9.3
THE ACIDIC ENVIRONMENT
1. Indicators were identified with the observation that
the colour of some flowers depends on soil composition.
1.2.1 – Classify common substances as acidic, basic or neutral.
1.2.2 – 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.
1.3.2 – Identify data and choose resources to gather information about the colour ranges of a range of
indicators.
Chemical substances can be classified as acidic, basic or neutral. This classification is usually
based on the concentration of hydrogen ions produced when the substance is dissolved in water.
The concentration of hydrogen ions can be determined by change in colour of dyes called
indicators. Examples of indicators include litmus, phenolphthalein, methyl orange and
bromothymol blue.
Methyl orange
Bromophenol Blue
Bromocresol Blue
Methyl Red
Bromothymol Blue
Phenol Red
Thymol Blue
Phenolphthalein
Red – Yellow
Yellow – Blue
Yellow – Blue
Pink – Yellow
Yellow – Blue
Yellow – Red
Yellow – Blue
Colourless – Red
pH 3.1 – 4.4
pH 3.0 – 4.6
pH 3.8 – 5.4
pH 4.4 – 6.0
pH 6.2 – 7.6
pH 6.8 – 8.4
pH 8.0 – 9.6
pH 8.3 – 10
1.2.3 – Identify and describe some everyday uses of indicators including the testing of soil acidity/basicity.
Everyday uses of indicators include the monitoring of the pH-level of fish-tank environments,
garden soil and swimming pools.
1.3.1 – Perform a first-hand investigation to prepare and test a natural indicator.
1.3.3 – Solve problems by applying information about the colour changes of indicators to classify some
household substances as acidic, neutral or basic.
A natural indicator is the dye in red cabbage leaves. The leaves were shredded and the dye
extracted by placing the shredded leaves into minimal hot water. The clear purple solution was
strained. A small amount of the solution was applied to test tubes containing sodium hydroxide
(reference base); hydrochloric acid (reference acid); ‘Windex’ (window cleaner); white vinegar;
generic shampoo; lemon juice and milk.
The colour of the reference base became green, then clear yellow after the introduction of the
indicator; the reference acid became a clear hot pink. ‘Windex’ went from a clear light blue to a
clear deep green (basic). White vinegar, lemon juice and shampoo all became pink (acidic). Milk
was unable to be determined.
-I-
2. While we usually think of the air around us as neutral, the atmosphere naturally
contains acidic oxides of carbon, nitrogen and sulfur. The concentrations of these
acidic oxides have been increasing since the Industrial Revolution.
2.2.1 – Identify oxides of non-metals which act as acids and describe the conditions under which they act as
acids.
2.2.2 – Analyse the position of these non-metals in the Periodic Table and outline the relationship between
the position of elements in the Periodic Table and acidity/basicity of oxides.
Non-metal oxides include sulfur dioxide, carbon dioxide, nitrogen dioxide and nitrous dioxide.
They act as acids when they ionise water. Generally, non-metallic oxides are acidic. However,
some non-metallic oxides are neutral.
2.2.3 – Define Le Chatelier’s principle.
Le Chatelier’s Principle states: ‘when a change is made to an equilibrium system, the system
moves to counteract the imposed change and restore the system to equilibrium’.
2.2.4 – Identify factors which can affect the equilibrium in a reversible reaction.
Factors which will affect equilibrium include concentration, gas pressure and temperature. If the
temperature increases, the endothermic reaction will be favoured. If the temperature decreases,
the exothermic reaction will be favoured. If the concentration of a certain reactant of the equation
is increased, the system will shift to convert that reactant into product(s). Vice-versa, if the
concentration is decreased, the system will shift to convert products back into the reactant. If the
pressure of a system increases, the shift will be towards the reaction with the least number of
moles.
2.2.5 – Describe the solubility of carbon dioxide in water under various conditions as an equilibrium
process and explain in terms of Le Chatelier’s principle.
The solubility of carbon dioxide in water under various conditions can be understood under Le
Chatelier’s principle. The equilibria involved are:
CO2(g)
CO2(g)
+

H2O(l)
CO2(aq)
+

H2CO3(aq)
heat
+
heat
As the pressure of the system increases, more carbon dioxide is dissolved, because the
equilibrium will favour the reaction with least amount of moles of products. As the pressure
decreases, carbon dioxide escapes (into the air); the system counteracts this by evolving more
carbon dioxide gas from solution.
