HSC Chemistry Revision Module 2
David Pham
#
1.
The Acidic Environment
classify common
substances as acidic,
basic or neutral
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
• Skill
Examples of common substances and their acidity:
Acidic
Neutral
Basic
Citric juices
Salt
Oven cleaner
Carbonated drinks
Sugar
Ammonia
Milk
Pure water
Detergent
Aspirin
Alcohol–water solutions Blood
Vinegar
Lactose solutions
Soda (Bicarb, washing)
Battery acid
Drain cleaner
Stomach acid
Soap
Lactic acid
Lime water
Indicators are strongly coloured substances which, in solution, changes colour
depending on the solution’s acidity, explored more fully below. Examples 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. They are particularly useful as
their indicator ranges cover the common titration ranges and real–life acidities.
• Litmus: Red
Blue, [4.8 – 8.1]
• Phenolphthalein: Colourless
Pink, [8.0 – 10.0]
Yellow, [3.1 – 4.4]
• Methyl Orange: Red
Blue, [6.0 – 7.6]
• Bromothymol Blue: Yellow
Acid–base indicators are weak organic acids (denoted as HIn, and are usually
pigments or dyes) that has a different colour to its conjugate (In–), with the
colour change occurring over a specific and relatively narrow pH change.
Typically, one or both of the forms are intensely coloured, so only a tiny amount
of indicator is needed, little enough not to disrupt the pH of the solution being
studied. Indicators change colour due to the equilibrium reactions that they set
up with the amount of H+ in solution.
HIn(aq) +H2O(l)
H3O+(aq) +In–(aq)
As ka is constant, the ratio of [HIn] to [In–] changes with the [H3O+]. The
experimenter typically sees the HIn colour if the ratio of [HIn]/[In–] is 10:1 and
the In– colour if the ratio is 1:10. Between these, the colours merge into some
intermediate hue, such as a purple to litmus at pH 6.5. The transition range may
shift slightly depending on the concentration of the indicator in solution and on
the temperature at which it is used.
Because of the subjective determination of color, pH indicators are susceptible
to imprecise readings. For applications requiring precise measurement of pH, a
pH meter is frequently used.
Some colour changes of other common indicators are:
Indicator
Colour Transition
Transition range
Thymol blue (first transition)
red
yellow
1.2–2.8
Methyl yellow
red
yellow
2.9–4.0
Bromophenol blue
yellow
purple
3.0–4.6
Bromocresol green
yellow
blue-green
3.8–5.4
Phenol red
yellow
red
6.8–8.4
Cresol Red
yellow
reddish-purple
7.2–8.8
Thymol blue (second transition) yellow
blue
8.0–9.6
Alizarine Yellow R
yellow
red
10.2–12.0
identify and describe some
everyday uses of
indicators including the
testing of soil
acidity/basicity
*
*
*
2.
solve problems by applying
information about the
colour changes of
indicators to classify some
household substances as
acidic, neutral or basic
perform a first–hand
investigation to prepare
and test a natural indicator
identify oxides of non–
metals which act as acids
and describe the
conditions under which
they act as acids
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
Indicators are used as solutions or can be adsorbed onto the surface of paper
such as filter paper, especially with litmus paper. Indicators are used at times in
everyday life, most often when dealing with the suitability of a substance for a
use. They are used to detect whether a substance is basic, acidic, or neutral.
For example:
• Domestic waste water and waste water from light industries is often tested to
ensure that waste water is not acidic so that it will not corrode sinks, drains
and sewerage pipes.
• In chemical research they are used to determine the acidity or basicity of a
solution, and to monitor changes in acidity during accurate volumetric
analysis (titrations).
Soil acidity is tested to ensure that the plants intended for the area have suitable
pH in which to live. Some plants have adverse effects when the soil is too acidic
or too basic, and plants intrinsically prefer an environment to the other. This
unsuitable pH can be corrected with suitable addition of calcium oxide (lime) to
increase pH or ammonium sulfate or compost to decrease it. Soil pH can be
measured electronically but can also be measured with universal or narrow
indicators, such as methyl orange or bromothymol blue.
To test, fill a test tube 1/3 with soil, add distilled water, stopper it, and agitate.
The supernatant water is removed with a Pasteur pipette into two test tubes of
universal indicator (to get an imprecise pH range) then a narrow–range indicator
to obtain an accurate pH.
Swimming pools are also monitored for their pH levels for their suitability for
human use, and this can be modified as necessary. The addition of NaOCl
produces HOCl which is used to kill microbes, but increases pH. HCl or NaOCl
is added as necessary to control pool pH levels, as a pH under 6.5 causes
metals to be attacked and pH over 8 can cause irritation to lungs. An ideal pH is
7.2 to 7.6.
Phenol red is used as an indicator, compared against standards, where low pH
(<6.8) is given by yellow, high (>8.4) given by red–purple, and satisfactory
levels given by pinkish orange. To test pool pH, collect a water sample, add
phenol red, and then compare with the colour chart to find the pH range.
• Skill
Indicators can be used to differentiate between the possible pH ranges of
substances. For example, litmus is a common indicator and its presence shows
whether the substance has pH less than 4.8 or higher than 8.1. Other indicators
can narrow this range, and thus indicators can be used to find the pH range of
a substance, classifying it as acidic, neutral or basic.
• See Attachments
• See Prac Book
Carbon dioxide (CO2), sulfur dioxide (SO2) and nitrogen dioxide (NO2) all
dissolve in water forming acid solutions. Most non–metal oxides (except for CO,
NO and N2O which are neutral) are said to be acidic in solution. They act as
acids upon reaction with water, creating H3O+ ions. Notably, the oxides of more
electronegative elements make strong acids. For example:
2HClO4
Cl2O7 + H2O
SO3 + H2O
H2SO4
N2O5 + H2O
2HNO3
P4O10 + H2O
4H3PO4
To detect that a non–metal oxide gas is acidic with indicator paper, the paper
must be moist. Moisture enables the gas to dissolve and form the acid that
produces hydrogen ions. Reaction of a hydrogen ion with an indicator causes
the colour change.
It can be identified that, in general, the more electronegative an element, the
more acidic its oxide is, when in contact with water. For example, the highly
electronegative S, N, and Cl form very strong (~100% ionisation) acids in water.
That is why non–metals form more acidic oxides. More weak semi–metals form
weaker acids, notably HSiO3, analogous to silicon dioxide dissolved in water –
but note that SiO2 doesn’t dissolved in water readily. This is because the weaker
acid, with their weaker A– group, does not release the hydrogen as easily. For
example:
SO3(g) + 2NaOH(aq)
Na2SO4(aq) + H2O(l)
On the other hand, metals form basic oxides. These have very little
electronegativity and as such form ionic solids with oxygen. These ionise
completely in water to form oxide ions, which react with water to form hydroxide.
The more ‘electropositive’ a metal is, the more it ionises in water to form a
strong base. Some metal oxides cannot be dissolved in water fully but that part
that does completely ionises if the metal is sufficiently ‘electropositive,’ forming
hydroxide ions. For example:
O2– + H2O
2OH–
CuO(s) + 2HCl(aq)
CuCl2(aq) + H2O(l)
Therefore, oxides of elements to the upper-right in the periodic table are most
acidic, while those in the bottom left are most basic – that is, the more metallic
an element, the more basic its oxide will be.
