|
 10. The proton theory of acids &
bases - weak or strong acids & bases
Doc
Brown's Chemistry GCSE/IGCSE/O level Science–Chemistry Revision Notes
pH scale of acidity and alkalinity,
acids, bases–alkalis, salts and neutralisation
Index of all my GCSE notes on acids, bases
and salts
All my
GCSE Chemistry Revision
notes
This is a BIG
website, you need to take time to explore it [ SEARCH
BOX]
Use your
mobile phone or ipad etc. in 'landscape' mode
email doc
brown
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. These revision notes on acid-base theory and the strength
of acids and bases should prove useful for the new AQA chemistry, Edexcel
chemistry & OCR chemistry GCSE (9–1, 9-5 & 5-1) science courses.
Note: aq or (aq) means water or
aqueous solution.
Doc Brown's
chemistry revision notes: basic school chemistry science GCSE chemistry, IGCSE chemistry, O level
& ~US grades 8, 9, 10 school science courses for ~14-16 year old science
students for national examinations in chemistry topics including acids
bases alkalis salts preparations reactions
10.
More on Acid–Base Theory and Weak and Strong Acids
-
Some compounds react will water to produce
acidic or alkaline solutions.
-
Water (aq) must be present for a substance to
act as an acid or as a base (at least at gcse level!).
-
A substance dissolving in water and
splitting up into ions undergoes an example of ionisation.
-
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.
-
e.g. hydrochloric acid:
HCl(g) + aq
==> H+(aq) + Cl–(aq)
-
or sulfuric acid:
H2SO4(l)
+ aq ==> 2H+(aq) + SO42–(aq)
-
Note:
+ aq is simply
'shorthand' for adding to water!
-
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.
-
e.g. sodium hydroxide:
NaOH(s) + aq ==> Na+(aq) +
OH–(aq)
-
or calcium hydroxide:
Ca(OH)2(s)
+ aq ==> Ca2+(aq) + 2OH–(aq)
-
Reminder: An alkali is a soluble base.
-
When alkalis and acids
react, the 'general word' and '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)
-
because all
acids form hydrogen ions in water and all alkalis (soluble
bases) form hydroxide ions in water. -
and, in this case,
the remaining ions e.g. sodium Na+(aq) and
chloride Cl–(aq)
become the salt crystals of sodium chloride 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).
-
The hydrogen ion (H+)
is equivalent to a proton - you can use either term.
-
e.g. here the hydroxide ion is the base
and accepts a proton from an acid to form water (as in
neutralisation).
-
H+(aq)
+ OH–(aq) ==> H2O(l)
-
This simple ionic equation
represents what happens when any soluble alkali (metal
hydroxides like NaOH) is mixed with any acid (like HCl etc.).
-
Water is formed, and the ions
left in solution constitute the ions that make up a salt that
can be crystallised out of solution by evaporating the water.
-
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, because non formed!)
-
HCl(g) +
NH3(g) ==> NH4+Cl–(s)
-
You can write an almost identical
equation when aqueous ammonia is mixed with dilute hydrochloric acid.
-
HCl(aq) +
NH3(aq) ==> NH4Cl(aq)
-
You can evaporate and crystallise out
ammonium chloride crystals.
-
or copper(II) oxide (base) +
sulfuric acid (acid) ==> copper(II) sulfate + water
-
CuO(s)
+ H2SO4(aq)
==> CuSO4(aq) + H2O(l)
-
ionically it is: Cu2+O2–(s)
+ 2H+(aq)
==> Cu2+(aq) + H2O(l)
-
The copper oxide is the base because
it accepts hydrogen ions from the acid.
-
The sulfate ion is a
spectator ion.
-
Copper(II) oxide is an insoluble
base, therefore can't be an alkali.
-
You can analyse what happens when a
carbonate dissolves in an acid
-
e.g.
magnesium
carbonate + sulfuric acid ==> magnesium sulfate + water +
carbon dioxide
-
Sulfuric acid donates hydrogen
ions to the base ion - the carbonate ion,
-
so in terms of ions, the reaction
between ANY carbonate and ANY acid can be written as:
-
2H+ +
CO32- ===> H2O +
CO2
-
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
However water is
an amphoteric oxide i.e. it reacts as both a proton acceptor and a
proton donator.
