9.2 The Acidic Environment – Yr 12 2008
Written by Sam Jiang
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
Description
Substance which in
solution produces H+
or more strictly, H3O+
( hydronium ion)
Substances
Acid
Vinegar (acetic acid)
Citric fruit (citric acid)
Aspirin (acetyl salicylic acid)
Vitamin C (ascorbic acid)
Tartaric acid
Car batteries (sulfuric acid)
Base
Substance
which Bitter taste
Caustic soda for oven/drain
either contains the Soapy feel
cleaners (NaOH)
2soluble oxide
O
or Conduct electricity in Washing soda (Na2CO3)
Alkali hydroxide ion OH or solution (alkalis)
Bicarbonate of soda (NaHCO3)
which in solution Turn red limits blue
Antacid (CaCO3, Mg(OH)2)
which produces the
Lime for mortar (Ca(OH)2)
hydroxide ion
Garden lime (CaCO3)
+
Neutral Substance which has same amount of H and OH- ions, e.g. Water

Common properties
Sour taste
Sting or burn skin
Conduct electricity
Turn blue limits red
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
Indicators are chemical dyes that consist of organic molecules or ions with different
coloured acidic and basic forms, they may change colour as the acidity of the surroundings
changes.
The range of acidity or alkalinity over which indicators change colour varies from one
indicator to another. They generally have one colour when in Hln state and another colour
when is ln-. When changing from one state to another, the indicator changes colour.
Added to base:
Hln + OH-  H2O + lnAdded to acid:
ln- + H3O+  Hln + H2O

identify and describe some everyday uses of indicators including the testing of soil
acidity/basicity
Chemical
Used by chemists in analytical work
analysis
Soil testing
Monitoring
and
maintaining
pools
Monitoring
waste
e.g. detecting end point in titrations
Soil is moistened
Universal indicator is mixed with the sample
Sprinkle white insoluble barium sulfate BaSO4 on the surface
Match colour seen against PH/ colour chart
Liquid ammonia NH3 and Agricultural lime CaCO3 increases basicity, Ammonium
salts NH4+ increases acidity.
Pool water’s acidity level need to be monitored
Ensure skin/membranes are not irritated
Avoid growth of green algal scam in pool
Pool chlorine, Na2CO3, NaHCO3 and HCl, HNaSO4 are used
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
Oxides are a class of compounds that often display acidic or basic properties. Many non
metal oxides are acidic oxide. Acidic oxide either reacts with water to form an acid or reacts
with bases to form salts. Basic oxide reacts with acid to form salts.
Common oxides
Acidic Basic Amphoteric Neutral
Valencies
CO2
Na2O ZnO
CO
SO32- CO32SO2
K2 O
Al2O3
NO
NO3- NH4+
SO3
MgO PbO
N2 O
PO33- OHNO2
CaO
SnO
H3O+ SO42NO3
CuO
GeO2
NH3
P 2 O5
Ag2O Sb2O3
P 2 O3
FeO
Cl2O
Fe2O3

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
Most metals and non metals can react with oxygen and produce oxide.
Oxides of metals are basic, they are all ionic compounds.
Oxides of non metals are acidic, they are covalent compounds
Amphoteric oxides are formed by transition elements close to the non metals in the
period table.
Group 1 oxides are strongly basic, group 2 oxides are basic, basicity increases down the
group.
Most transition metal oxides are basic while some are amphiproteic
Group 3-6, trend down group from acidic, amphoteric to basic oxides.
Group 7 is strongly acidic, acidity decrease down the group

Define Le Chatelier’s principle
If the conditions of a system at equilibrium are changed, the system will readjust itself
to oppose the change and re-establish equilibrium.
If a closed system at equilibrium is disturbed, then the system adjusts itself so as to
minimise the disturbance.

