7
Weak acid
The question should give the degree of ionisation as a percentage.
E.g. Calculate the pH of 0.1 mol/L ethanoic acid if only 1.3% ionises:
 [H+] = 1.3% of 0.1 = 0.0013
 pH = -log10(0.0013) = 2.9
* For more practice, refer to Faulder’s sheets or past HSC questions
4. Because of the prevalence and importance of acids, they have been used and studied for
hundreds of years. Over time, acid and base have been refined
Outline the historical development of ideas about acids including those of: Lavoisier, Davy, and
Arrhenius
Antoine Lavoisier was the first chemist to create theory of acids. He showed that many non-metal
compounds containing oxygen produced acids as they dissolved in water
 Lavoisier hypothesised that the presence of oxygen in the non-metal compounds gave the
compounds their acidic properties
 However, his theory of acids did not explain why oxides of metals were not acidic
Humphry Davy electrolysed samples of hydrochloric acid solution and showed that it produced
hydrogen and chlorine gas; no oxygen was formed.
 Thus, Davy proposed that the presence of hydrogen in acids gave them their acidic
properties
 However, his theory did not explain why many compounds of hydrogen were not acidic (e.g.
methane)
Svante Arrhenius noted that all acidic solutions were formed when acids ionised/dissociated into
ions (hydronium ion and an anion) as they dissolved in water
 Arrhenius recognised that some weaker acids did not ionise as completely as the strong
acids in water
 He also proposed that a base was a substance that produced hydroxide ions when dissolved
in water
 However, his theory also only applied to aqueous solution
 It also did not explain why ammonia (NH3), metal oxides (e.g. MgO) and carbonates are basic;
but these compounds do not contain hydroxide ions.
Outline the Bronsted-Lowry theory of acids and bases
In 1923, Johannes Bronsted (a Danish chemist) and Thomas Lowry (an English chemist)
independently proposed a new theory of acids and bases now known as the Bronsted-Lowry theory.
Bronsted and Lowry defined acids and bases as follows:
 Acids are proton donors
 Bases are proton acceptors
Thus, a substance cannot act as an acid (proton donor) without another acting as a base (proton
acceptor).
The Bronsted-Lowry theory also explains why acids dissolve in water to produce ions. A proton is
donated from the acid to the water molecule to produce a hydronium ion. The presence of the
hydronium ion gives the solution its acidic properties. Water is not only behaving as a solvent, but
also as a BL-base.
Similarly, a base can be dissolved in water to produce ions. In this instance, water acts as the BL-acid ,
donating a proton.
Describe the relationship between an acid and its conjugate base and a base and its conjugate acid
In an equilibrium reaction between a BL-acid and a BL-base, the BL-acid donates a proton to the BLbase, thus becoming proton deficient after the reaction. The proton deficient species is the
conjugate base, as it can accept protons to return back to its previous state. The original BL-base has
accepted a proton, and the new species is known as its conjugate acid (as it can donate a proton to
revert back).
Conjugates are the acids/bases on the right-hand side of the reaction. A conjugate acid/base pair
differs only by one proton (e.g. HF and F-)
Strong acids (e.g. hydrochloric acid and nitric acid) have very weak conjugate bases. The conjugate
acids of strong bases such as the hydroxide ion are also very weak. Weaker acids/bases have
relatively strong conjugate acids/bases. (Only incompetent/very weak acids and bases have actual
‘strong’ conjugates)
Identify conjugate acid/base pairs
Acid
HCl
HNO3
H3O+
HSO4HF
HNO2
CH3COOH
H2CO3
NH4+
HCO3
Conjugate base
ClNO3H2O
SO42FNO2CH3COOHCO3NH3
CO32-
Base
O2OHS2CO32NH3
HCO3HSCH3COOFSO42-
Conjugate acid
OHH2O
HSHCO3NH4+
H2CO3
H2S
CH3COOH
HF
HSO4-
Identify a range of salts which form acidic, basic or neutral solutions and explain their acidic,
neutral or basic nature
Basic salts refer to those that form basic solutions; Acidic salts refer to those that form acidic
solutions. Other salts are neutral.