There are some elements which form amphoteric oxides, which can act as
either an acid or base, depending on the conditions of reaction. For example:
2AlCl3(aq) + 3H2O(l)
Al2O3(s) + 6HCl(aq)
where aluminium oxide acts as a base, or here, where it acts as an acid:
2NaAl(OH)4
Al2O3(s) + 2NaOH(aq) + 3H2O(l)
Many transition metals form amphoteric oxides, such as zinc, vanadium,
chromium, and manganese. For example:
K2[Zn(OH)2](aq)
ZnO(s) + 2KOH(aq)
ZnO(s) + 2HCl(aq)
ZnCl2(aq) + H2O(l)
Here, zinc oxide can be seen to be able to neutralise strong acids and bases.
Also, metals with higher oxidation states tend to form oxides which are more
acidic. For example, Cr oxides can be basic(+2), amphoteric(+3), or acidic(+6).
describe, using equations,
examples of chemical
reactions which release
sulfur dioxide and chemical
reactions which release
oxides of nitrogen
Sulfur dioxide:
Action of bacteria and subsequent oxidation
2SO2(g) +2H2O(g)
2H2S(g) + 3O2(g)
Iron/Copper/Zinc Smelting
4FeS2(s) + 11O2(g)
2Fe2O3(s) + 8SO2(g)
2CuFeS2(s) + 5O2(g) + 2SiO2(s)
2Cu(l) + 4SO2(g) + 2FeSiO3(l)
2ZnS(s) + 3O2(g)
2ZnO(s) + 2SO2(g)
Oxides of Nitrogen:
Lighting strikes/areas of high temperature
2NO(g)
N2(g) + O2(g)
N2O(g) + NO2(g)
3NO(g)
2NO(g) + O2(g)
2NO2(g)
O2N(g) + NO2(g)
N2O4(g)
NO2(g)
NO(g) + O(g)
*
identify natural and
industrial sources of sulfur
dioxide and oxides of
nitrogen
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
assess the evidence which
indicates increases in
atmospheric concentration
of oxides of sulfur and
nitrogen
explain the formation and
effects of acid rain
Sulfur dioxide is formed by either combustion of fossil fuels with sulphur
impurities, or via the processes of roasting and extraction of copper from
chalcopyrite and zinc from zinc sulphide (though this is siphoned off to make
sulfuric acid these days). Also, SO2 comes from oxidation of H2S, smoke from
bushfires and the major contributor to SO2 emission is volcanoes.
2SO2(g) +2H2O(g)
2H2S(g) + 3O2(g)
4FeS2(s) + 11O2(g)
2Fe2O3(s) + 8SO2(g)
2CuFeS2(s) + 5O2(g) + 2SiO2(s)
2Cu(l) + 4SO2(g) + 2FeSiO3(l)
2ZnS(s) + 3O2(g)
2ZnO(s) + 2SO2(g)
Nitrogen oxides are formed in car engines and other high–temperature
combustion environments (particularly high voltages) as these high temperatures
cause the nitrogen and oxygen in air to react. This also occurs at oil refineries
and electrical power production sites, where arcing is present. In Sydney, around
85% of total NOx emissions comes from vehicles, which combust impure fossil
fuels. It can also be produced via the nitrogen cycle and bacterial action.
2NO(g)
N2(g) + O2(g)
3NO(g)
N2O(g) + NO2(g)
2NO(g) + O2(g)
2NO2(g)
O2N(g) + NO2(g)
N2O4(g)
NO2(g)
NO(g) + O(g)
Sulfur and nitrogen oxides are detrimental to the environment for several
reasons. NO2 causes irritation of the eyes, particularly in young children and
older people. At higher concentrations it causes extensive tissue damage. It also
leads to ozone formation, as photochemical smog. Ozone can have adverse
effects also, such as irritation of respiratory systems. Sulfur oxides are irritating,
poisonous gases, particularly affecting asthmatics.
H2SO3(aq)
SO2(g) + H2O(l)
H2SO3(aq) + O2(g)
2H2SO4(aq)
SO3(g) + H2O(l)
H2SO4(aq)
2NO2(g) + H2O(l)
HNO3(aq) + HNO2(aq)
2HNO3(aq)
2HNO2(g) + O2(g)
4NO2(g) + 2H2O(l) + O2(g)
4HNO3(aq)
In addition, they can cause acid rain (rain with a higher [H+] than normal) through
dissolution in water, forming strong acids which can adversely affect the
environment and humans. Sulfur dioxide (and its further oxidised state, SO3),
react with water vapour to form sulphurous and sulfuric acids, which are weak
and strong respectively. NO2 reacts to form nitrous and nitric acid, which are of
similar strength. These chemicals are highly corrosive.
The release of sulfur and nitrogen oxides into the environment is detrimental to
not only to the environment and its complex interactions, but also to humans
and society. Thus, the emission of these acidic oxides into the atmosphere is
undesirable, and steps should be further taken to reduce their effect on the
atmosphere.
• See attachments (P131-133, 153 of Jacaranda Chemistry)
Normal rain is already slightly acidic due to the dissolved carbon dioxide
present, even in unpolluted rain, producing carbonic acid. It has a pH of around
5.6
H2CO3(aq)
CO2(aq) + H2O(l)
H2CO3(aq)
H+(aq) + HCO3–(aq)
Acid rain is rain which is more acidic than usual, and has a pH of lower than 5.
Acid rain is mainly the result of rain dissolving non–metal oxides present in the
atmosphere, particularly sulfur and nitrogen oxides. These produce acids
(sulfurous, sulfuric, and nitric acids, though dilute), and hence acid rain:
H2SO3(aq)
SO2(g) + H2O(l)
H2SO3(aq) + O2(g)
2H2SO4(aq)
SO3(g) + H2O(l)
H2SO4(aq)
2NO2(g) + H2O(l)
HNO3(aq) + HNO2(aq)
define Le Chatelier’s
principle
identify factors which can
affect the equilibrium in a
reversible reaction
2HNO2(g) + O2(g)
2HNO3(aq)
4NO2(g) + 2H2O(l) + O2(g)
4HNO3(aq)
Sulfurous and nitrous acids are weak but nitric and sulfuric acids are strong.
They can be formed upon oxidation of the weaker acidic oxides, dissolved in
water. These oxides, their sources, and recent rising levels are discussed
above. In areas of North America and Europe rainfall of pH 4 is common and in
some areas a pH of 2 has been detected. This is mainly due to the fact that they
are densely populated and hence highly industrialised regions, which produce
many of these gases.
Acid rain has varied and adverse effects on the environment and human
society. Firstly, it causes surface waters and lakes to become acidic, often with
drastic effects on aquatic ecology. Many thousands of lakes in industrialised
regions are now too acidic to support fish life. Acid rain disrupts the CO2
gaseous/aqueous equilibrium and stresses fish life. In addition, lowered pH
irritates external contacts, such as skin and gills. Also, aluminium which has
been leached from soils by the acidic rain is toxic to marine life. Biological
magnification of acid rain through the food chain can seriously damage
predators. Secondly, acid rain causes damage to plants, including crops and
forests, via their corrosive nature. Pine needles, for example, lose their waxy
coating, or acid rain can cause defoliation. This adverse effect nature is also
amplified due to the acid’s effect on soil, particularly as seedlings can become
damaged. Minerals such as K, Ca, and Mg can be removed from the ground by
dissolution via acid rain, removing plant nutrients and possible causing toxic
levels of metal ions. Acid rain affects structure made of stone or metal,
corroding them. In particular, the historical heritage in the world is slowly being
destroyed by the effects of acid rain on the marble (limestone) building and
statues. Organics and other substances are also affected.