Identify factors which can affect the equilibrium in a reversible reaction
Temperature, increase will favour endothermic reaction, decrease will favour exothermic
Volume, Increase favour the side with more moles of gas
Pressure, increase favour the side with less moles of gas
Concentration, favour to the side with less concentration
Surface area for heterogeneous reaction

Describe the solubility of carbon dioxide in water under various conditions as an
equilibrium process and explain in terms of Le Chatelier’s principle
CO2 + H2O  H2CO3 Exothermic
Pressure
If the pressure increase, the equilibrium is disturbed, therefore the reaction tries to oppose
this change by favouring the side with less number mole of gas, and the reaction will shift
toward to the product side, more CO2 is dissolved, producing more H2CO3 and decrease the
pressure to minimise the change.
If the pressure decrease, the equilibrium is disturbed, therefore the reaction tries to oppose
this change by favouring the side with more number mole of gas, and the reaction will shift
toward to the reactant side, releasing more CO2 gases into the air and increase the pressure
to minimise the change.
Temperature
If the temperature increase, the equilibrium is disturbed, since the reaction is exothermic, it
will oppose this change by favouring the reverse reaction that absorbs heat, decreasing the
temperature to minimise the change. Therefore the reaction will shift toward the reactant
side.
If the temperature decrease, the equilibrium is disturbed, since the reaction is exothermic, it
will oppose this change by favouring the forward reaction that liberates heat, increasing the
temperature to minimise the change. Therefore the reaction will shift toward the product
side.
Adding alkali
CO2 + H2O  H2CO3  H+ + HCO3-  2H+ + CO3-
When a base is added, it will neutralise the H+ and form water, decreasing the concentration
of H+. In order to re-establish equilibrium, more H2CO3 will move to the right to become H+,
increasing the concentration of H+ , however this will decrease the concentration of H2CO3
and thus CO2 need to move to the right and form H2CO3.

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

Identify natural and industrial sources of sulfur dioxide and oxides of nitrogen

Explain the formation and effects of acid rain
sulfur dioxide
oxides of nitrogen
industrial sources
Power plant combusting fossil High temperature combustion in
fuels or sulfur containing fuels
motor vehicles (petrol engine),
Combustion of fossil fuels in homes, industries and power plant.
motor vehicles
Metallic ore smelting facilities
Wood and paper industrial
processors
Oil and gas refineries
Natural sources
Volcanoes
Lightening
Hot spring
Combustion of organic matter,
e.g. bushfire
reasons for concern React with water in atmosphere and precipitate as acid rain.
Harm or kill marine animals such as fish by increasing PH in water
and creating a marine environment unsuitable for reproduction and
survival. PH below 5 fish will stop laying eggs, lower PH will kill
them, affecting all animals in the food chain.
Reduces crop yields both by damaging foliage and by causing the
leaching of minerals and important nutrients needed for plant growth
from the soil and releasing toxic heavy metals. It is causing the
defoliation of pine plantations, which is environmentally detrimental
and destroying large amount of forest.
Buildings made of steel are damaged by acid rain, steel bridge and
building needs to be painted to prevent this.
Building, statues stones and Lung damage, other respiratory
monuments made by marble are illness particularly for asthma
damaged by acid rain, forming patients.
gypsum
Fatigue, nausea, cough and
CaCO3(s) + H2SO4(aq)  vomiting. Lower resistance to
CaSO4(aq) + CO2(g) + H2O(l)
respiratory infections such as
Calcium sulfate will flakes off, influenza.
damaging the building.
Photochemical smog
Difficulty in breathing, choking
sensation. Chronic lung, heart
However the major source of oxides of nitrogen and sulphur dioxide is from burning fossil
fuels for industrial operations and travelling both on the road and on air. These operations
have developed society enormously, and created huge impacts on transport and living
standards. Thus the negative effects of emissions must be balanced with the benefit received.
Alternative energy sources are researched at the moment and methods to reduce the level of
acidic oxides in the environment

Describe, using equations, examples of chemical reactions which release sulfur
dioxide and chemical reactions which release oxides of nitrogen
2ZnS(s) + 3O2(g)  2ZnO(s) + 2SO2(g)
(burning sulfur compound)
S(s) + O2(g)  SO2(g)
H2O(l)+ SO2(g)  H2SO3(aq)
2SO2(g) + O2(g)  2SO3(g)
H2O(l)+ SO3(g)  H2SO4(aq)
N2(g) + O2(g)  2NO(g) (lightening)
2NO(g) + O2(g)  2NO2(g)
H2O(l) + 2NO2(g)  HNO2(aq) + HNO3(aq)