The acidity/basicity of salts can be shown through hydrolysis reactions (see if the salt acts as a BLacid or BL- base as it reacts with water – neutral salts do not react with water). Hydrolysis refers to
the breaking of a bond with the addition of water
Weak Acid + Strong Base  Basic Salt + Water
Weak Base + Strong Acid  Acidic Salt + Water
Strong Acid + Strong Base  Neutral Salt + Water
Neutral salts include NaCL, KBr, NaNO3. Acidic and basic salts can be identified through the strength
of the acid/base reactants. An alternate method would be to analyse the origins of the cation and
anion to determine whether they are neutral, acidic of basic.
Identify amphiprotic substances and construct equations to describe their behaviour in acidic and
basic solutions
An amphiprotic species is one that can both donate or accept a proton.
H2O is amphiprotic
 H2O (l) + H2O (l)
H3O+ (aq) + OH- (aq)
HCO3‫ ־‬is amphiprotic:
- HCO3‫( ־‬aq) + H3O+ (aq)
-
HCO3‫( ־‬aq)
+ OH
‫־‬
(aq)
H2CO3 (aq) + H2O (l)
CO32‫( ־‬aq) + H2O (l)
Identify neutralisation as a proton transfer reaction which is exothermic
All neutralisation reactions are exothermic; they all liberate heat energy
The amount of heat liberated per mole in neutralisation reactions is almost the same no matter how
strong the acid/base: ΔH ≈ -56 kJ/mol. Although note that strong acid/base reactions produce
slightly more heat per mole
This similarity is understandable as the same reaction occurs in each neutralisation reaction:
-
H+ (aq) + OH ‫( ־‬aq)
H2O (l)
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.
It is important to note that neutralisation reactions are exothermic and thus release large quantities
of heat. Care must be taken, including the use of safety goggles and lab coats (in the laboratory).
 Strong or concentrated acids and bases must never be used to neutralise spills; if an excess
is used, the spill will become dangerous again.
When neutralising an acid or a base the following procedure is followed:
 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).
 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. The neutralised product (water, and a
soluble salt) is then absorbed using paper towels and disposed.
The most common substance used to neutralise spills in laboratories is powdered sodium
bicarbonate (NaHCO3); this is because the hydrogen carbonate ion (HCO3-) is an amphiprotic species,
and it is cheap and readily available substance.
In large chemical spills, inert sand or vermiculite is used to prevent the spillage area from further
contamination. The acidic sand/vermiculite is then placed in a chemical waste container and
removed for neutralisation off site (using sodium bicarbonate).
Describe the correct technique for conducting titrations and preparation of standard solutions
Volumetric analysis is a quantitative technique that involves the determination of unknown
concentration of a solution through a chemical reaction with a standardised solution. In acid-base
analysis, the reaction involved is a neutralisation reaction.
Apparatus
Volumetric flasks are flasks that hold an accurately known volume of solution. This volume is
indicated by a line etched into the neck of the flask (calibration mark). Volumetric flasks are used to
the preparation of solutions whose concentrations must be accurately known (primary or secondary
standards)
A pipette is used to accurately deliver a specific volume of solution such as 5.00mL, 10.00mL or
20.00mL. The pipette is used to deliver the appropriate volume of solution (an aliquot) into a conical
flask prior to titration.
Burettes are used to accurately deliver variable volumes of solution Burettes are usually graduated
from 0.0mL to 50.0mL and the reading on the scale is estimated to +-0.05mL. The difference
between the initial and final readings on the burette indicates the volume of solution delivered in a
titration.
*for diagrams, refer to book
Other terminology:
A known concentration volume of titrant reacts with a solution of titrand to determine
concentration. The volume of titrant reacted is called the titre (usually delivered by the burette).
The equivalence point is the point at which the neutralisation reaction is complete – i.e. when a
stoichiometric ratio of reactants have been added.
The end point of the titration is the point at which the indicator changes colour. It is best to have the
end point as close to the equivalence point as possible.