Ca2+(aq) + CO2(g) + H2O(l)
CaCO3(s) + 2H+(aq)
Fe(s) + H2SO4(aq)
FeSO4(aq) + H2(g)
In 1885, the French chemist, Le Chatelier, put forward a principle for predicting
the effect of change on reversible reactions:
‘If a chemical system at equilibrium is disturbed, the system will adjust to re–
establish equilibrium in such a way to minimise the effect of the disturbance.’
A change in concentration:
• Removing A or B or adding C or D will increase the backward reaction rate as
there is less reactant, relative, to before the change.
• Removing C or D or adding A or B will increase the forward reaction rate as
there is less product, relative, to before the change.
However, in an equilibrium, changing the amount of liquids or solids will not
affect the equilibrium as their concentration remains unchanged.
Changing the concentration or amount of aqueous or gaseous reactant systems
will cause the equilibrium to shift in such a way that the ratio of all species is as
close as it can be to before the change. However, they can never reach levels
before the change.
A change in pressure (volume) for gaseous systems:
If we assume n > m, and A,B are gases
• Increasing pressure will increase the production of B as more B means less
A, lowering the total number of moles of gas as the volume is less.
• Decreasing pressure will decrease the production to B to create more
molecules to expand into the extra space.
• Volume changes can also be thought of pressure changes but note that
concentration changes as well.
• Adding an inert gas will not changed the equilibrium position.
If m > n, the direction of preferred reactions are the opposite.
Changing the pressure (or volume) of a gaseous system will cause the system
to shift in such a way that the number of particles per unit volume is as close as
it can be to before the change, but it cannot achieve previous levels, as with any
other equilibrium change.
A change in temperature:
describe the solubility of
carbon dioxide in water
under various conditions
as an equilibrium process
and explain in terms of Le
Chatelier’s principle
*
*
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 25oC
and 100kPa
calculate volumes of gases
given masses of some
If we assume the forward reaction is has positive ∆H, heat is absorbed in the
forward reaction
• Increasing temperature will favour the forward reaction, which increases heat
input
• Decreasing temperature will favour the reverse reaction, which increases
heat output
If ∆H is negative, heat is released in the forward reaction.
• Decreasing temperature will favour the forward reaction, which increases
heat output
• Increasing temperature will favour the reverse reaction, which increases heat
input
Changing the temperature for an equilibrium will result in the temperature
changed being partially counteracted by increasing or decreasing heat output as
necessary, so that the heat of the surroundings is as close as it can be to before
the change.
In answering equilibrium problems, the following format should be used as to
avoid any confusion or assumptions:
The system acts to minimise the disturbance/oppose the change (Le Chatelier’s
principle) – that is, [insert]. The reaction that does this is [insert]; thus this
direction of reaction is preferred. The equilibrium position thus shifts [insert].
Subsequently, answer any extra parts, such as qualitative concentrations.
Pressure: In a closed system such as a soft drink bottle, CO2(g) is in equilibrium
with CO2(aq).
CO2(g)
CO2(aq), ∆H<0 (1)
As the left side has much more volume relative to the right side, an increase in
pressure shifts the equilibrium position to favour the forward reaction. This is
done in bottling plants where pressurised conditions are used to supersaturate
the solution with carbon dioxide. When the bottle is opened, the pressure is
decreased and the reaction favours the evolution of gas, seen as bubbles, and
eventually the drink becomes flat.
Temperature: The dissolution of carbon dioxide in water is exothermic. If the
solution is warmed, the solution favours the endothermic reaction, which
evolves gas. Heating a carbonated solution will eventually result in flatness.
Cooling it will increase the solubility.
pH: The dissolved CO2 reacts with water to form a slightly acidic solution of
carbonic acid.
H2CO3(aq) (2)
CO2(aq) + H2O(l)
This carbonic acid ionises in water in two steps, the second one producing little
product
H2CO3(aq)
H+(aq) + HCO3–(aq) (3)
–
HCO3 (aq)
H+(aq) + CO32–(aq) (4)
This explains rain water’s slight acidity and soda pop’s tangy taste.
If an additional source of hydrogen ions is added, it drives both (4) and (3)
equilibria to the left and increases the amount of carbonic acid. The increased
H2CO3 shifts the equilibria in (2) and creates more CO2(aq). This then drives
equilibrium (1) to produce CO2(g). As such, an increase in [H+] will lead to less
CO2 being dissolved. Conversely, in basic conditions, CO2 dissolves more
readily as more H2CO3 is consumed in reaction.
• See Attachments
• See Prac Book
•
• Skill
Taught in Preliminary Course
3.
*
substances in reactions,
and calculate masses of
substances given gaseous
volumes, in reactions
involving gases at 0oC
and 100kPa or 25oC and
100kPa
define acids as proton
donors and describe the
ionisation of acids in water
gather and process
information from
secondary sources to write
ionic equations to
represent the ionisation of
acids
identify acids including
acetic (ethanoic), citric (2–
hydroxypropane–1,2,3–
tricarboxylic), hydrochloric
and sulfuric acid
At 0o, 100kPa, 1 mole of gas occupies 22.71L of volume (STP).
At 25o, 100kPa, 1 mole of gas occupies 24.79L of volume (SLC).
By Brønsted–Lowry’s definition of acids and bases, acids are proton (H+) donors
and bases are proton acceptors. In water, an acid HA donates a proton to
water, forming H3O+ and A–, and this process is called ionisation.
HA + H2O
H3O+ + A–
It can be alternatively written as
HA
H+ + A–
• Skill
All monoprotic (singly ionisable hydrogen) acids (to some extent) ionise via
HA
H+ + A–
for strong acids, and
HA
H+ + A–
for weaker acids. Examples are:
H+ + Cl–
and CH3COOH
H+ + CH3COO–
HCl
Polyprotic acids ionise in step, though with each step the readiness of ionisation
decreases drastically, e.g.
H+(aq) + HSO4–(aq)
H2SO4(aq)
–
HSO4 (aq)
H+(aq) + SO42–(aq)
• Skill
Some common acids are:
• Acetic (ethanoic) acid – CH3COOH. Acetic acid occurs naturally in the
decomposition of biological material, such as oxidation of wine alcohol, but
most acetic acid used by humanity is manufactured industrially. It is a weak
acid, ionising to a small degree at standard conditions only.
H+ + CH3COO–
CH3COOH
• Citric (2–hydroxypropane–1,2,3–tricarboxylic) acid. Citric acid occurs naturally
in citrus fruits such as oranges, limes, and lemons, but it is also manufactured
to be added to food as preservatives. It is a weak triprotic acid, ionising in 3
steps, all of which proceed little and are progressively weaker.
3H+(aq) + C3H5O(COO)33–(aq)
C3H5O(COOH)3(aq)
• Ascorbic acid, known commonly as Vitamin C, is necessary in prevention of
scurvy. It is found in fresh fruit and vegetables and forms part of their taste.
• Hydrochloric acid (HCl) is an important acid. It is found in human stomach
acid and aids digestion, by activating enzymes. It is a strong acid due to its
ease of ionisation, caused by its highly polar covalent bonding. It ionises, to
all intents and purposes, 100% at standard conditions.