Assess the evidence which indicates increases in atmospheric concentration of
oxides of sulfur and nitrogen
“Increased” erosion of stone, marble monuments and buildings.
Average concentration of SO2 and NO2 in most large cities worldwide is about 0.01ppm for
each gas. This compares with measurements of 0.001ppm in other populated parts of the
earth and as low as 0.00005 ppm in areas well away from human activity, electrical storms
and volcanic action. This implies that concentration has increased.
3. Acids occur in many foods, drinks and even within our stomachs

Define acids as proton donors and describe the ionisation of acids in water
Acids are proton donors because in solution, they give up a hydrogen in or hydronium ion to
the solution. Most hydrogen atoms contain no neutrons, but only a single proton in their
nucleus. In explaining acidity, the terms “hydrogen ion” and “proton” are used
interchangeably. Acids ionise in water to produce hydronium ion and a negative ion.

Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-
tricarboxylic), hydrochloric and sulfuric acid
acetic (ethanoic), CH3COOH
citric (2-hydroxypropane-1,2,3- tricarboxylic), C6O7H8
hydrochloric, HCl
sulfuric acid, H2SO4

Describe the use of the pH scale in comparing acids and bases
PH is an expression for the concentration of H+ ions in solution. However, PH measures the
free hydrogen ion concentration in solution, not the total ions in solution.
Neutral substances have pH around 7, acid is below 7, acidity increase as pH decrease, base
is above 7, basicity increase as pH increase.

Describe acids and their solutions with the appropriate use of the terms strong,
weak, concentrated and dilute

Describe the difference between a strong and a weak acid in terms of an
equilibrium between the intact molecule and its ions
Strong acids are those in which all the acid molecule present in solution has ionised to
hydrogen ions, it has a high degree of ionisation and ionise completely, containing no intact
molecules. The reactions are non reversible go to completion.
HCl, HI, HBr, HNO3, H2SO4
NaOH, KOH, Ba(OH)2, Ca(OH)2, Na2O
Weak acid are those in which only some of the acid molecules present in the solution has
ionised to form hydrogen ions. There is an equilibrium holds between intact molecule and
hydrogen ions as the reaction does not go to completion. Degree of ionisation increases for
weak acid as they are diluted.
CH3COOH, C6O7H8, HF, HNO2, H3PO4
NH3, CO32-,
Concentrated solution means the solution is really concentrated, high number of moles of
the acid is dissolved in an unit volume.
Dilute solution means the solution has a low concentration, low number of moles of acid is
dissolved in an unit volume.

Identify pH as -log10 [H+] and explain that a change in pH of 1 means a ten-fold
change in [H+]
PH = -log10 [H+]
[H+] = 10-PH
14
PH = 14 + log10 [OH-] = -log10 10 
[OH ]
Water partially ionises to produce the following equilibrium
H2O  H+(aq) + OH-(aq) [self ionisation equilibrium]
The ionisation constant of water
KW = [H+][OH-]
=1.0  10-14 at 25 degree

Gather and process information from secondary sources to explain the use of
acids as food additives
A diverse range of acids are used as food additives for various purposes:
Acetic acid is the major component of vinegar. It is used for pH reduction, controlling
microbial growth (fungi, moulds, yeasts, etc.) and flavouring.
Citric acid is used for flavouring, microbial growth and pH reduction, also used for artificial
lemon juice, citric flavoured drinks. Prevent fruit from browning by inactivates the enzymes
that cause discolour.
Malic acid imparts a smooth, tart taste and has an acid flavour, helping to mask the
aftertastes of low or non-caloric sweeteners. It is primarily used in carbonated beverages,
powdered juice drinks, jams and jellies, canned fruits and vegetables and confectionery.
Phosphoric acid is primary used in cola, root beer and similarly flavoured carbonated
beverages.
Tartaric acid occurs in grapes and wine. It contributes a strong tart taste, which enhances
fruit flavours, particularly grape and lime. It is often used as an acidulant in grape, wine and
lime-flavoured beverages, gelatine desserts, jam and jellies and hard sour confectionery.
As food additives, acids have the following main uses:
Flavouring: each acid has a set of taste characteristics. Generally, weak acids have a
stronger taste than strong acids because they exist in the primarily un-dissociated state.
PH regulation: lowering the pH of the food system enhances the effectiveness of food
preservatives. PH also determines the flavour and the microbial stability.
Buffering: this is important because a fluctuating pH can have an adverse effect on flavour,
colour and microbial stability. Common buffers: citric acid-sodium citrate.
Preservation: acids help slow down the growth of spoilage organisms, especially bacteria.
This effect very much depends on the pH value of the acid. Yeasts and moulds grow within
a large pH range, and the pH must be lowered to halt their growth.