As the temperature of the system increases, the endothermic reaction is favoured. Therefore,
more carbon dioxide gas will be formed. As the temperature decreases, the exothermic reaction
is favoured; therefore more carbon dioxide is dissolved.
The solubility will increase in the presence of a base such as hydroxide ions. Hydroxide ions (OH –)
will decrease the concentration of hydrogen ions (H +) formed by carbonic acid (H2CO3(aq)).
Therefore the system will shift to produce more acid and dissolve more carbon dioxide.
- II -
2.2.6 – Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen.
2.2.7 – Describe, using equations, examples of chemical reactions which release sulfur dioxide and chemical
reactions which release oxides of nitrogen.
2.3.2 – Analyse information from secondary sources to summarise the industrial origins of the above gases
and evaluate reasons for concern about their release into the environment.
Industrial origins:
The sulfur dioxide in the air comes from the burning of coal in coal-powered/oil-fired power
stations. The smelting of metals such as zinc, copper and lead also produces sulfur dioxide from
their ores. Sulfur is oxidised into sulfur dioxide by the following equation:
–S–(s)
Sulfur in Compound
+
O2(g)
Oxygen in air
→
SO2(g)
Sulfur Dioxide
The incomplete combustion of fossil fuels releases large amounts of carbon-monoxide. The fact
that the air is not 100% O2 contributes to incomplete combustion in closed or unventilated
environments. Carbon monoxide is formed by cars and other combustion reactions with the
following equation:
–C–(l)
+
→
½O2(g)
CO(g)
Due to the high temperatures of an internal combustion engine, and the rapid cooling of exhaust
gas, NO(g) is prevented from decomposing. Nitrogen dioxide is formed by a secondary reaction of
nitrogen oxide with oxygen:
2NO(g)
+
→
O2(g)
2NO2(g)
Reasons for Concern:
Sulfur dioxide and the oxides of nitrogen, along with carbon dioxide, are largely responsible for
acid rain:
+
H2O(l) →
Water
H2SO3(aq)
Sulfurous Acid
2NO2(g)
+
Nitrogen dioxide
H2O(l) →
Water
HNO3(aq)
Nitric Acid
CO2(g)
+
Carbon dioxide
H2O(l) →
Water
H2CO3(aq)
Carbonic Acid
SO2(g)
Sulfur dioxide
+
HNO2(aq)
Nitrous Acid
The introduction of acidic solutions into rain water can cause it to drop in pH to about pH2. Long
term exposure to this is damaging to flora, since they continually absorb acid rain via diffusion. To
freshwater animals, the pH can release toxic metal ions which will kill them.
To the built environment, acid rain can attack structures made of marble, metal and sandstone.
For example, bridges are subject to corrosion in this way:
Fe(s)
Iron in bridge
+
H2CO3(aq)
→
Carbonic Acid
FeCO3(s)
+
Oxidation of Iron
H2(g)
Hydrogen
Further damage is caused with an increase of photochemical smog and respiratory problems.
Evaluation:
- III -
It is clearly seen that the introduction of non-metal oxides into the atmosphere contributes to the
destruction of the natural and built environment. If unchecked, the devastation will only continue
to grow as the burning of fossil fuels increases. Therefore, the release of non-metal oxides needs
to be restricted and, if possible by further industrial reactions, eliminated completely.
2.2.8 – Assess the evidence which indicates increases in the atmospheric concentrations of these gases.
The evidence which indicates increases in atmospheric concentrations of these gases have been
gathered from measurements on bubbles of ancient air trapped in ice sheets of Antarctica and
Greenland. This is then compared to more modern levels. This evidence is supported because
the bubbles of air remain unchanged. Concentration of air is constant around the world. The only
problem which may arise is slow diffusion of gases may occur in the air pockets.
2.2.9 – Calculate volumes of gases given masses of some substances in reactions, and calculate masses of
substances given gaseous volumes, in reactions involving gases at 25°C and 100kPa.
2.3.1 – Identify data, plan and perform a first hand investigation to decarbonate soft drink and gather data
to measure the mass changes involved and calculate the volume of gas released at 25°C and 100kPa.