HCl
H+ + Cl–
• Sulfuric (H2SO4) acid is produced industrially on a large scale. Most sulfuric
acid is manufactured, but it can also occur naturally. For example, most sulfur
dioxide released into the earth’s atmosphere is oxidised and dissolved in
water to form the sulfuric acid in acid rain. If the acid rain results from volcanic
eruption it could be regarded as natural, but if acid rain results from smelting
of sulfide ores, it could be regarded as manufactured. It is industrially
*
identify data, gather and
process information from
secondary sources to
identify examples of
naturally occurring acids
and bases and their
chemical composition
identify pH as –log10 [H+]
and explain that a change
in pH of 1 means a ten–
fold change in [H+]
describe the use of the pH
scale in comparing acids
and bases
*
process information from
secondary sources to
calculate pH of strong
acids given appropriate
hydrogen ion
concentrations
important, being involved in fertiliser, explosive, and petroleum production. It
is a strong acid, due to the sulfate group being very electronegative and thus
allowing easy ionisation.
H2SO4(aq)
H+(aq) + HSO4–(aq)
HSO4–(aq)
H+(aq) + SO42–(aq)
• Carbonic acid (dissolved CO2 in water) – H2CO3. This occurs naturally, in
seawater and rainwater, which causes normal rain to be slightly acidic. As it is
a relatively weak acid, it is used in our bodies as a buffer system, explored
later on more fully.
H2CO3(aq)
H+(aq) + HCO3–(aq)
–
HCO3 (aq)
H+(aq) + CO32–(aq)
The definitions of acids can be used to identify acids, as well as pH<7.
Some examples of naturally occurring bases and acids are:
Acid
Base
Stomach acid (HCl)
Ammonia (NH3)
Vinegar (CH3COOH)
Nicotine (C8H14N2)
Citric acid
Limestone (CaCO3)
Acidic rain (H2CO3 and others)
From this equation we can see that a change in one unit of pH means a tenfold
changed in the concentration of H+, as the scale is logarithmic. It was chosen
thus as acidity strengths are not proportional with concentration, as citric acid,
1000 times weaker than stomach acid, does not affect us 1000 times less. Thus
the logarithmic scale was chosen for convenience of calculations.
A p[something] means the negative base ten logarithm of that [something].
Remember to use log10, not ln or loge
The pH scale offers a logarithmic scale for the relative strengths of acids and
bases (their degree of ionisation). This is used rather than the molarity of
hydrogen ions as the use of indices is cumbersome. In pure water without any
dissolved gas, [H+] = [OH–] = 10–7 mol L–1 and so pH =7. In an acidic solution,
[H+] > 10–7 mol L–1 and pH < 7. In a basic solution, [H+] < 10–7 mol L–1 and pH >
7. The following table relates pH to the hydrogen ion concentration, [H+], and
provides examples of common aqueous solutions for each pH value. Similarly, a
pOH scale can be used to represent the [OH–] of a solution. Note that, at
standard conditions, pH + pOH = 14.
[OH–]
Aqueous solution example
pOH
pH
[H+]
0
–14
0
10 = 1
10
1 M hydrochloric acid
14
1
10–1
10–13
0.1 M hydrochloric acid
13
2
10–2
10–12
stomach acid, lemon juice
12
3
10–3
10–11
soda water, wine
11
4
10–4
10–10
tomato juice, beer
10
5
10–5
10–9
acid rain
9
–6
6
10
10–8
urine, milk
8
7
10–7
10–7
pure water without any dissolved gas 7
8
10–8
10–6
sea water
6
9
10–9
10–5
baking soda solution
5
10
10–10
10–4
detergent solution, soap
4
11
10–11
10–3
milk of magnesia
3
12
10–12
10–2
household ammonia, bleach
2
13
10–13
10–1
0.1 M sodium hydroxide
1
–14
0
14
10
10 =1
1 M sodium hydroxide, oven cleaner
0
• Skill
All strong acids ionise 100% in water for their first proton. Thus, the pH is
usually the same as their concentration. Answer to 2 decimal places unless
instructed otherwise. Remember to use log10, not ln or loge
For weaker acids, the degree of ionisation must be given, as use of ka is not
covered in the syllabus. Find [H+] and use the same above formula for pH.
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
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
For calculations involving a hydroxide concentration rather than H+, use the fact
that pH + pOH = pKw = 14.
Strong, concentrated, weak and dilute have different meanings to their
everyday use when referring to acidic solutions. They refer to degree of
ionisation and concentration, rather than strength.
• A strong solution is one in which a high degree (~100%) of the acid in the
solution is ionised (but not necessarily concentrated), e.g. hydrochloric acid.
• A concentrated solution is one in which there are a high amount of acid (but
not necessarily ionised) per unit volume, e.g. any solution with high molarity
• A weak solution is one in which a low degree of acid in solution is ionised,
e.g. vinegar.
• A dilute solution is one in which there are few acid molecules per unit volume,
even though they might be all ionised.
From this, it can be seen that a dilute strong solution can still have a lower pH
than a concentrated weak solution.
Note that there is no arbitrary cut–off between weak and strong or concentrated
and dilute. When asked to give quantitative values, one should use extreme,
indisputable concentrations or acid strengths, such as 0.001M or 5M, or 100%
dissociation versus 1% dissociation.
A strong and weak acid are differentiated by the degree of ionisation they
undergo in solution. The more H+ per molecule, the stronger the acid is.
For any given acid, an equilibrium is set with the surroundings
HA(aq)
H+(aq) + A–(aq)
Reiterating: the strength of an acid is dependant on this degree of ionisation, or,
in other words, the position between the intact molecule and its ions – [H+]/[HA].
A strong acid’s equilibrium is (almost, and for all intents and purposes is) shifted
all to the right, giving a very high ratio, while in a weak acid the forward reaction
proceeds very little, leading to a very low ratio.
For example;
HCl(aq)
H+(aq) + Cl–(aq)
CH3COOH(aq)
H+(aq) + CH3COO–(aq)
Hydrochloric acid thus produces much more H+ than acetic acid and thus is
stronger, having a lower pH. This is due to the fact that HCl has a very high
equilibrium ratio of [H+]/[HA] compared to acetic acid, which has a very low
position.
For the equation
HA(aq)
H+(aq) + A–(aq)
The magnitude of Ka shows how much the reaction shifts to the right hand side,
and depicts the relative strengths of acids.
Ka of acetic acid = 1.8 x 10–5
Ka of citric acid = 7.4 x 10–4 ; 1.7 x 10– 5; 4.0 x 10–7 (citric acid is triprotic)
Ka of hydrochloric acid = large
From this we see that acetic acid is the weakest of the three, then citric acid,
and hydrochloric acid is by far the strongest. This is due to the degree of
ionisation of the molecules in water; less strong acids ionise less readily in
water and require a much higher concentration to equal that of a stronger acid.
Acetic acid and citric acid usually ionise to less than 5% (depending on their
concentration) but hydrochloric acid ionises, for all intents and purposes, 100%
*
*
*
plan and perform a first–
hand investigation to
measure the pH of
identical concentrations of
strong and weak acids
use available evidence to
model the molecular
nature of acids and
simulate the ionisation of
strong and weak acids
at standard conditions. Note that the 2nd and 3rd ionisations of citric acid do not
affect the pH significantly, as they are orders of magnitude less than the first.
For example, at 0.01M solutions:
Acid
Hydrochloric acid Citric acid
Ethanoic acid
Degree of ionisation
100%
27.5%
4.2%
pH
2
2.56
3.38
For 1M solutions:
Acid
Hydrochloric acid Citric acid
Ethanoic acid
Degree of ionisation
100%
2.75%
0.42%
pH
0
1.56
2.37
Thus it can be seen that HCl is stronger than citric acid, which is in turn stronger
than acetic acid, due to their degrees of ionisation.
