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Brown's Chemistry GCSE/IGCSE Science-Chemistry Revision Notes
pH scale of acidity and alkalinity,
acids, bases-alkalis, salts and neutralisation
Part
10.
More on acid-base theory and weak &
strong acids
Part 10 introduces students to the more advanced
theory of acids and bases. Acids are defined as proton donors and bases are
defined as proton acceptors. The terms 'weak' and 'strong' are explained when
referring to e.g. weak acids, strong acids, weak bases or strong bases. The
theories of Arrhenius and Bronsted-Lowry are described and examples given and
fully explained.
GCSE/IGCSE Sub-index:
Index of all pH, Acids, Alkalis, Salts Notes 1.
Examples of acid-alkali chemistry : 2.
pH scale, indicators, ionic theory of acids-alkali neutralisation
: 3. pH examples of
acid, neutral or alkaline
solutions : 4. Acid reactions with
metals/oxides/hydroxides/carbonates and neutralisation reactions : 5.
Reactions of bases-alkalis
like sodium hydroxide : 6. Four methods
of making salts : 7. Changes in pH in a
neutralisation : 8. Important formulae, salt
solubility and water of crystallisation : 9. Further examples of word/symbol equations
for salt preparations :
10.
More on Acid-Base Theory and Weak and Strong Acids
See also
Advanced Level Chemistry Students Acid-Base Revision
Notes - use index
10.
More on Acid-Base Theory and Weak and Strong Acids
-
Some compounds react will water to produce
acidic or alkaline solutions.
-
Water must be present for a substance to
act as an acid or as a base (usually at gcse level!).
-
Acids in aqueous solution produce
hydrogen H+
ions. The H+ ion is a proton. In water this proton is hydrated
(associated with water and more correctly expressed as H3O+(aq))
but H+(aq) is adequate here. The greater the
concentration of hydrogen ions the more acid the solution and the lower the pH.
-
Alkalis in aqueous solution produce OH-(aq)
hydroxide ions. The greater the concentration of
hydroxide ions the more alkaline the solution and the higher the pH.
-
When alkalis and acids
react,
the 'general word' and e.g. 'molecular formula' neutralisation equation might be ...
-
ACID
+ ALKALI ==> SALT
+ WATER
-
e.g.
-
 hydrochloric
acid + sodium hydroxide ==> sodium
chloride + water
-
HCl(aq)
+ NaOH(aq)
==> NaCl(aq) + H2O(l)
-
BUT
the ionic equation for ANY neutralisation is
-
H+(aq)
+ OH-(aq) ==> H2O(l)
-
and the remaining ions
e.g. Na+(aq) and Cl-(aq) become
the salt crystals NaCl(s) on evaporating the water.
-
Acids can be defined as proton donors.
A base can be defined as a proton acceptor (Bronsted-Lowry theory).
-
e.g. here the hydroxide ion is the base
and accepts a proton from an acid.
-
or here the hydrogen chloride is the
acid and the ammonia is the base when ammonium chloride is formed when
the two gases are mixed. The acid hydrogen chloride donates a proton
to the base ammonia. (note: no water present!)
-
or copper(II) oxide (base) +
sulphuric acid (acid) ==> copper(II) sulphate + water
-
CuO(s)
+ H2SO4(aq)
==> CuSO4(aq) + H2O(l)
-
ionically it is: Cu2+O2-(s)
+ 2H+(aq)
==> Cu2+(aq) + H2O(l)
-
Acids are characterised by having at
least one replaceable hydrogen atom in forming a salt, the H is
replaced by a metal ion (Na+, Mg2+ etc.) or the
ammonium ion (NH4+):
-
Incidentally water is a
neutral oxide because its pH is 7
-
Several scientists have made contributions
to ionic and acid-base theory e.g.
-
Arrhenius (1887), was one of the first
scientists to suggest that substances could split into free positive
and negative ions when
dissolved in water, the so called 'electrolytic dissociation'
giving rise to electrically conducting solutions. His theory was
considered a bit revolutionary, and he was given a low rating for his PhD
at Paris at first! - however the 'professors' recanted when
other scientists decided it was a good idea and in 1903 he was awarded
the Nobel Prize for his ionic theory work!
-
Lowry and Bronsted (1923) took
further the work of Arrhenius and applied ionic theory to the concept
of acids and bases - that is, that acids are proton donors
and bases are acceptors.
-
e.g. the reaction of ammonia
with acids ...
-
in the reaction ...
-
HCl(g) +
NH3(g)
==> NH4+Cl-(s)
-
hydrogen chloride is the acid
- a proton donor leaving the chloride ion Cl-
-
AND
-
ammonia is the base - a
proton acceptor to form the ammonium ion NH4+
-
It should be noted that the work of Arrhenius took much longer to be accepted than the work of
Lowry and Bronsted because there was no pre-existing (and proven) theory of
ion formation.
-
Acids and alkalis are
further classified by the
extent of their ionisation in water.
-
They are described as strong or weak
depending on their degree of ionisation in water.
-
Do not confuse
the terms weak
and strong about how far the 'molecules' become ionised in
water with the terms dilute and
concentrated, they mean different things!
-
Dilute and
concentrated refer to the concentration of the acid or
alkali in terms of how much (i.e. a little or a lot) of the original material is dissolved
in water as measured by concentration e.g. molarity.
-
A strong acid or alkali is one that
is that is nearly or completely 100% ionised in water
(not an equilibrium situation)
-
examples of strong acids
are hydrochloric,
sulphuric and nitric acids.
-
e.g. the maximum (or nearly) hydrogen
ion concentration results in the lowest pH ...