Preparing standard solutions
A primary standard is prepared using chemicals that are pure and satisfy the following criteria:





High level of purity
Accurately known composition
Free of moisture
Stable and unaffected by air during weighing
Readily soluble in pure (distilled) water
 High molar weight solid to reduce percentage error in weighing
 Reacts instantaneously and completely
Many pure substances do not meet such requirements: concentrated HCl fumes and loses HCl gas,
concentrated sulphuric acid absorbs water from the atmosphere, NaOH absorbs moisture from the
air (hygroscopic) and reacts with CO2.
Hydrated sodium carbonate is also unsuitable as it loses water as it Is being weighed; however,
anhydrous sodium carbonate is suitable. It is firstly dried in a drying oven then cooled in a desiccator.
The final product is free of water and able to use for a primary standard.
A secondary standard is a solution whose concentration has been determined by reacting with a
primary standard.
Method:
1. Clean all apparatus (rinse volumetric flask with distilled water several times)
2. Weigh the watch glass
3. Weight out 1.3g of anhydrous sodium carbonate and record its exact mass
4. Transfer sodium carbonate to the beaker.
5. Wash the sodium carbonate into beaker using a wash bottle
6. Stir until all the sodium carbonate is completely dissolved
7. Pour the solution into the volumetric flask using a filter funnel
8. Rinse the beaker, stirring rod and funnel several times with the wash bottle, ensuring all solution
goes into the volumetric flask
9. Add distilled water to the volumetric flask until the level is almost to the graduated mark on the
flask (e.g. 5mL away)
10. Using a dropper, add distilled water drop by drop until the bottom of the meniscus is level with
the graduation mark
11. Calculate the exact concentration of the solution you have made and write it on the flask
Titrating the primary standard against HCl acid
1. Clean all apparatus (rinse several times with distilled water – for burette and pipette rinse it twice
after with small quantities with the solution to be transferred in them; any water droplets will dilute
the solution. Conical flask can be left wet as any water remaining will not change the number of
moles of solution it holds)
2. Set up the burette in its stand. Check it is vertical and tap is off
3. Pour hydrochloric acid through a funnel into the burette and take a reading
4. Pipette 25mL of standard solution into a conical flask
5. Place a white tile under the conical flask to make it easier to observe colour changes of the
indicator
6. Add 3 drops of indicator to the conical flask
7. Adjust burette until the tip is just inside the top of the conical flask
8. Open the tap and slowly add hydrochloric acid to the conical flask. Use left hand to turn the
stopper and right hand to gently swirl the conical flask as acid is added. Look down at the colour in
the flask against the white tile
9. Stop as soon as there is a permanent change in the colour of the indicator
10. Read the level in the burette
11. Refill burette and repeat steps 5-10 until you have 3 titres within 0.1mL with each other
12. Ignore the first rough titre. Using the three consistent titres, calculate the average value
13. Calculate the concentration of the hydrochloric acid
Risks
Acid/base in eyes
Glassware can break and cut you
Spillages can cause slipping
Precautions
Safety glasses
Take care
Need to wipe up as soon as possible
*Note that the unknown and known can be placed in either conical flask or burette – it doesn’t
matter.
Choice of indicator
The indicator is chosen such that the pH at
the end point matches as closely as
possible with the pH at the equivalence
point of the titration.
The equivalence point is approximately in
the middle of the steep slope (inflexion
point). Note that weak acid to weak base
reactions are generally avoided due to the
fact that there is no rapid change in pH
even at the equivalence point.
Qualitatively describe the effect of buffers with reference to a specific example in a natural system
A buffer is a solution that contains comparable amounts (roughly molar equivalent) of a weak acid
and its conjugate base, and is therefore able to maintain a relatively constant pH when small
quantities of acid or base are added. The equilibrium involved is:
HA + H2O
H3O+ + A-
If hydronium ions are added, then by LCP, the equilibrium shifts to the left; the base A- will combine
with much of the added hydronium ions to form HA, in an attempt to minimise the change in
hydronium ion concentration. More of the un-ionised (weak) molecular acid is produced, thus
minimising pH change.