• See Attachments
• See Prac Book
• See animation on Jacaranda CD
A molymod kit can be used to model strong/weak/concentrated/dilute acids
The hydrohalic acids HCl and HF were modelled with a molymod kit. HF is a
weak acid and ionises little in solution, while HCl has essentially complete
ionisation. H is a black circle, F a red circle, and Cl blue.
In their non-solution state they exists as molecular gases, as below.
In solution, however, they ionise to form H+ ions and their conjugate base.
*
*
*
solve problems and
perform a firsthand
investigation to use pH
meters/probes and
indicators to distinguish
between acidic, basic and
neutral chemicals
gather and process
information from
secondary sources to
explain the use of acids as
food additives
Note the differences between their degrees of ionisation – HCl, being a strong
acid, completely ionises, while HF, as a weak acid, ionises to a small extent
only. The evidence for this small extent of ionisation comes from testing pH
using a pH meter – HF’s pH is higher than HCl’s at the same concentration,
meaning that the HF solution has less H+ ions, meaning it ionises less.
• Skill
Acidic chemicals have a pH of less than 7 at 25o. Basic chemicals are those
with a pH of more than 7 at 25o, and neutral substances have a pH of 7 at 25o.
• See Attachments
• See Prac Book
Many acids, biological and otherwise, can be used as food additives for a wide
range of functions.
• They can be added for taste in food or drinks, especially malic, acetic, citric
and tartaric acids. These add a tart (sharp) taste to foods. Phosphoric acid
acidulates colas.
• They inhibit growth of microbes such as bacteria and mould, due to the low
pH preventing excess enzyme action. Propanoic acid is used for bread,
potato crisps, and cake mixes.
• Acting as antioxidants, they prevent spoilage of foods in use as
preservatives. Citric acid is used as a preservative in soft drinks, and vitamin
C is also used.
• Acetic acid (4% solution) is used as vinegar to preserve foods, such as
pickling. Tartaric acid is used in jams, fruits, pickles, and soft drinks.
• They act as leavening agents (substances which react with NaHCO3 to
produce CO2 gas). Tartaric acid acts as one in desserts.
Acidic Food Additive
Chemical Formula
Information
Acetic acid
CH3COOH
used as vinegar (4% solution) to preserve
foods (e.g. pickling); flavour enhancer
Citric acid
HOOCCH2COH(COOH)CH2COOH flavouring and preservative (anti-oxidant),
especially in soft drinks; antacid ingredient
Malic acid
HOOCCH2CHOHCOOH
flavour enhancer particularly in fruit fillings in
bakery products; improves aftertaste; boosts
savoury tastes; preferred acidulant in noncarbonated drinks to provide sour taste; used in
diet drinks and diet candy to reduce the intense
sweetness of the artificial sweeteners.
Tartaric acid
HOOCCHOHCHOHCOOH
antioxidant and flavouring; preservative in
jams, fruits, pickles and soft drinks; emulsifying
agent in bread making; leavening agents in
desserts
Lactic acid
CH3CHOHCOOH
production of dairy products such as cheese
and yoghurt; acidity regulator
Phosphoric acid
H3PO4
acidulation of soft drinks (particularly colas);
manufacture of cottage and processed cheese;
pH control in diet jellies
Propanoic acid
CH3CH2COOH
controls bacteria and mould growth, particularly
in bread, potato crisps and cake mixes
Ascorbic acid (vitamin C)
CH2OHCHOH (C4H3O4)
antioxidant to prevent spoilage; added to
increase vitamin C in many foods
Lavoisier proposed that acids contained oxygen, and gave oxygen its name,
4. outline the historical
development of ideas
which means ‘to give rise to acids,’ as his knowledge was mostly restricted to
about acids including those oxyacids. As a result of his studies, Lavoisier showed that many non–metal
of:
oxides, when dissolved in water, formed acids. He eventually went on to name
oxygen, from ‘acid–former.’ Although the oxides of many non–metals such as
• Lavoisier
SO2 and CO2 form acids in water, metal oxides in fact form bases.
• Davy
Humphry Davy, in 1810, was able to show that the hydrohalic acids, H2S, and
• Arrhenius
H
2Te did not contain oxygen. He then proposed that all acids contained
* gather and process
hydrogen
rather than oxygen, and by 1830 several other acids which did not
information from
contain
oxygen
were discovered, such as HBr, HF, HI, and HCN. However, this
secondary sources to trace
still
did
not
explain
the characteristic properties of acids. Also, this theory did not
developments in
explain
why
some
compounds
containing hydrogen, such as methane, are not
understanding and
acidic, but could even be basic, as in the case of ammonia.
describing acid/base
In 1838, German chemist Justus von Liebig extended Davy’s theory by
reactions
proposing that acids had ‘replacable hydrogen.’ He reasoned when acids
attacked, metals the metals replace the hydrogen to form a salt and hydrogen
gas. However, this still failed to account for some other properties of acids, such
as production of NO2 and hydrogen from nitric acid on metal.
Svante Arrhenius developed the theory which explained acids’ properties, by
stating that “acids form hydrogen ions in aqueous solution, while bases form
hydroxide ions”. Also, acids neutralised bases and vice versa. The hydrogen
ions thus cause the acidic properties. For example:
H+(aq) + Cl–(aq)
HCl(aq)
This explained the fact that acids can conduct electricity, and the differing
strengths of acids, which was explained by degree to which the forward reaction
occurs. For example, CH3COOH only produces a small amount of hydrogen
ions, and is this weaker than HCl.
H+(aq) + CH3COO–(aq)
CH3COOH(aq)
Arrhenius also suggested that reactions between acids and bases, called
neutralisation, produced water when H+ and OH– react.
H+(aq) + OH–(aq)
H2O(l)
In water, hydrogen ions, due to their small size and high charge (and status as a
single proton), react with water molecules to form H3O+, a coordinate covalent
compound. They are represented thus as H+(aq) or H3O+(aq) even though the latter
is more correct.
Arrhenius also proposed that a base was a substance that, upon dissolution in
water, produced OH– ions. For example:
NaOH(aq)
Na+(aq) + OH–(aq)
Ba+(aq) + 2OH–(aq)
Ba(OH)2(s)
Although the Arrhenius definition is suitable for many common acids and bases,
it has its limitations. Many substances which behave as acids or bases, such as
NH3 and Na2CO3, do not contain an OH group, and their reactions with acids
could not be explained. Some substances which react with acids and have OH
in their structure, but are insoluble in water (e.g. some group two hydroxides).
These couldn’t be classified as bases, as Arrhenius’ theory only applied to
aqueous media. It also does not fully explain or account for the relative
strengths of acids and bases (i.e. why they ionise to different levels). Also, it
could not explain why certain salts, when dissolved in water, created acidic or
basic solutions, such as NaS or ZnCl while NaCl is neutral.
The Brønsted–Lowry theory of acids and bases addressed these shortcomings.
Further development led to the Lewis (based on electrons) theory which is more
generalised, and incorporates the solvent–system definition.
outline the Brønsted–
Lowry theory of acids and
bases
describe the relationship
between an acid and its
conjugate base and a base
and its conjugate acid
Brønsted and Lowry independently developed a more general theory of acids
and bases, which involves proton transfer and acceptance in acid–base
reactions, and addressed shortcomings of previous theories. An acid is a proton
donor (H+) while a base is a proton acceptor, for example
HCl(g) + H2O(l)
H3O+(aq) + Cl–(aq)
In this reaction HCl is donating a proton to H2O, which means that HCl is acting
as an acid and H2O is acting as a base. Similarly
NH4+(aq) + OH–(aq)
NH3(g) + H2O(l)
Here ammonia acts as a base by accepting H+ while water, which donates the
proton, acts as an acid. Here, we can see that while ammonia contains no OH–
in its structure, it creates it upon reaction. Thus we can see that the Brønsted–
Lowry theory can adequately explain many more observed reactions and
properties of acids and bases, though the Arrhenius definition works quite well
for other applications.