-
Amphoteric means
something that can act either as an acid or as a base.
-
To illustrate water
functioning as both an acid and a base ...
-
e.g. water
acting as a base – proton acceptor with a stronger acid like the
hydrogen chloride gas
-
HCl(g)
+ H2O(l) ==> H3O+(aq)
+ Cl–(aq)
-
more simply: HCl(g)
+ aq) ==> H+(aq)
+ Cl–(aq)
-
This is how
hydrochloric acid is formed which you write simply as HCl(aq).
-
e.g. water
acting as an acid – proton donor with a weak BUT stronger base like
the alkaline gas ammonia
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 at the time there was no pre–existing (and proven) theory of
ion formation.
Some extra important NOTES on the pH Scale
that you need to know
(a) pH is a measure of the hydrogen ion (H+)
concentration
The lower the pH, the higher the
hydrogen ion concentration, the more acidic is the solution.
I know this seems confusing, but that's the way the pH
scale has been defined historically.
(b) Each pH unit change is equivalent to a 10x change in
concentration of the hydrogen ion
For example changing the pH of a
solution from pH 3 to pH 2 makes the solution 10x more acidic.
Changing a solution's pH from 1 to 3 makes it 100x
less acidic (10 x 10), this is similar to comparing solutions of a
strong acid with that of a weak acid - its all about the extent of
ionisation and the resulting concentration of hydrogen ions - read on!
(c) The pH scale and acid
concentration
Its quite easy to relate the pH
of an acid solution from its concentration of hydrogen ions
(without going too much into
the maths!).
The concentration of hydrogen
ions (mol/dm3) = 10-pH (ONLY advanced level
students need to know all this maths)
So, if the pH is 3.0, the
hydrogen ion concentration = 10-3 or 0.001 mol/dm-3
If the pH is 1.0, the
hydrogen ion concentration = 10-1 or 0.1 mol/dm-3
You can then work the argument
the other way: (pH = -log10(H+ concentration))
If the concentration of
hydrochloric acid is 0.5 mol/dm3, strong acid,
so hydrogen ion concentration
is also 0.5 mol/dm3
the pH is -log10(0.5)
= 0.30
Please note that GCSE students
do not have to do these calculations - but perhaps make an estimate?
Advanced A level
chemistry notes - equilibrium section 5
5.1
Lewis and Bronsted–Lowry acid–base theories
5.2
self–ionisation of water and pH scale
5.3
strong acids - examples -
Ph calculations
5.4
weak acids - examples & pH-Ka-pKa calculations
5.5
strong bases -examples -pH calculations
5.6
weak bases - examples & pH-Kb-pKb calculations
WEAK ACIDS and STRONG ACIDS
-
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.
-
Ionisation in this context means on
dissolving a substance in water, it splits into positive and negative ions.
-
Strong means a high degree of
ionisation (~100%) and weak means a low degree of ionisation (often
<<100%).
-
Examples are explained below, but
after avoiding misunderstanding and misusing terms like dilute and
concentrated.
-
BUT FIRST ....
-
Do not confuse
the terms weak
and strong, which is about how far the 'molecules' become ionised in
water with the terms dilute and
concentrated, they mean completely different things!
-
Dilute and
concentrated refer to the concentration of the acid or alkali in
terms of how much of the original material is dissolved
in water as measured by concentration e.g. molarity i.e. a little
or a lot in a given volume of solution, low concentration or high
concentration.
-
The original material might be: HCl,
HNO3, H2SO4, NaOH, Ca(OH)2,
and dilute and concentrated refer to how many particles of the formula
are in a given volume e.g.
-
 
-
The above diagrams illustrate a less
concentrated (more dilute) solution on the left and a more concentrated
solution (less dilute) on the right. The fact that the diagrams show a
mixture of chemicals A and B is irrelevant.
-
It is
completely independent of what concentration of hydrogen ions (in an
acid) or concentration of hydroxide ions (in an alkali) is formed when
the substance is dissolved in water - but that's when the terms 'strong'
and 'weak' are necessary.
-
So, what do we mean by the terms 'weak' and 'strong' when talking about
acids (and soluble bases-alkalis)
-
A strong acid or strong alkali is one that
is that is nearly or completely 100% ionised in water
(not an equilibrium situation)
-
Examples of strong acids
are hydrochloric,
nitric and sulfuric acids.