A 390mL soft drink was heated in water until it was approximately 25°C. It was then weighed and
mass recorded on the triple beam balance. The soft drink was agitated by shaking and opened to
release carbon dioxide gas. Subsequent weighing, shakings and recordings of mass were done
until the mass reached a constant.
To calculate the volume of CO2(g) released, the mass of CO2(g) was recorded (by subtraction). It
was then substituted into the following formula:
n(CO2)
=
=
≈
m(CO2)
3.7
0.084 moles
÷
÷
Mm(CO2)
44
v(CO2)
=
=
=
n(CO2)
0.084
2.06 litres
×
×
SLC
24.47
2.2.10 – Explain the formation and effects of acid rain.
The formation of acid rain is in the evolution of non-metal oxides such as the production of SO 2(g)
by the burning of coal, the production of CO2(g) and CO(g) by the combustion of fossil fuels and the
production of NO2(g) by the combustion of coal, oil and petrol. Nitrogen dioxide also is formed by
soil processes, volcanic activity and lightning.
These gases, when reacted with water in the atmosphere, result in acidic solutions, and as rain,
acid rain. For example:
SO2(g)
Sulfur dioxide
+
H2O(l) →
Water
H2SO3(aq)
Sulfurous Acid
The effects of acid rain include the destruction of plants, which roots are constantly bathed in acid
rain through diffusion; the destruction of fresh water animals, whose metabolism and well-being
are affected by their limited water environment and pH-change intolerance; the destruction of the
human built environment, such as buildings with sandstone, limestone, marble and metals. For
example, a marble statue can be eroded by acid rain:
CaCO3(s)
+
Calcium carbonate
Acid Rain
→
- IV -
Ca(HCO3)2(aq)
Calcium hydrogen carbonate (soluble)
3. Acids occur in many foods, drinks and even within our stomachs.
3.2.1 – Define acids as proton donors and describe the ionisation of acids in water.
Acids can be defined as proton donors because they contribute a proton (a hydrogen ion) to
water when they ionise. For example:
CH3COOH(aq)
Acetic acid
+
H2O(l) 
Water
H3O+(aq)
+
Hydronium ion
CH3COO–(aq)
Acetate ion
3.2.2 – Identify acids including acetic (ethanoic), citric (2–hydroxypropane–1, 2, 3–tricarboxylic),
hydrochloric and sulfuric acid.
Acetic Acid:
(Ethanoic acid)
CH3COOH(aq)
Citric Acid:
(2–hydroxypropane–1, 2, 3–tricarboxylic acid)
(COOH)CH2CH(OH)(COOH)CH2COOH(aq)
Hydrochloric Acid
(Hydrochloric acid)
HCl(aq)
Sulfuric Acid
(Sulfuric acid)
H2SO4(aq)
3.2.3 – Describe the use of the pH scale in comparing acids and bases.
3.2.5 – Identify pH as –log10[H+] and explain that a change in pH of 1 means a ten-fold change in [H+].
The pH scale is used to compare acids and bases, based on the concentration of [H +] in solution.
pH is measured from 0 to 14 (acidic to basic). The pH scale is a logarithmic relationship, the
formula being:
pH
=
–log10[H+]
The pH of a base is derived by the following:
pOH
pH
=
=
–log10[OH–]
14
–
pOH
where [OH–] is the concentration of hydroxide ions in the solution.
3.2.4 – Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated
and dilute.
3.2.6 – Compare the relative strengths of equal concentration of citric, acetic and hydrochloric acid and
explain in terms of the degree of ionisation of their molecules.
Acids ionise in water and become proton donors, forming [H +] ions in water. The greater the
concentration of [H+], the greater is the strength of the acid. The negative logarithmic relationship
only works for strong acids, those which ionise completely in water. A weak acid is an acid which
only partially ionises in solution. A concentrated acid can be either strong or weak. Concentration
and weakness refer to the level of dilution of the solution using a solute, usually water.
Between hydrochloric, acetic and citric, the strong acid, hydrochloric, is strongest. The weakest
acid is the weak acid, acetic, and citric acid is between the strengths of the two.
3.2.7 – Describe the difference between a strong and weak acid in terms of an equilibrium between the
intact molecule and its ions.