If instead, hydroxide ions are added, it would react with the hydronium ions, causing pH to rise;
however, by LCP the equilibrium will shift to the right and more A- will be produced. Since more of
the hydronium ions is then produced, the pH change is minimised.
Example in natural system
An example of a natural buffer system is the carbon acid system. It occurs natural in freshwater lakes
and rivers (that contain carbonate rocks from which HCO3‫ ־‬ions can be formed), maintaining the
constant pH of 6.5-7.5 needed for life to exist. The equilibrium involved is:
H2CO3 (aq) + H2O (l)
H3O+ (aq) + HCO3‫( ־‬aq)
Carbon dioxide from the air dissolves in the water, forming carbonic acid:
CO2 (aq) + H2O (l)
H2CO3 (aq)
Meanwhile, its conjugate base, HCO3‫־‬, is present as ions leeched out of rocks and minerals
containing the lake/river. These processes produce comparable amounts of H2CO3 / HCO3‫ ־‬to
produce a buffer system.
There needs to be a source of HCO3‫ ־‬apart from dissolved CO2 in order to produce this buffer system.
Therefore, rainwater is not a buffer (as the only source of HCO3‫ ־‬is through the ionisation of H2CO3)
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
A functional group is a classified group of atoms in a molecule which is responsible for the
characteristic properties of that molecule.
The same functional group with undergo the same/similar chemical reactions regardless of the size
of the molecule. Alkanols and alkanoic acids are part of the functional group: organic compounds.
Alkanols
Alkanols are a homologous series (a series of compounds with the same general formulae) that are:
 Derived from alkanes
They have the general formula:
CnH2n+1OH
 Contain the hydroxyl group (-OH)
 Neutral
Alkanoic acids
Alkanoic acids are considered to be weak organic acids that are:
 Derived from alkanes
 Contain the carboxylic functional group (-COOH)
 Acidic (H+ ions disassociate from COO- in solution)
Alkanol
- Possesses a bent geometry around the
oxygen atom
- Two bond dipoles: C-O and O-H bonds
- Polar molecules
Alkanoic acid
- An oxygen is double bonded to a central
carbon, and an –OH group is singlebonded to the same carbon. (Note the –
OH here is not called a hydroxyl group)
- The carboxyl group is always at the end
of the chain
- More polar than alkanol due to polar CO, O-H C=O bonds
Explain the difference in melting point and boiling point caused by straight-chained alkanoic acid
and straight-chained primary alkanol structures
The strength of intermolecular forces determines the boiling and melting points of a molecule. The
stronger the intermolecular forces, the more energy needed to break these forces (i.e. higher MP
and BP)
 For the same number of carbons in a straight carbon-chain, the highest BP’s and MP’s
belong to: the alkanoic acids, then the alcohols and then the parent alkanes.
 There is a general trend of increasing BP/MP with increasing molecular weight. This is
because as MW increases along a homologous series, the SA increases and therefore so too
does the number and strength of dispersion forces.
Alkanes
The only intermolecular forces
are dispersion forces. They are
very weak; hence the BP and
MP’s of alkanes are low.
Alkanols
Alkanols experience dispersion
forces, but also dipole-dipole
interactions due to the polar CO and O-H bonds creating a net
dipole. Furthermore, alkanols
can interact through hydrogen
bonding (due to the OH bond).
The intermolecular forces in
alkanol molecules require more
energy to overcome – giving
rise to higher MP and BP’s.
Alkanoic acids
Like alkanols, alkanoic acids
have dispersion forces, dipoledipole interactions and
hydrogen bonds; however the
additional double bonded C=O
means it is even more polar
than alkanols – resulting in
stronger dipole-dipole forces.
The intermolecular forces in
alkanoic acids require the most
energy to overcome – thus it
has the highest MP and BP’s.