Acids are defined as proton donors, while bases are proton acceptors. This
theory also added a role to the solvent, and mainly focuses on water as an
ionising solvent.
H3O+(aq) + OH–(aq)
H2O(l) + H2O(l)
However, it has also allowed for chemists to venture out of aqueous chemistry
to the realm of non–aqueous or gas–phase reactions.
In the Brønsted–Lowry theory, a protonated base has the potential to act as an
acid, and, similarly, a deprotonated acid is a potential base. For example,
ammonia’s reaction can be written backwards:
NH4+(aq) + OH–(aq)
NH3(g) + H2O(l)
–
+
NH3(g) + H2O(l)
NH4 (aq) + OH (aq)
+
In this reaction NH4 is now donating a proton to water and is thus an acid, while
OH– is the accepting base. Also,
CH3COOH(aq) + OH–(aq)
H2O(l) + CH3COO–(aq)
–
+
CH3COO (aq) + H3O (aq)
CH3COOH(aq) + H2O(l)
These reactions show that ethanoic acid acts as a weak acid while the
identify conjugate
acid/base pairs
identify a range of salts
which form acidic, basic or
neutral solutions and
explain their acidic, neutral
or basic nature
ethanoate ion acts as a weak base. Thus, CH3COOH/CH3COO– and NH4+/NH3
are conjugate acid/base pairs. In general, with acid HA,
H+(aq) + A–(aq)
HA(aq)
–
HA is the acid, and A is its conjugate base. The stronger a particular acid, the
weaker its conjugate base. Similarly, the stronger a base, the weaker its
conjugate acid is.
The strength of an acid and its conjugate base are related. Strong acids such as
HCl and HNO3 have very weak conjugate bases, and conjugate acids of strong
bases such as the hydroxide ion (water) are very weak. This can be attributed to
the equilibrium set up between the acid/base and solvent:
HA(aq)
H+(aq) + A–(aq)
If HA is a strong acid, then the equilibrium position lies to the right–hand side,
then A– does not accept protons well, and is thus a weak base. Thus, A– is a
weak acid if HA is a strong acid.
H+(aq) + A–(aq)
HA(aq)
If A– is a strong base, then the equilibrium lies to the right–hand side as it
accepts protons well, meaning that HA does not donate them well. Thus, HA is
a weak acid if A– is a strong base.
If one is of intermediate strength then so is the other.
• Skill
Conjugate acid–base pairs are defined by the Bronsted–Lowry acid–base
theory. An acid is a substance which readily donates a proton (H+ ion) in order
to become its conjugate base. Similarly, bases accept protons to become the
conjugate acid. To change one to the other, simply add H+ (for base to acid) or
remove one (for acid to base). Some examples:
The reaction of a salt with water to produce a change in pH is called hydrolysis.
Any salt consists of two ions, a cation and an anion. Each of these, by the B–L
definition of acids and bases, has the capacity to act as an acid or base,
reacting with water. The pH of the solution depends on the degree to which
these ions act as B–L acids or bases.
Neutral solutions: A salt consisting of the anion of a strong acid and the cation
of a strong base yields a neutral solution because the ions do not react
appreciably with water. The anions of strong acids are all halide ions (except F–)
and strong oxoanions such as NO3– and ClO4–. The cations of strong bases are
those from Group 1A(1) and Ca2+, Sr2+, and Ba2+ from group 2A(2). Salt
containing only these ions, such as NaCl, Ba(NO3)2, and KBr yield neutral
solutions because no reaction with water takes place.
Acidic solutions: There are 3 main groups of ions which produce acidic
solutions. A salt consisting of the anion of a strong acid and the cation of a weak
base yields an acidic solution because the cation acts as a weak acid, and the
anion does not react. For example, NH4Cl produces an acidic solution because
*
*
choose equipment and
perform a first–hand
investigation to identify
the pH of a range of salt
solutions
identify amphiprotic
substances and construct
equations to describe their
behaviour in acidic and
basic solutions
the NH4+ ion is a weak acid, forming NH3, and the Cl– ion does not react as it is
the anion of a strong acid.
NH3(aq) + H3O+(aq)
NH4+(aq) + H2O(l)
Other examples include NH4NO3 and CH3NH3Br.
Small, highly charged metal ions also yield H3O+ in solution. For example,
Fe(NO3)3 produces an acidic solution because the hydrated Fe3+ ion acts as a
weak acid whereas the NO3– ion does not react.
Fe(NO3)3(aq) + 6H2O(l)
Fe(H2O)63+(aq) + 3NO3–(aq)
3+
Fe(H2O)6 (aq) + H2O(l)
Fe(H2O)5OH2+(aq) + H3O+(aq)
Other examples include CrCl3, FeBr3, ZnSO4, Cu(H2O)42+, and Al(NO3)3.
A third group of salts which yield H3O+ in solution consists of cations of strong
bases and anions of polyprotic acids that still have 1 or more ionisable protons.
For example NaH2PO4 yields an acidic solution as Na+ does not react while
H2PO4– is a weak acid. Note that HCO3– and HPO42– create basic salts.
H2PO4–(aq) + H2O(l)
HPO42–(aq) + H3O+(aq)
Other examples include KHSO4 and NaHSO3. Note that HCO3– and HPO42–
create basic salts.
Basic solutions: A salt consisting of the anion of a weak acid and the cation of a
strong base yields a basic solution in water as the anion acts as a weak base.
The anion of a weak acid accepts a proton from water to yield an OH– ion.
Sodium ethanoate, for example, yields a basic solution as Na+ does not react
with water but the CH3COO– ion will act as a weak base:
CH3COO–(aq) + H3O+(aq)
CH3COOH(aq) + H2O(l)
Other examples include KF and Na2CO3.
HCO3–(aq) + OH–(aq)
CO32–(aq) + H2O(l)
The hydrogen carbonate ion is a stronger acid than water so the equilibrium lies
to the left, producing hydroxide ions, which creates a basic solution.
Unknown: Salts of weakly acidic cations and weakly basic anions yield different
conditions, depending on their relative strength, and degree of reaction with
water. For example, NH4HS is basic as the production of OH– is greater than
that of H3O+, as NH4+ is a weaker acid that HS– as a base [kb(HS–) > ka(NH4+)].
However, NH4CH3COO is virtually neutral as both their acidic and basic qualities
are nearly the same. We can look at table of acid strengths.
Approximate pH for 0.01M solutions at standard conditions:
Acid
pH
Base
pH
HCl
1.0
NaOH
13
NaHSO4
1.4
Na3PO4
11.7
H2SO3
1.5
Na2CO3
11.5
H3PO4
1.5
NH3
11.1
HF
2.1
Na2SO3
9.8
CH3COOH
3.0
Na2HPO4
9.2
NaH2PO4
4.5
NaHCO3
8.4
NH4Cl
4.6
NaCH3COO
8.4
• See Attachments
• See Prac Book
In chemistry, a substance is described as amphiprotic if it can either (and both)
donate or accept a proton, thus acting either like an acid or a base (according to
Brønsted–Lowry theory of acids and bases: acids are proton donors and bases
are proton acceptors). They donate or accept protons depending on the
conditions of their reaction (usually the acidity of other substances in the medium
determines the nature of proton movement.)