Compare the relative strengths of equal concentrations of citric, acetic and
hydrochloric acids and explain in terms of the degree of ionization of their
molecules
Hydrochloric acid is stronger than citric acid, which is stronger than acetic acid.
HCl is a strong acid that completely ionises in solution, and thus it is stronger than citric and
ethanoic acid, which are weak acids.
Citric acid is a triprotic acid, which contains three replaceable hydrogen ions, and this
ionises to a greater extent than ethanoic acid when placed in solution. Thus, citric acid is
stronger than acetic acid.

Identify data, gather and process information from secondary sources to identify
examples of naturally occurring acids and bases and their chemical composition
HCl - produced by glands in the lining of our stomachs to form an acidic environment for
the efficient operation of the enzymes that break complex food molecules into easily
transportable small molecules that are absorbed into the blood stream when they pass into
the intestine.
Acetic acid (CH3COOH) – is present in vinegar(4%)which is commonly made from wine by
oxidation of ethanol, find naturally in apple, coca, grape, etc. Clear, colourless, distinctive
odour, dissolve in water.
Citric acid (2-hydroxypropane – 1,2,3 – tricarboxylic acid C6H8O7 ) – occurs in citrus fruit
such as organs, lemons and lime. It is soluble white powder.
Ascorbic acid (C6O8O6) – vitamin C. it occurs widely in fruits and vegetables and is an
essential part of our diet. Used as an antioxidant in food preservation.
Carbonic acid (H2CO3) – formed when carbon dioxide dissolves in water. Present in soda
water
Formic acid (HCO2H) – present in the stings of ants, bees and nettles
Lactic acid (CH3CH(OH)CO2H) – present in milk and yoghurt. Produced by muscles during
exercise.
Bases
Caffeine – C8H10N4O2 it is white crystalline powder with bitter taste found in coffee,
colanuts, guarana, tea.
Nicotine - C8H14N4 it is liquid obtained from tobacco leaves.
Quinine (CH3NH2 – natural base found in rotting fish
Ammonia (NH3) - Component of decaying organic matter, volcanic gases. It is colourless
with sharp odour.
Very few soluble bases exist freely in nature. However, calcium carbonate magnesium
carbonate are found in rocks such as limestone, marble and dolomite.
4. Because of the importance of acids, they have been used and prevalence and studied
for hundreds of years. Over time, the definitions of acid and base have been refined

Outline the historical development of ideas about acids including those of:
-Lavoisier
Antoine Lavoisier, in about 1779 proposed that acids were substances than contained
oxygen since he discovered that oxygen exists in many compound that displayed acidic
properties when dissolved in water. However this theory was soon disproved because many
oxygen containing substances were basic and some acidic substances were acidic but
contained no oxygen. He focused on composition rather than property of acids.
The presence of oxygen in compounds formed from non-metals causes acidity.
-Davy
Humphrey Davy, in 1810 suggested that acids were substances that contained hydrogen.
Bases were substances that reacted with acids to form salts and water. This redefined acids
in terms of the element hydrogen. These definitions worked quite well for most of that
century, yet this is still a definition by composition and did not explain why acids/bases
behave the way they do.
All acids contain hydrogen.
-Arrhenius
Svante Arrhenius in 1887 proposed that an acid was a substance which ionised in solution to
produce hydrogen ions and a base was a substance that in solution produced hydroxide ions.
This introduced the neutralisation equation and he proposed that it was these ions that gave
acids their characteristics properties. However this definition applies only to aqueous
solutions and excludes substances that are distinctly basic such as NH3 NaCO3 with no OHAn acid is a substance that provides H+ ions in aqueous solution
A base us a substance that provides OH- ions in aqueous solution

Gather and process information from secondary sources to trace developments in
understanding and describing acid/base reactions