-
e.g. when dissolving in
water (aq) the maximum (or nearly) hydrogen
ion concentration results in the lowest pH ...
-
hydrochloric acid:
HCl(g) + aq ==> H+(aq) + Cl–(aq)
-
nitric acid:
HNO3(l) + aq ==> H+(aq) +
NO3–(aq)
-
sulfuric acid:
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 e.g. pH 0–1 for a
given concentration.
-
All three of these acids
are ~100% ionised (~100% dissociated into ions) in aqueous solution, which is why they are called
strong acids.
-
For a given concentration, the
lower the pH, the stronger the acid, irrespective of how strong
or weak the acid is.
-
Its all about the extent to
which the original acid molecules are dissociated into ions.
-
Examples of strong alkalis
(soluble strong bases) are sodium hydroxide or potassium hydroxide
etc. (usually Group 1 like NaOH & KOH etc. or Group 2 hydroxides like
calcium hydroxide, Ca(OH)2).
-
e.g. when dissolving in
water (aq) the maximum (or nearly)
hydroxide ion concentration results in the highest pH ...
-
potassium hydroxide:
KOH(s) + aq ==> K+(aq) + OH–(aq)
- calcium hydroxide Ca(OH)2(s)
+ aq ==> Ca2+(aq) + 2OH–(aq)
- The greater the
concentration of hydroxide ions the higher the pH, so strong alkalis
make the most alkaline solutions e.g. pH 13–14.
- Again, all of these alkalis are
~100% ionised (dissociated into free ions) in aqueous solution, that's why they are described as
strong bases (alkalis).
-
For a given concentration, the
higher the pH, the stronger the alkali/base, irrespective of how
strong or weak the soluble base (alkali) is.
-
A
weak acid or alkali is only partially ionised in water
(often just a few %)
-
Ionisation is usually far less than 100%
for strong acids or alkalis, so far less hydrogen ions or hydroxide ions formed.
-
Examples of weak acids are ethanoic
acid (in vinegar), citric acid (in citrus fruits) and carbonic
acids ('soda water', fizzy drinks) and these weak acids ionise via
reversible reactions
and an equilibrium is reached when they dissolve in water.
-
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) ...
-
ethanoic acid: CH3COOH(aq) CH3COO–(aq)
+ H+(aq)
-
undissociated acid ~98% <===> dissociated
acid ions ~2%
-
this gives a pH of
around 3.
-
The ionisation of a weak acid is
a reversible reaction and equilibrium is set up.
-
The equilibrium position is
mainly on the left-hand side, the unionised (undissociated) ethanoic
acid.
-
carbonic acid: H2CO3(aq)HCO3–(aq)
+ H+(aq)
-
undissociated acid <===> dissociated
acid ions
-
This gives water a pH of ~3-5, it would be lower in fizzy drinks that
rainwater!
-
You can use a
reversible sign like
 to
indicate the weak acid equilibrium is much more on the left side
than the right side.
-
carbon dioxide in water ('carbonic
acid') forms a weakly acid solution (H2O + CO2
==> H2CO3).
-
The pH of unpolluted rainwater is
about pH 5.5 due to carbon dioxide dissolved from the atmosphere.
-
The steady rise in atmospheric CO2
level, is causing concern to environmentalists because along with
climate change, the oceans are becoming more acidic and this does
affect marine ecosystems.
-
Many organic acids are weak
acids e.g. citric acid from fruit.
-
An example of a weak
soluble base
(weak alkali) is ammonia
solution, only about 2% changes to the ionic forms on the right of the
equation as written below. Sodium carbonate is also a soluble weak
base in aqueous solution. Again these weak soluble bases (alkalis)
ionise via reversible reactions and an equilibrium
is reached when they dissolve in water.
-
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 is still over 7, typically pH 9 to pH 11) ...
-
NH3(aq) + H2O(l)NH4+(aq)
+ OH–(aq)
-
Only about 2% of the
ammonia ionises to form ammonium ions and hydroxide ions giving a pH of ~10.
-
The ionisation of a weak
base/alkali is a reversible reaction and equilibrium is set up.