*the hydrogen bonding in alkanols and alkanoic acids also allows for their high solubility in water
Identify the IUPAC nomenclature for describing the esters produced by reactions of straightchained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8
Naming esters:
 Alkanol part first
 Alkanoic part last
 Write as two separate words (alkyl alkanoate)
Drawing esters (includes structural and condensed structural formulae)
 Alkanoic acid first
 Alkanol part second
Example
Above – Propyl acetate (CH3COOCH2CH2CH3)
*Exceptions when naming alkanoic acids are: methanoic and ethanoic acid, for which the preferred
IUPAC names are formic and acetic acid respectively
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. It involves an acid-catalysed condensation reaction
between an alkanoic acid and an alkanol (or more generally between a carboxylic acid and an
alcohol)
Esters are sweet-smelling, volatile organic compounds which contain the ester functional group
“-COOC-”
Note, important steps when writing out equations of esterification
 Write equilibrium arrow
 Write catalyst on top of the equilibrium arrow
 Write water as a product
Example of esterification
Describe the purpose of using acid in esterification for catalysis
Explain the need for refluxing during esterification
Esterification reactions are usually quite slow and the reaction does not proceed to completion as it
is a reversible reaction in equilibrium. In order to increase the rate of reaction, we:
 Add a strong, concentrated acid as a catalyst
 Heat the reaction mixture
Purpose of the acid
The purpose of the acid is to decrease the time for the reaction to reach equilibrium and to
increase yield.
The concentrated, sulphuric acid acts as a catalyst as well as a dehydrating agent. The acid speeds
up the rate of reaction by lowering the activation energy. Furthermore, by removing water from the
reaction, the equilibrium is shifted towards the products; thus increasing the yield of product –
although this increase is quite small.
The need for refluxing
By heating the reaction mixture (increasing the
kinetic energy of the reaction), increasing the
amount of collisions between molecules and thus
increase the rate of reaction. Refluxing, therefore,
maximises the rate of reaction by allowing the
mixture to be heated at B.P. without losing any of
the reactants.
The reactants and products of the esterification
reaction are flammable and volatile, and so we
cannot use a naked flame such as that of a Bunsen
burner (flame is too hot, and could ignite the
reactants/products). These issues can be avoided
by heating the mixture using a reflux apparatus.
The mixture is heated by a hot-water bath supported by an electric hot plate. A water condenser is
mounted above the reaction flask (round bottomed flask), and cold water circulates to cool the hot,
rising vapours. The vapours condense back to liquid state and drip back into the reaction flask. This
process therefore allows the heating of the reaction at BP, without losing any of the reactants or
products.
The system is also open to the atmosphere to avoid build- up of pressure due to the production of
vapours. Small, boiling chips are dropped in the reaction mixture to enable a large SA for
vaporisation to occur without the risk of superheating and the explosive ejection of vapours.
Outline some examples of the occurrence, production and uses of esters
Esters occurs widely in both nature and are produced synthetically in industry.
In nature, esters are found in
 ATP (important in reactions in cells)
 Fats and oils
 Natural waxes
Esters are also manufactured for a wide range of industrial applications




Lubricants
Solvents
Plasticisers
Flavouring/fragrances
Occurrence
Natural fatty acids
Phthalate esters
Production
Produced from natural oils such
as soybean and canola using
transesterification. An ester is
formed by reacting vegetable oil
with sodium hydroxide dissolved
in methanol. After processing, it
can be used as biofuel in diesel
engines.
Produced from the reaction
between Phthalic acid and a
dialkyl.
Use
Alternate source of fuel (biofuel)
Plasticiser (added to plastics to
soften and increase flexibility)
Process information from secondary sources to identify and describe the uses of esters as flavours
and perfumes in processed foods and cosmetics
The characteristic tastes and smells of fresh fruits and flowers are due to a complex mixture of
chemicals, including esters. As a result, esters are often used as flavour enhancers and in fragrances
in cosmetics.
For example:
 The ester octyl acetate, which is found in oranges and other citrus fruits, is often used in
lollypops or soft drinks to give a fruity orange flavour. The esterification reaction involves 1octanol and acetic acid.
 Methyl salicylate (wintergreen oil) is an organic ester produced by wintergreens, and is
synthetically produced for use in lotions and creams to soothe sore muscles. It is an ester of
salicylic acid and methanol.