Outline the Brönsted-Lowry theory of acids and bases
There are some inadequacies with Arrhenius definition of an acid, because it does not give
due recognition to the role of the solvent. Ionisation of the acid is not something the acid
does in isolation: rather it is a reaction between the acid molecule and the solvent. Whether
an acid is strong or weak depends not only upon the nature of the acid itself, but also upon
the nature of the solvent it is dissolved in: HCl in water is a strong acid, but when dissolved
in diethyl ether it is quite weak. Also, some compounds that do not contain hydroxide ions
are able to neutralise acids, such as ammonia, some salts can act as acids or bases and some
substances can act as both an acid and a base.
In 1923 two chemists. Lowry from Britain and Bronsted from Denmark, independently
proposed new definitions for acids and bases in terms of proton donors and acceptors: these
definitions overcame the difficulties regardless of solvent, and are the most widely used
definition today, as it explains acid/base by property and behaviour rather than composition.
An acid-base reaction is one in which a proton is transferred from an acid to base.

Describe the relationship between an acid and its conjugate base and a base and
its conjugate acid

Identify conjugate acid/base pairs
A base after it has received a proton has the potential to react as an acid, this is called
conjugate acid. Similarly an acid which has donated a proton is potentially a base, called
conjugate base.
Acid
Base
Strongest Acid
HCl
Cl
Weakest Base
H2SO4
HSO4
HNO3
NO3H 3 O+
H2 O
HSO4
SO42H3PO4
H2PO4CH3COOH
CH3COOH2CO3
HCO3H2 S
HSNH4+
NH3
H2 O
OHHSS2Weakest Acid
OHO2Strongest Base
Strong acid has a weak conjugate base
Strong base has a weak conjugate acid
Weak acid has a strong conjugate base
Weak base has a strong conjugate acid

Identify a range of salts which form acidic, basic or neutral solutions and explain
their acidic, neutral or bases basic nature
Salt can be acidic, basic or neutral. Acidic and basic salts contain an ion which behaves as
an acid or base.
Ammonium chloride: acidic, combination of a strong acid and a weak base.
HCl + NH3  NH4Cl
Ammonium nitrate: acidic, combination of a weak base and a strong acid
HNO3 + NH3  H3O+ + NO3- + NH3
Sodium carbonate: basic, combination of a strong base of a weak acid
H2CO3 + 2NaOH  Na2CO3 + 2H2O
Sodium acetate: basic, combination of a strong base and a weak acid
CH3COOH + NaOH  NaCH3COO + H2O
Sodium hydroxide: Neutral, combination of a strong base and a strong acid
Na+ + OH-  NaOH
Sodium chloride: Neutral, combination of a strong base and a strong acid
2HCl + NaOH  2NaCl + H2O

Identify amphiprotic substances and construct equations to describe their
behaviour in acidic and basic solutions
A substance which can act both as a Brönsted-Lowry acid (proton donor) and base (proton
acceptor) is called an amphiprotic substance.
Amphiprotic substances: H2O, HSO-, HCO3-, H2PO4-, HPO42-

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
Acids and alkalis are corrosive, they can damage skin, clothing and benches. When
accidents or spills occur, neutralisation reaction can be utilised to minimise damage.
Neutralisation is an exothermic process, therefore it cannot be used for spills on skin as it
will cause even more heat and intensify the burn.
Spills on floors and benches can be cleaned by neutralisation reaction. Weak base such as
sodium hydrogen carbonate NaHCO3 and weak acid such as acetic acid can be added to
neutralise the spill follow by a large amount of water. Being a weak acid/base, it won’t
produce much heat.
Sodium carbonate is widely used to neutralise acidic spills or effluents because:
It is a stable solid which is easily and safely handled and stored
It is the cheapest alkali available
If too much of it is used there is less danger than excess of sodium hydroxide calcium
hydroxide.
Amphiprotic substances such as sodium hydrogen carbonate can be used to neutralise both
acid and base.
Never add water to acid but acid to water.

Identify neutralisation as a proton transfer reaction which is exothermic
Neutralisation reactions are proton transfer reactions which are exothermic with final
products as water and a salt.