-
The equilibrium position is
mainly on the left-hand side, the unionised ammonia.
-
sodium carbonate: CO32–
+ H2O(l)
HCO3–(aq) + OH–(aq)
-
both of which, when
dissolved in water, produce hydroxide ions (OH–) giving an alkaline solution, despite the fact that
OH doesn't appear in their
formulae!
-
Take care in combing the two sets of terms e.g. you
can have:
-
a dilute solution of a weak acid e.g. 0.1
mol/dm3 solution of ethanoic acid (CH3COOH(aq)CH3COO–(aq)
+ H+(aq))
-
a concentrated solution of a weak acid
e.g. a 10 mol/dm3 solution of ethanoic acid
-
a dilute solution of a strong acid e.g. a
0.1 mol/dm3 solution of hydrochloric acid (HCl(g) + aq
==> H+(aq) + Cl–(aq))
-
a concentrated solution of a strong acid
e.g. a 10 mol/dm3 solution of hydrochloric acid
-
How can you distinguish a weak
acid from a strong acid?
-
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.
-
You could ...
-
(i) Compare the pH of
solutions of equal concentration (equal molarity) and measure the pH
with an accurately calibrated pH meter.
-
(ii) You could get a rough estimate
from universal indicator solution or paper, but that's not very
accurate.
-
The differences between strong
and weak acids shows up in other sorts of experiments e.g.
-
The
pH of solutions of equal
concentration e.g. of molarity 1.0 mol/dm3
-
The
rate of reaction with metals.
(1 molar means 1.0 mol/dm3, sometimes written as 1M for
shorthand))
-
If you put magnesium ribbon into 1 molar
solutions of hydrochloric acid (strong, high % ionisation so high H+(aq)
concentration) and 1 molar solution of ethanoic acid (weak, low percentage ionization so much lower H+(aq)
concentration), you can
see the difference in the fast and slow 'fizzing' rates!
-
You can repeat the experiment
using calcium carbonate (limestone granules) instead of magnesium ribbon.
-
You can do simple
rate of reaction experiments comparing has
fast the gas is evolved from the reaction mixture.
-
-
-
The above graph shows the sort
of results you might expect by adding the same masses of magnesium ribbon or
calcium carbonate granules to the same volume of ethanoic acid, CH3COOH,
or hydrochloric acid, HCl, of equal concentration e.g. both acids
with a concentration of 1.0 mol/dm3.
-
The principal observation is
that the rate of reaction with the strong hydrochloric acid is much
greater than the rate of reaction of the weak ethanoic acid. You can
tell this from the gradient of the rate of reaction, particularly at the
start of the reaction.
-
The acid equilibrium in the case
of hydrochloric acid, is 100% to the right in forming hydrogen ions.
-
In the case of ethanoic acid,
the equilibrium is ~98% to the left-hand side, so far fewer hydrogen ions
are produced to react with the magnesium or limestone.
-
These percentages mean that the
hydrogen ion concentration in hydrochloric acid is about 50 times that in
ethanoic acid (100% to 2%) AND it is the hydrogen ion that actually reacts
directly with the metal or the carbonate.
-
So the rate observations can be
interpreted as a function of the hydrogen ion concentration.
-
The greater the hydrogen ion
concentration the faster the reaction because there is an increased
probability of a hydrogen ion hitting the surface of the magnesium ribbon or
a limestone granule i.e. a greater collision frequency, greater chance of a
fruitful collision of the right kinetic energy to change reactants into
products.
-
Therefore the collision
frequency of the active ingredient (hydrogen ion) is much greater in the
strong acid than the weak acid.
-
On this simple numerical basis,
you might expect the hydrochloric acid to react 50 times faster than the
ethanoic acid (100% to 2% ionisation), applying the simple concentration
rule for reaction speed.
-
What you also find, in both
cases, the same final volume of gas is the same in both cases, but the
strong acid completes the reaction in a much shorter time.
-
The little arrows show when the
reaction ceased in each case, that is when the graph of gas volume first
becomes horizontal.
-
So, you might ask the question -
if ethanoic acid only ionises about 2%, how can it still produce the same
volume of hydrogen or carbon dioxide as hydrochloric acid by the end?
-
Well the answer is 'quite
simple', the ionisation of the weak acid continues as hydrogen ions are used
up.