Also, since they can act like an acid or a base, they are amphoteric. Amphoteric
substances, however, are not necessarily amphiprotic.
Water can react as a base in reaction with HCl and as an acid in reaction with
NH3. In autoionisation of water, one molecule donates a proton and the other
accepts it.
OH–(aq) + H3O+(aq)
H2O(l) + H2O(l)
–
H2O(l) + HS (aq)
OH–(aq) + H2S(aq)
identify neutralisation as a
proton transfer reaction
which is exothermic
*
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
qualitatively describe the
effect of buffers with
reference to a specific
example in a natural
system
H2O(l) + NH4+(aq)
NH3(aq) + H3O+(aq)
In the second and third reactions water is acting primarily as either a proton
donor or acceptor, depending on its conditions.
In similar ways the HCO3– ion (hydrogen carbonate ion) can react with water or
hydroxy or hydronium ions, as a base or acid, donating or accepting protons.
H2CO3(aq) + OH–(aq)
HCO3–(aq) + H2O(l)
–
HCO3 (aq) + H2O(l)
CO32–(aq) + H3O+(aq)
+
–
HCO3 (aq) + H3O (aq)
H2CO3(aq) + H2O(l)
HCO3–(aq) + OH–(aq)
CO32–(aq) + H2O(l)
2–
–
Similarly HPO4 and HSO4 are amphiprotic.
When an acid and a base meet, they undergo neutralisation. The meaning of
acid–base reactions have changed along with the definitions of acid and base,
but in the Arrhenius sense, neutralisation occurs when the H+ ion from the acid
and the OH– ion from the base combine to form H2O.
HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
For a net ionic equation (holds true for any monoprotic neutralisation reaction):
H2O(l)
∆rxnHo = –55 to 60kJ/mol
H+(aq) + OH–(aq)
It was observed that all neutralisation between strong acids and base had very
similar heats of reaction. This theory explained that the only reaction occurring
was that of neutralisation, producing the same amount of heat each time.
In the new Brønsted–Lowry definition, neutralization is now a proton–transfer
reaction, explaining its exothermic nature.
Examples of these similar values are:
∆rxnHo = –57.9 kJ/mol
HCl + NaOH
∆rxnHo = –57.3 kJ/mol
HNO3 + KOH
CH3COOH + NaOH
∆rxnHo = –56.1 kJ/mol
Strong acids and bases are a liability to safety in the laboratory and the
commercial world. Neutralisation reactions can be used to clean up after acids
or bases which have been accidentally spilled.
In a school setting, the immediate action to take is to make sure that the acid is
not on any person. If so, then immediate flushing is required and first aid is to be
administered. Isolation of the spill is required next; this is usually done with sand
or vermiculite, to prevent the liquid from spilling out and contaminating other
areas. It is recommended that strong acids are diluted before being mixed with
a base, as neutralisation is a very exothermic process and much heat is
released. Neutralisation is often carried out with Na2CO3 or NaHCO3. This is
due to the fact that they release CO2, and the fizzing allows us to monitor
reaction progress.
Larger spills use a similar process; spills are prevented from escaping by sand
or vermiculite, which is removed for off-site neutralisation. For alkaline
materials, NaHCO3 or Na2HPO4 can be used to neutralise the solution as they
are amphiprotic, and react with alkalis as well as acids and neutralise them.
Copious amounts of base and water are then used to minimise any pH changes
which may occur with excess ions, from being discharged.
Neutralisation is an exothermic reaction which can be dangerous, due to the
amount of heat evolved upon large-scale neutralisations. However it is
otherwise a safe and elegant technique to neutralise acid and base spills, and
as such is a suitable technique to an otherwise difficult problem.
Buffers are solutions which resists changes in pH when small quantities of acid
or base are added to them. By restricting the pH range of organic body fluids,
these buffers ensure that biochemical reactions are produced at their required
(optimum) rate. For example, our saliva pH must be maintained at 6.4 to 7.0,
pancreas cells at 8.5, and stomach acid at 1.6, or otherwise enzymes and juices
will not function properly. Buffers are also used in a chemical laboratory to
calibrate pH meters, and providing an environment for chemical processes.
Buffer solutions usually contain weak B-L acids and its conjugate base or a
weak B-L base and its conjugate acid, in correct amounts. This precise control
of concentrations is allows the pH of the solution to be kept within limits, as
described by Le Chatelier’s principle. Consider the weak acid HA:
A-(aq) + H3O+(aq)
HA(aq) + H2O(l)
This equilibrium lies well to the left due to the weakness of the acid. Its
conjugate base, A- is added (e.g. as NaA salt) then this equilibrium is
established.
HA(aq) + OH-(aq)
A-(aq) + H2O(l)
There are now two equilibria. Addition of either strong acid or base (in moderate
amounts) will not shift the pH excessively. For example, if we add OH-, eq2 will
shift to the shift to the left and minimise the increase in pH. Addition of a strong
acid, which produces H3O+, would cause eq1 to shift to the left, minimising the
acidity increase. Also, the A- and HA can react with the H3O+ and OH- to
neutralise the substances, converting them to much weaker bases. As such,
addition of strong acids or bases in moderate amounts will not modify the pH
excessively, but it is controlled. This can especially be seen in the weak-acid
strong-base titration curve (and the resultant flipped weak-base strong acid
titration curve) where there are areas of ‘flatness’ in which addition of moderate
amounts of chemical do not significantly change the pH. For example:
*
*
describe the correct
technique for conducting
titrations and preparation
of standard solutions
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
Human blood has a pH of about 7.4. A condition known as acidosis develops if
the pH of blood falls below 7.35. Below 7.0 the person will enter a coma, and if
the pH drops below 6.8, death may result. Similarly, if the pH rises above 7.45 a
condition known as alkalosis occurs, and above 7.8 this condition is life
threatening. The presence of buffers in the blood maintains the pH between 7.35
and 7.45. The main buffer system used to control the pH of blood is the carbonic
acid / hydrogencarbonate ion buffer system.
HCO3-(aq) + H3O+(aq) (1)
H2CO3(aq) + H2O(l)
The diffusion of oxygen into cells and the lungs and carbon dioxide out affects
the pH of the blood, lowering the pH with every inhalation and increasing it with
exhalation, which can cause problems due to fluctuations in pH on a regular
basis. Thus the buffer is needed to regulate these changes to a small degree.
Excess changes are regulated by the buffer system as described above; excess
H3O+ shifts the equilibrium to the left, increasing pH, and excess OH- removes
H3O+ from the system by neutralisation, causing the equilibrium to shift to the
right to decrease pH. Either way, the changes in pH are minimised.
The buffer system also works to control excess addition from external sources
of either acid or base, effectively controlling the pH of the blood to manageable
values.
• See Attachments
• See Attachments
• See Prac Book
• Skill
*
*
5.
reaction between selected
acids and bases
perform a first–hand
investigation to determine
the concentration of a
domestic acidic substance
using computer–based
technologies
describe the differences
between the alkanol and
alkanoic acid functional
groups in carbon
compounds
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
explain the difference in
melting point and boiling
point caused by straight–
chained alkanoic acid and
straight–chained primary
alkanol structures
• See Attachments
• See Prac Book
Alkanol are hydrocarbons in which the functional group is a hydroxy substituent.
For primary alkanols with one hydroxy group, the homologous series is given by
R–OH, where R is CnH2n+1. An example is ethanol.