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

Describe the correct technique for conducting titrations and preparation of
standard solutions
Volumetric analysis is a form of chemical analysis in which the concentration (molarity) or
amount of a substance (unknown) is determined by measuring the volume of a solution of
known concentration of another substance (the standard) which is just sufficient to react
with all of the sample of the first substance
Procedure of adding the standard solution to the unknown solution until reaction is
completed is called the titration.
The point when the reaction is complete is called the equivalence point. With an acid-base
titration, there is no colour change and no precipitate formed to show us when the reaction
is completed, so an indicator is used. The point at which the indicator changes colour is
called the end point. There are a number of indicators available, and they need to be chosen
according to the equivalence point.
Methyl orange (3.1 – 4.4 pH) Strong acid with weak base
Bromothymol blue (6.2 – 7.6 pH) Strong acid with strong base
Litmus (5.0 – 8.0 pH) Weak acid with weak base
Phenolphthalein (8.3 – 10.0 pH) Strong base with weak acid
The equivalence point of acid-base titration can be found by measuring electrical
conductance of the solution
A standard solution is a solution the concentration of which is accurately known, it is
prepared by primary standard, which has the following characteristics:
Available at high purity, high molecular mass, unreactive with components of air such as
water vapour and carbon dioxide.
1. Wright out appropriate mass of the primary standard into a beaker
2. Add demineralised water until completely dissolved
3. Carefully transfer solution into a volumetric flask, and rinse beaker thoroughly
4. Fill volumetric flask with demineralised water to the mark
5. Stopper flask and invert twenty times to produce homogenous solution
Titration
Beakers, pipette and burette should be rinsed with the solution it is to deliver
Conical flask should be rinsed with demineralised (distilled) water
1. Fill a burette with a solution and adjust the solution level in the burette to the zero mark
(ensure no air bubbles)
2. Using the pipette, place a solution of the other sample in a flask under the burette
3. Add two to four drops of a suitable indicator to the flask (indicator is slightly
acidic/basic, and thus too much shouldn’t be added)
4. Place a piece of white tile under the flask
5. Run the solution from the burette into the flask with constant swirling rapidly until the
end point is close, then run the solution drop by drop carefully until the indictor just
changes colour
6. Read the volume (use white background and avoid parallax error)
7. Calculate the required concentration or amount

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(base) and
its conjugate base(acid) and which is therefore able to maintain an approximately constant
pH even when significant amounts of strong acid or strong bases are added to it.
CO2 + H2O  H+ + HCO3This is a buffer system in lakes whereby acids and bases are neutralised. When acid is added,
carbonic acid is produced, increasing H+ concentration and carbon dioxide escapes into the
atmosphere, neutralising the solution due to Le Chatelier’s principle. When hydroxide ion is
added, H+ concentration reduced and the reaction shifts to the right to oppose the change.
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
Functional group: an atom, group or group of atoms that reacts in a characteristic way in
different carbon compounds
Alkanol: -OH hydroxyl group
Alkanoic acid: -COOH carboxylic acid group
Because of the extra double bonded Oxygen atom alkanoic acid is more polar than alkanol

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
Methanol
CH3OH
Methanoic
HCOOH
Ethanol
C2H5OH
Ethanoic
CH3COOH

Propanol
C3H7OH
Propanoic
C2H5COOH
Butanol
C4H9OH
Butanoic
C3H7COOH
Pentanol
C5H11OH
Pentanoic
C4H9COOH
Hexanol
C6H13OH
Hexanoic
C5H11COOH
Heptanol
C7H15OH
Heptanoic
C6H13COOH
Octanol
C8H17OH
Octanoic
C7H15COOH
explain the difference in melting point and boiling point caused by
straight-chained alkanoic acid and straight-chained primary alkanol structures
Alkanoic acid because of the addition polar C=O bond is more polar than alkanoics,
therefore from stronger dipole-dipole bonds and hence stronger intermolecular force.
Alkanoic acid because of the oxygen also has a higher molecular weight, therefore it has
higher melting point and boiling point.

identify esterification as the reaction between an acid and an alkanol and describe,
using equations, examples of esterification
Esterification is the reaction between an acid and an alkanol. Alkanols react with alkanoic
acids in a acid-catalysed reaction to form esters. Esters are molecules of strong aroma.

describe the purpose of using acid in esterification for catalysis
Usually concentrated sulphuric acid is used as a catalyst (18M). It is a dehydrating agent
and is used to lower the activation energy, increase the rate of reaction and achieve
equilibrium at a much faster rate.