-V-
In a strong acid, such as hydrochloric acid, an equilibrium is formed during ionisation:
HCl(aq)
+
Hydrochloric acid
H2O(l)
Water

H3O+(aq)
+
Hydronium ion
Cl–(aq)
Chloride ion
In the equilibrium of the strong acid, the equation completely lies on the right side (near 100%
ionisation). The molecule of the strong acid completely ionises.
In a weak acid, such as acetic acid, an equilibrium is formed during ionisation:
CH3COOH(aq)
Acetic acid
+
H2O(l) 
Water
H3O+(aq)
+
Hydronium ion
CH3COO–(aq)
Acetate ion
In the equilibrium of the weak acid, the equation lies mostly on the left (partial ionisation). The
molecule of the weak acid is in solution with few of its ions.
3.3.1 – Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to
distinguish between acidic, basic and neutral chemicals.
3.3.2 – Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong
and weak acids.
4.3.2 – Choose equipment and perform a first-hand investigation to identify the pH of a range of salt
solutions.
The three dot points were simultaneously completed by a pH recording of a range of acids, bases
and salts using a pH probe connected to a data-logger and various indicators.
It was clearly seen that the acidic solutions have their pH below 7. This was shown accurately by
the probe and qualitatively by the indicators. The results for pH were accurate and registered the
0.1M solutions of strong acids (HCl, H2SO4 and HNO3) at about pH 1. The results also clearly
showed the difference in pH between strong and weak acids. Since they were in identical
concentration, the difference in ionisation levels attributes to a difference in pH.
The pH of salts was tested with 0.1M solutions. The ‘neutral’ salts were identified (NaCl, Na 2SO4
and KNO3) and were neutral because of the inert nature of their ions. A basic salt, CH3COONa,
was identified. It is basic because of the presence of hydroxide ions during its dissociation:
CH3COONa(aq) →
CH3COO–(aq)
+
CH3COO–(aq)
H2O(l) →
- VI -
+
Na+(aq)
CH3COOH(aq) +
OH–(aq)
4. Because of the prevalence and the importance of acids,
they have been used and studied for hundreds of years.
Over time, the definitions of acids and base have been refined.
4.2.1 – Outline the historical development of ideas about acids including those of: – Lavoisier, –Davy, –
Arrhenius.
4.2.2 – Outline the Brönsted-Lowry theory of acids and bases.
4.3.1 – Gather and process information from secondary sources to trace development in understanding and
describing acid/base reactions.
The first definition of acids and bases were derived from our taste of them. Acids tasted sour,
hence it was derived from the Latin word: ‘acer’, meaning sharp. Bases, to humans, tasted bitter.
In 1776, the Frenchman Antoine Lavoisier discovered that many compounds with oxygen (nonmetallic oxides) developed acidic properties in solution. He defined an acid as a non-metal
compound containing hydrogen.
In 1810, Humphry Davy showed that HCl has acidic properties and contained no oxygen. He
redefined an acid as a substance containing hydrogen.
In 1894, Svante Arrhenius redefined the theory to explain that acids contain H + ions which are
liberated in solution. A base, he said, is a substance which contains OH – ions in solution.
However this theory does not account for all bases, such as sodium carbonate (Na2CO3).
In 1923, the Danish scientist Johannes Brönsted and Englishman Thomas Lowry proposed
independent theories that an acid is a proton donor, a base a proton acceptor.
4.2.3 – Describe the relationship between an acid and its conjugate base and a base and its conjugate acid.
4.2.5 – Identify conjugate acid/base pairs.
When Brönsted and Lowry proposed their theories they established a relationship between acids,
bases and their conjugate pairs. This definition recognised the important role of the solvent in
determining acidity or basicity. Acids ionise in water due to interaction with water molecules. The
acidic proton is donated to the water molecule (a base) to form the hydronium ion (the conjugate
base). For example:
HCl(aq)
ACID
+
H2O(l)
BASE
→
H3O+(aq)
CONJUGATE ACID
+
Cl–(aq)
CONJUGATE BASE
Each acid has a conjugate base, and each base a conjugate acid.
4.2.6 – Identify amphiprotic substances and construct equations to describe their behaviour in acidic and
basic solutions.
In Brönsted-Lowry theory some substances are capable of behaving as either acids or bases.