-
nitric acid is:
HNO3(l) + aq ==> H+(aq) +
NO3-(aq)
-
and sulphuric acid is:
H2SO4(l) + aq ==> 2H+(aq) + SO42-(aq)
-
The greater
the concentration of hydrogen ions the lower the pH, so strong
acids make the most acidic solutions.
-
examples of strong alkalis
(soluble strong bases) are sodium hydroxide or potassium hydroxide
etc. (usually Group 1 or 2 hydroxides).
-
A weak acid or alkali is only partially
ionised in water.
-
examples of weak acids are ethanoic, citric and carbonic
acids.
-
e.g. for ethanoic about 2% ionises
(forward reaction to the right), the
equilibrium lies mainly to the un-ionised form on the left and for the
weaker carbonic acid even less is ionised. So only a relatively low
concentration of free hydrogen ions form giving a less acidic higher pH
solution than
strong acids (but pH still less than 7) ...
-
An example of a weak alkali/base
(weak soluble base) is ammonia
solution, about 2% changes to the ionic forms on the right. So only a
relatively low concentration of free hydroxide ions form giving a less
alkaline solution, so the pH is less than a strong base/alkali (but
pH still over 7) ...
-
NH3(aq) + H2O(l) NH4+(aq)
+ OH-(aq)
-
or sodium carbonate: CO32-
+ H2O(l)
HCO3-(aq) + OH-(aq)
-
both of which, when
dissolved in water, produce hydroxide ions giving an alkaline solution, despite the fact that OH doesn't appear in their
formulae!
-
You can distinguish between strong and weak acids of the same concentration by
using the pH scale and observations from a variety of experiments
support the low or high of ionisation theory.
-
e.g. by the rate of reaction with metals.
-
If you put magnesium ribbon into 1 molar
solutions of hydrochloric acid (strong, high % ionisation so high H+(aq)
concentration) and ethanoic acid (weak, low percentage ionization so much lower H+(aq)
concentration), you can
see the difference in the fast and slow 'fizzing' rates!
-
Since stronger/weak
acid solutions (or alkalis) contain more/less hydrogen ions, they are
better/poorer conductors of electricity.
-
e.g. If you carry out
electrolysis experiments with the same two solutions, you get a much greater
volume of hydrogen collected at the cathode from the hydrochloric acid
compared to the ethanoic acid.
-
You must use solutions
of the same concentration and electrolysed them for the same time before
measuring the gas volumes (Electrolysis
methods 1a and 1b).
-
Remember that its
the H+ ion that is the active chemical species in acid
solutions NOT a 'HCl' or a 'H2SO4' or a 'CH3COOH'
molecule.
-
More on pH scale and
indicators in section 2.
-
The pH is dependent on the relative concentrations of the H+(aq)
and the OH-(aq) concentrations.
-
a high H+(aq)
concentration means a low pH
-
lower H+(aq)
concentration means higher pH and higher OH-(aq)
concentration, less acid
-
a high OH-(aq)
concentration means a high pH
-
lower OH-(aq)
concentration means lower pH and higher H+(aq)
concentration, less alkaline
-
In general:
-
Neutralisation ionically is: H+(aq)
+ OH-(aq)
==> H2O(l)
(exothermic)
-
The pH of a solution, or determining
the neutralisation point, can be measured with
-
When mixing an acid and alkali the
neutralisation end-point can also be determined by
-
SUMMARY of the advanced BRONSTED
LOWRY THEORY
-
a Bronsted-Lowry
acid is defined as a proton donor (H+),
-
and a Bronsted-Lowry
base is defined as a proton acceptor e.g. two examples
-
(i) hydrogen chloride gas +
ammonia gas ==> ammonium chloride solid
-
HCl(g) + NH3(g)
==> NH4Cl(s)
-
acidic hydrogen chloride
gives a proton to the ammonia molecule base to give the ammonium ion (NH4+).
-
(ii) copper oxide dissolves
in acid solutions
-
copper(II) oxide + sulfuric
acid ==> copper(II) sulfate + water.
-
CuO(s) + H2SO4(aq)
==> CuSO4(aq) + H2O(l)
-
copper oxide is the base
because it reacts with protons from the acid to form water.
-
Incidentally water
is a neutral oxide because its pH is 7.
-
However water is
an amphoteric oxide i.e. it reacts as both a proton acceptor and a
proton donator.
-
e.g. water
acting as a base - proton acceptor with a stronger acid like the
hydrogen chloride gas
-
e.g. water
acting as an acid - proton donor with a weak BUT stronger base like
the alkaline gas ammonia
-
Further work-study
links:
-
Multiple choice revision quizzes and other worksheets
-
GCSE/IGCSE foundation-easier multiple choice quiz on pH, Indicators, Acids,
Bases, Neutralisation and Salts
-
GCSE/IGCSE higher-harder multiple choice quiz on pH, Indicators, Acids,
Bases, Neutralisation and Salts
-
GCSE/IGCSE Structured question worksheet on Acid
Reaction word equations and
symbol
equation questions
-
Word
equation answers and
symbol
equation answers)
-
GCSE/IGCSE word-fill worksheet on Acids,
Bases, Neutralisation and Salts
-
GCSE/IGCSE
matching pair quiz on Acids, Bases, Salts and pH
-
See also
Advanced Level Chemistry Students Acid-Base Revision
Notes - use index
Revision KS4 Science GCSE/IGCSE/O level
Chemistry Information Study Notes for revising for AQA GCSE Science, Edexcel
360Science/IGCSE Chemistry & OCR 21stC Science, OCR Gateway Science
WJEC gcse science chemistry CCEA/CEA gcse science chemistry O Level
Chemistry (revise courses equal to US grade 8, grade 9
grade 10) tuition help science chemistry courses revision guides
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NH4+(aq)
+ OH-(aq)