H4
Acids and Bases in Everyday Life
This chapter investigates the every day uses of acids and bases, the environmental
effects and safety aspects of using acids, as well as using pH to describe the acidity of
substances. Many acid-base reactions are examples of equilibrium reactions and the
chapter looks at the application of Le Chatelier’s principle to these reactions.
Acids
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Bases
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Indicators
Substance that produce hydrogen ions (hydronium ions
H3O+) in solution (Arrhenius definition)
Acids have a sour taste
Acids sting or burn the skin
Acids conduct electricity
Turn blue litmus red
Substance that contain the oxide ion or hydroxide or
which may produce the hydroxide ion in solution
(Arrhenius definition)
Alkalis are soluble bases
Alkalis have a soapy feel
Alkalis have a bitter taste
Alkalis are good conductors of electricity in solution
Alkalis turn red litmus blue
Substances that change colour in acidic or alkaline solutions.
Common indicators include
 Methyl orange
 Bromothymol blue
 Litmus
 Phenolphthalein
Indicators may be used to
 Test the acidity or alkalinity of soils
 Test the acidity of home swimming pools
 Monitor wastes from photographic processes.
Neutralisation
The reaction of an acid with a base
acid + base
salt + water
Salt
An ionic compound that forms when an acid reacts with a
base.
Equations for acidbase reactions
Neutral species equation
NaOH(aq) + HCl(aq)
NaCl(aq)
Complete ionic equation
Na+ + OH- + H+ + Cl-
Na+ + Cl- + H2O(aq)
Net ionic equation
H+ + OH-
+ H2O(aq)
H2O(aq)
Spectator ions
Ions that are not involved in acid-base reactions. The sodium
ion and the chloride ion in the reaction of sodium hydroxide
with hydrochloric acid are examples of spectator ions.
Common acids
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Hydrochloric acid
Sulfuric acid
Sulphurous acid
Nitric acid
Nitrous acid
Carbonic acid
Phosphoric acid
Acetic (ethanoic) acid
Hydrocyanic acid
Hydrogen sulfide
Acidic oxides
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React with water to form an acid
React with bases to form salts
Oxides of non-metals
Covalent
Basic oxides
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React with acids to form salts
Do not react with alkali solutions
Metal oxides
Ionic
Amphoteric oxides
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React with acids to form salts and also react with alkalis.
Oxides of zinc, aluminium, lead and tin
Ionic
Neutral oxides
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React with neither acids nor alkalis.
Carbon monoxide, nitrous oxide and nitric oxide.
Covalent
(HCl)
(H2SO4)
(H2SO3)
(HNO3)
(HNO2)
(H2CO3)
(H3PO4)
(CH3COOH)
(HCN)
(H2S)
Le Chatelier’s
Principle
If a system at equilibrium is disturbed, then the system adjusts
itself so as to minimise the disturbance.
A system is disturbed if
 The concentration (or pressure) of one or more of the
species involved is changed.
 The total pressure acting upon a reaction that involves
gases is changed.
 The temperature is altered.
The carbon-dioxide
water equilibrium
The equilibrium can be described by the equation
CO2(g) + H2O(l)
H2CO3(aq)
The forward reaction is favoured by
 Increases amounts of carbon dioxide gas
 Increased pressure
 Decreased temperature (the forward reaction is
exothermic)
Acidic oxides in the
atmosphere
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Acid rain
Rain that has a higher hydrogen ion concentration than
normal, higher than about 10-5 mol L-1.
Carbon dioxide (360ppm)
Sulfur dioxide (0.001 ppm)
Nitrogen dioxide (0.001 ppm)
The effects of acid rain include
 Increasing acidity of lakes
 Damage to pine forests
 Erosion of marble and limestone of building surfaces and
decorations
Avogadro’s
hypothesis
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When measured at the same temperature and pressure,
equal volumes of gases contain the same number of
molecules.
Equal numbers of molecules of different gases occupy the
same volume, at the same temperature and pressure.
Avogadro’s number
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The number of particles in one mole.
6 x 1023 particles
Molar volume of
gases
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22.4 litres at STP (0oC and 101.3kPa)
24.46 litres at SLC (25oC and 101.3kPa)
Synthetic acids
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Sulfuric acid
Nitric acid
Phosphoric acid
pH
Defined as the negative logarithm (to the base 10) of the
hydrogen ion concentration.
pH = -log10[H3O+]
[H3O+] = 10-pH
Self ionisation of
water
H2O
H3O+
+ OH-
[H3O+][ OH-]
= 10-14
+ H2O
Kw
K
w
=
is the ionic product of water
[H3O+][ OH-]
pH of alkaline
solutions
pH of solutions
pH
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= 14.00
= 10-14
+ log[OH-]
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A neutral solution is defined as a solution in which the
concentration of hydrogen ions is equal to the
concentration of hydroxide ions.
A neutral solution has pH = 7.00
An acidic solution has pH < 7.00
A basic solution has pH > 7.00
Measuring pH
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Use a variety of indicators
Using a mixture of indicators such as universal indicator
Indicator papers
pH meter
Strong acid
All the acid present in the solution has ionised to form
hydrogen ions, there are no neutral acid molecules present.
Weak acid
An acid in which only some of the acid molecules present in
the solution have ionised to form hydrogen ions.
Degree of ionisation
The fraction of the molecules that have ionised.
Concentrated acid
Solution in which the total concentration of solute species is
high (above 5 mol L-1)
Dilute solution
Solution in which the total concentration of solute species is
low (less than 2 mol L-1)