Alkanoic acids are acidic hydrocarbons where the functional group is a
carboxylic acid substituent. In general, the homologous series is given by R–
COOH, where R is Cn–1H2n–1. An example is ethanoic (acetic) acid (above).
Both molecules are polar and can form hydrogen bonds due to the presence of
oxygen, hydrogen, and hydroxy groups.
• Skill
Alkanol + Alkanoic Acid
Alkyl Alkanoate + Water
The alkyl group comes from the alkanol and the alkanoate comes from the
alkanoic acid. Their naming is based on their originating compounds. For
example
Ethyl Ethanoate + Water
Ethanol + Ethanoic Acid
C2H5OH + CH3COOH
C2H5OCOCH3 + H2O
This is an example of a condensation reaction via the expulsion of a water
molecule from the OH of the acid and H of the alcohol, similar to other processes
of condensation, such as organic polymerisation.
Alkanoic acids have higher melting and boiling points than the alkanol with most
similar molecular mass, which is in turn harder to melt/boil than the
corresponding alkane.
That is:
alkanoic acid > alkanol > alkane
This can be explained in terms of the intermolecular bonding between the
molecules, and the extent to which strong hydrogen bonds occur between
molecules. We consider molecules with as similar molecular masses as
possible to minimise the differentials from dispersion forces, so the changes are
due to the presence of other bonds only.
Therefore, alkanes, which are linked by dispersion forces only, have the lowest
melting/boiling points of the three. Alkanols and alkanoic acids are linked mainly
by hydrogen bonding (for smaller molecules, though larger molecules are still
more strongly linked than alkanes) but their melting and boiling points differ due
to the different numbers of possible hydrogen bonds between molecules.
Alkanoic acids, with their two oxygen atoms, can bond with other to a greater
degree than alkanols with their single oxygen, as shown:
identify esterification as the
reaction between an acid
and an alkanol and
describe, using equations,
examples of esterification
This is represented by the gaseous form of acetic acid, where they form strong
dimers rather than act as discrete molecules.
In addition, alkanoic acids are slightly more polar than alkanols. Esters,
however, do not present sufficient area to hydrogen bond or polar bond
efficiently, and there are no OH groups. It can thus be seen that
alkanoic acid > alkanol > ester > alkane
• See attachments for graphs + tables of MP/BP versus n(C) and molar mass
Esterification is the reaction of an alkanoic acid with an alkanol to produce an
alkyl alkanoate. In the laboratory, it is performed under reflux with a H2SO4
catalyst (covered in detail later). The naming of such compounds is covered
above.
A simple example:
Ethyl Ethanoate + Water
Ethanoic Acid + Ethanol
CH3COOH + C2H5OH
CH3COOC2H5 + H2O
Propan–1–oic Acid + Butanol
C2H5COOH + C4H9OH
Butyl Propanoate + Water
C2H5COOC4H9 + H2O
As can be seen, the hydroxy group leaves the carboxylic acid, combining with
hydrogen of the alkanol’s hydroxy group. Esterification is, in general:
describe the purpose of
using acid in esterification
for catalysis
explain the need for
refluxing during
esterification
+
+ H2O
Where R represents a carbon chain.
Concentrated H2SO4 acts as a catalyst for esterification. It reduces the
activation energy required for the reaction to proceed, providing and alternative
pathway. The catalyst does not participate in the reaction itself. It however
reduces the time for the reaction to reach equilibrium, as the reaction is very
slow.
It also acts as a dehydration agent. When water is removed from the reaction
equilibrium, the equilibrium shifts to the right and increases the yield.
*
*
*
identify data, plan, select
equipment and perform a
first–hand investigation to
prepare an ester using
reflux
outline some examples of
the occurrence, production
and uses of esters
process information from
secondary sources to
identify and describe the
uses of esters as flavours
and perfumes in processed
foods and cosmetics
The reactants and products of the esterification reaction are volatile, and readily
vaporise on heating, causing possible combustion, which is inherently unsafe.
To avoid loss of material from the reaction flask, the mixture is heated using a
reflux apparatus.
A water condenser is mounted above the reaction flask (normally a roundbottomed Pyrex flask) and cold water circulates around the hot, rising vapours.
These condense back to the liquid state and drip back into the reaction flask so
that the reaction can continue without the loss of reactants or products.
Because the system is open to the atmosphere there is no build up of pressure
due to the production of vapours. The organic liquids and vapours are also
flammable, and care must be taken to avoid fires and explosions. The reaction
vessel is normally heated by a hot-water bath (supported on an electric hotplate)
or by using a special electrical heating mantle in which the round-bottom flask is
clamped. Naked flames from a Bunsen burner are avoided to reduce the risk of
fire.
Another feature of the reflux procedure is the use of small boiling chips that are
normally pieces of crushed ceramic. These boiling chips prevent a process
called ‘bumping’ as they provide a large surface area on which vaporisation can
occur without the risks of sudden superheating and the explosive ejection of
vapours.
• See Attachments
• See Prac Book
Esters occur in nature, and can be formed from inorganic acids, such as
phosphate esters. Some of these (ATP) are vitally important in energy transfer
reaction in cells. Fats and oils are triesters of glycerol and long-chain carboxylic
acids, commonly known as triglycerides.
Esters are manufactured as jet engine lubricants, due to their low viscosity and
clean high-temperature operation.
Esters are good solvents and are highly biodegradable compared to oils. Many
natural and synthetic esters are used in wood lacquers, thinners, and a variety
of coatings. Some low molecular weight esters have high volatility and are
useful components of some solvents used in surface finishes. Higher weight
esters are used in latex paints, due to lower volatility. Esters of acetic acid are
used in solvents for personal care and cosmetic products.
Phthalate esters are added to hard plastics to soften them and increase
flexibility. However, animal studies have implicated these esters to reproductive
diseases, including cancer.
Esters of natural fatty acids are used in biofuels, commonly known as biodiesel.
They are biodegradable and are good engine lubricants due to higher viscosity.
Esters are commonly used as flavours, fragrances, and emulsifiers, due to their
characteristic tastes and smells. Methyl salicylate (oil of wintergreen) is used in
lotions to soothe sore muscles. High molecular weight esters are used as
emulsifiers, such as citric acid esters, used to stabilise mayonnaise, margarine
and coffee whitener.
Specific ester uses are:
Flavouring/Perfumes Esters
Ester
Flavour
1-octyl acetate; CH3COOC8H17
Orange
Butyl butanoate; CH3CH2CH2COOC4H9
Pineapple
Ethyl butanoate; CH3CH2CH2COOC2H5
Methyl butanoate; CH3CH2CH2COOCH3
Apple
Ethyl formate; HCOOC2H5
Rum-like taste and odour
Methyl acetate; CH3COOCH3
Mint
Linalyl acetate
Lavender, sage
Other Esters
Ester
1-propyl acetate;
Uses
volatile solvent, wood lacquers, aerosol
CH3COOC3H7
1-butyl acetate;
CH3COOC4H9
1-pentyl acetate (amyl acetate);
CH3COOC5H11
1-butyl propanoate;
CH3CH2COOC4H9
1-pentyl propanoate;
CH3CH2COOC5H11
sprays, cosmetics and personal care
solvent
fragrance solvent, coatings for plastics,
cosmetics and personal care solvent
coatings, cleaning fluids, extraction solvent
for pharmaceuticals, printing, finishing
fabrics
appliance coatings, automotive refinishing,
enamels, lacquers
slow evaporating solvent, appliance
coatings, automotive refinishing, enamels,
lacquers, personal care products