explain the need for refluxing during esterification
Refluxing is the process of heating a reaction mixture in a vessel with a cooling condenser
attached in order to prevent loss of any volatile reactant or product. Esterification is a
moderately slow process at room temperature and does not go to completion. Refluxing
allows the reaction to be carried out at a higher temperature (near the boiling point of the
alcohol) than would otherwise possible. The water-cooled condenser condenses volatile
vapour to prevent the loss of both reactant and product, increases the time that reagents are
together and allows equilibrium to be reached faster. It also prevents the possibility of an
explosion if the reaction is carried out in a closed vessel.

outline some examples of the occurrence, production and uses of esters
Esters have pleasant, fruity odours and occur widely in nature as perfumes and flavouring
agents.
E.g. apricot: methyl butanoate
orange: octyl ethanoate
Esters are a substantial industry for developing and manufacturing synthetic flavours and
perfumes. The first step is often to identify the constituents of the natural flavour and then to
synthesise similar mixtures of esters which reproduce this flavour. Such artificial flavours
are often cheaper than natural extracts, and provide they contain only substances that occur
in the natural flavours, the represent little health hazard.
Ethyl ethanoate is widely used as a solvent in industry. It is also the common solvent in nail
polish remover. High molecular weight (non-volatile) esters such as dialkyl phthalates are
used as plasticisers in some plastics such as PVC; they make the materials soft and pliable.
Esters are used as artificial perfumes or scents as they emit a sweet smell
Esters are used in making artificial food flavours that are added in many edible items like
ice creams, soft drinks, sweets, etc
Esters are used as industrial solvents for making cellulose, fats, paints and varnishes
Esters are used as solvents in pharmaceutical industries
Esters are used as softeners in plastic industries and molding industries

process information from secondary sources to identify and describe the uses of
esters as flavours and perfumes in processed foods and cosmetics
If a low molecular weight ester is encountered, the odour is usually pleasant. They are the
very popular artificial scents and flavours added to a wide array of consumer products such
as soft drinks, chewing gum, scented wax candles, food flavourings, etc.
In many cases, although not exclusively so, the characteristic flavours and fragrances of
flowers and fruits are due to compounds with the ester functional group. An exception is the
case of the essential oils. The organoleptic qualities (odours and flavours) of fruits and
flowers are often due to a complex mixture in which a single ester predominates. Food and
beverage manufactures are thoroughly familiar with these esters and often use them as
additives to spruce up the flavour or odour of a dessert or a beverage. Many times such
flavours or odours do not even have a natural basis, as is the case with the juicy fruit
principle, isopentenyl acetate. An instant pudding that has the flavour of rum may never
have seen its alcoholic namesake – this flavour can be duplicated by the proper admixture,
along with other minor components, of ethyl formate and isobutyl propionate.
A single compound is rarely used in good-quality imitation flavouring agents. A formula for
an imitation pineapple flavour includes ten esters and carboxylic acids that can easily be
synthesised in the laboratory. The remaining seven oils are isolated from natural sources.
Although the fruity tastes and odours of esters are pleasant, they are seldom used in
perfumes or scents that are applied to the body. The reason for this is chemical. The ester
group is not as stable to perspiration as the ingredients of the more expensive essential-oil
perfumes. The latter are usually hydrocarbons (terpenes), ketones, and ethers extracted from
natural sources. Esters, however, are used only for the cheapest toilet waters, since on
contact with sweat, they undergo hydrolysis, giving organic acids. These acids, unlike their
precursor esters, generally do not have a pleasant odour. Butyric acid, for instance, has a
strong odour like that of rancid butter (of which it is an ingredient) and is a component of
what we normally call body odour. It is this substance that makes foul-smelling humans so
easy for an animal to detect when he is downwind of them. It is also of great help to the
bloodhound, which is trained to follow small traces of this odour. Ethyl butyrate and methyl
butyrate, however, which are the esters of butyric acid, smell like pineapple and apple,
respectively.
A sweet fruity odour also has the disadvantage of possibly attracting fruit flies and other
insects in search of food. Isoamyl acetate, the familiar solvent called banana oil, is
particularly interesting.
Fourteen