Such substances are said to be amphiprotic (either accepting or donating a proton). An example
of this is:
HCO3–(aq)
+
H3O+(aq)
→
H2CO3(aq)
+
H2O(l)
BASE
ACID
ACID
BASE
HCO3–(aq)
ACID
+
OH–(aq)
BASE
→
- VII -
CO32–(aq)
BASE
+
H2O(l)
ACID
Water, HSO4– and HPO4– are also amphiprotic.
Amphoteric substances are those which react with both acids and bases.
4.2.7 – Identify neutralisation as a proton transfer reaction which is exothermic.
Neutralisation is a proton transfer reaction in which an acid and a base react to form a salt and
water. Neutralisation reactions are exothermic. The amount of heat liberated when neutralisation
occurs depends on the strengths of the acid and the base.
4.2.8 – Describe the correct technique for conducting titrations and preparation of standard solutions.
4.3.3 – 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.
The volumetric analysis that was conducted was to find the molarity of an unknown sodium
hydroxide solution.
Firstly 250mL of primary standard of 0.05 mol L–1 sodium carbonate was needed. Anhydrous
sodium carbonate was spooned into a 100mL beaker on a 2 decimal place zeroed electronic
balance. The mass was calculated to be approximately 1.325 grams. Water was added and the
solution was swirled to dissolve the sodium carbonate. It was then transferred to a volumetric
glass. More water was added to the beaker by a wash bottle to ensure all of the sodium
carbonate went into the volumetric flask. The volumetric flask was then swirled and filled to the
250mL mark. This was the primary standard.
0.05 mol L–1
0.05
×
0.0125 moles
Molarity required of Na2CO3
In 250mL, n(Na2CO3)
=
=
=
Mass
number of moles
0.0125
1.325 grams
=
=
≈
0.25
×
×
Molar mass
105.99
The secondary standard was a solution of approximately 1 molar hydrochloric acid solution. The
acid was diluted to make 250mL of approximately 0.1 mol L –1 hydrochloric acid solution. This was
our secondary standard.
A pipette and conical flask was rinsed with small amounts of the primary standard. 25mL of the
primary standard was added by pipette to the conical flask. 2 drops of the indicator methyl red
was added to the flask. A small amount of the secondary standard was used to rinse the burette.
50mL of the secondary standard was then added to the burette and used to titrate the primary
standard.
Titrating, it was found that 23.35mL of ~0.1M HCl reacted with 25ml of Na2CO3.
CHCl
×
VHCl
=
2
×
CHCl
=
=
2 × (0.05) × (25) ÷ 23.35
0.107 mol L–1
CNa2CO3
×
VNa2CO3
A fresh burette was rinsed with the unknown sodium hydroxide solution. 50mL of sodium
hydroxide was then added to it. The secondary standard was used to rinse a new conical flask,
and 10mL of it was added by pipette to the flask. 2 drops of phenolphthalein was used as
indicator.
- VIII -
Titrating, it was found that 13mL of NaOH reacted with 10mL of HCl
CHCl
×
VHCl
=
CNaOH
×
VNaOH
=
=
VNaOH
10 × (0.107) ÷ 13
0.082 mol L–1 (3dp)
By using volumetric analysis, titration and several indicators, the concentration of HCl could be
standardised and thus unknown concentration NaOH.
4.2.9 – Qualitatively describe the effect of buffers with reference to a specific example in a natural system.
A buffer solution is a solution which contains comparable amounts of a weak acid and its
conjugate base and which is therefore able to maintain an approximately constant pH even when
significant amounts of strong acid or strong base are added to it.
The following buffer is found in human blood:
H2CO3(aq)
+
H2O(l)

H3O+(aq)
+
HCO3–(aq)
If an acid is added the concentration of the H 3O+ ions suddenly rises. By Le Chatelier’s Principle,
the equilibrium shifts to the left to counteract the change. This uses up the H 3O+ ions and returns
the pH to where it was.
If a base is added the concentration of the H 3O+ ions suddenly falls. The equilibrium shifts to the
right to counteract the change. This uses up the OH – ions present in the base, and returns the pH
to where it was.
4.3.3 – Perform a first-hand investigation to determine the concentration of a domestic acidic substance
using computer-base technologies.
This experiment was conducted similarly to the volumetric analysis experiment. A known
concentration of sodium hydroxide solution was used to titrate vinegar (acetic acid). The pH
probe and data-logger were set up and the program logged the pH by a chart and graph. As
expected the pH of the equivalence point was greater than 7 in this weak acid/strong base
titration.
4.3.4 – 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.
Neutralisation can be used as a safety measure or to minimise damage in spills but only weak
acids or bases should be used. An example of a weak base to clean up an acidic spill would be
calcium carbonate in powder form. Calcium carbonate is a safe solid powder; powder has a larger
surface area than chips, making it faster to react. The problem with neutralisation is that heat will
be built up, causing more damage.
- IX -
5. Esterification is a naturally occurring process
which can be performed in the laboratory.
5.2.1 – Describe the differences between the alkanol and alkanoic acid functional groups in carbon
compounds.
5.2.3 – Explain the difference in melting and boiling point caused by straight-chained alkanoic acid and
straight-chained primary alkanol structures.
Both alkanols and alkanoic acids are polar molecules due to the presence of the –OH groups and
–COOH groups. The presence of hydrogen bonding in both types of molecules leads to higher
melting and boiling points in these molecules. As the chain length increases, dispersion forces
increase, leading to higher melting and boiling points. When alkanoic acids and alkanols of the
same number of carbons are compared, the more extensive hydrogen bonding and an increase
in dipole-dipole interactions mean that the alkanoic acid has a higher melting and boiling point
than its respective alkanol.
5.2.2 – Identify the IUPAC nomenclature for describing the esters produced by the reactions of straightchained alkanoic acids from C1 to C8 and straight-chained primary alkanols from C1 to C8.
5.2.4 – Identify esterification as the reaction between an acid and an alkanol and describe, using equations,
examples of esterification.
5.2.5 – Describe the purpose of using acid in esterification for catalysis.
Esterification is the process where an ester is formed by an alkanol and an alkanoic acid. They
are named as alkyl alkanoates, where the alkyl is the abbreviated form of the alkanol used and
alkanoate the form of the alkanoic acid. For example:
Methanol
+
Butanoic Acid

Methyl Butanoate
+
Water
A catalyst, concentrated sulfuric acid, is used to assist in the dehydration of water and to lower
the activation energy. It has been shown that it is the OH of the alkanoic acid that joins with the H
of the alkanol to produce water.
5.2.6 – Explain the need for refluxing during esterification.
5.3.1 – Identify data, plan, select equipment and perform a first-hand investigation to prepare an ester
using reflux.
About 15mL of 1–pentanol (C5H12O) and 15mL of ethanoic acid were added to a pear-shaped
flask, and 10 drops of concentrated (30M) sulfuric acid and porcelain boiling chips were added.
The flask was secured by a retort stand and bathed in a water bath. An inverted reflux
condensing tube attached to the flask. The equipment was heated by a bunsen set on a blue
flame for 15 to 20 minutes, when two layers were clearly visible.
The solution was poured into a separating funnel and 100mL of water was added. The mixture
was shaken vigorously and the lower aqueous layer was discarded by separation. Sodium
carbonate was added to neutralise excess acid and the funnel was swirled gently. The plug was
released to allow carbon dioxide to escape. Two layers again formed and the lower discarded.
Sodium carbonate was added again and the procedure repeated.
The remaining liquid, the ester, was poured into a conical flask and a teaspoon of calcium
chloride was added to remove excess water. After a few minutes, the ester was decanted. The
ester was then distilled with a distillation kit to further purify it.
The volatility of the reagents of esterification means that the extended heating must be performed
under reflux. As the flask is heated the components vapourise and are then condensed due to the
cooling provided by the condensing tube. The reagents thus return to the flask and eventually
-X-
equilibrium is achieved without lost of material. The final product, pentyl-ethanoate (a bananaflavoured ester), was produced.
5.2.7 – Outline some examples of the occurrence, production and uses of esters.
5.3.2 – Process information from secondary sources to identify and describe the uses of esters as flavours
and perfumes in processed foods and cosmetics.
Esters are added in many products, such as perfumes and cosmetics. Esters are used to add an
artificial ‘fruity smell’ to foods such as pentyl-ethanoate (banana), octyl-ethanoate (orange) and
pentyl-butanoate (apricot). Esters also are used as solvents for polar and non-polar compounds
and soaps. Aspirin (acetylsalicylic acid) is an example of an ester used in medicine.
- XI -