-
As the hydrogen ions are used
up, more ethanoic acid ionises to replace those used up in the reaction to
try and maintain the position of the equilibrium.
-
This is a general rule about
reversible reactions and a chemical equilibrium. The system will always try
to maintain a balance in the ratio of reactants to products (reactants
products), if you remove a component from the equilibrium, the system tries
to replace it.
-
This contrasts with the strong
acid solution where all possible hydrogen ions exist right from the start of
the reaction.
-
The equations for these
reactions producing the gases hydrogen or carbon dioxide are as follows ...
-
(fast) magnesium + hydrochloric
acid ==> magnesium chloride + hydrogen
-
(slow) magnesium + ethanoic acid
==> magnesium ethanoate
-
(fast) calcium carbonate +
hydrochloric acid ==> calcium chloride + water + carbon dioxide
-
(slow) calcium carbonate +
ethanoic acid ==> calcium ethanoate + water + carbon dioxide
-
-
-
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 of hydrochloric
acid and ethanoic acid, you get a greater
rate 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 length time at
the same voltage (potential difference, p.d.) before
measuring the gas volumes of hydrogen formed. (Electrolysis
methods).
-
In the solution undergoing
electrolysis, the current is carried by the charged particles e.g. hydrogen
ions, hydroxide ions, chloride ions, ethanoate ions etc., therefore the
greater the ionisation of the acid the greater the concentration of ions.
-
The greater concentration of
ions in the strong acid solution reduces the electrical resistance and more
current flows via the greater number of ions present to carry it, hence more
hydrogen ions are reduced at the cathode to form hydrogen.
-
They will of course, both
produce hydrogen in electrolysis cells, because all acids produce some
hydrogen ions in water, but the rate of electrolysis is highly dependant on
the hydrogen ion concentration, which in turn is dependant on the strength
of the acid using solutions of the same concentration (as measured by
molarity).
-
Remember that its
the hydrogen ion, the H+ ion, is the active chemical species in acid
solutions NOT a 'HCl' or a 'H2SO4' or a 'CH3COOH'
molecule.
pH titration curves
for weak acids and weak soluble bases (alkalis) - relating to
feasibility of a quantitative titration
In
section 7. pH titration curves were introduced by looking
what happens when you add a strong acid to a strong alkali.
For titrations involving weak acids
or weak soluble bases (weak alkalis) things are a bit more complicated.
You need to add sections together
from the graph lines to describe what happens for a particular
combination AND order of addition.
Graph A adding an alkali to
an acid - pH changes
Lines (1) + (4) is adding a
strong alkali to a strong acid - this titration gives a distinct end
point in a titration
e.g. adding sodium hydroxide
solution to hydrochloric acid
Lines (1) + (3) is adding a weak
alkali to a strong acid - this titration gives a poorer end point in
a titration
e.g. adding ammonia solution
to hydrochloric acid
Lines (2) + (4) is adding a
strong alkali to a weak acid - this titration gives a poorer end
point in a titration
e.g. adding sodium hydroxide
solution to ethanoic acid solution
Lines (2) + (3) is adding a weak
alkali to a weak base - this titration gives a very poor end point
in a titration - not practical
e.g. adding aqueous ammonia
to ethanoic acid solution
Graph B adding an acid to an
alkali - pH changes
Lines (1) + (4) is adding a
strong acid to a strong alkali - this titration gives a distinct end
point in a titration
e.g. adding hydrochloric acid
to sodium hydroxide solution
Lines (1) + (3) is adding a weak
acid to a strong alkali - this titration gives a poorer end point in
a titration
e.g. adding ethanoic acid
solution to sodium hydroxide solution
Lines (2) + (4) is adding a
strong acid to a weak alkali - this titration gives a poorer end
point in a titration
e.g. adding hydrochloric acid
to aqueous ammonia solution
Lines (2) + (3) is adding a weak
acid to a weak alkali - this titration gives a very poor end point
in a titration - not practical
e.g. adding ethanoic acid
solution to ammonia solution.
A summary of some points about
pH, acids and alkalis
-
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 a 'rough' guide to the
'strength' of an acid or alkali:
-
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 (hydrogen ion 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
GCSE/IGCSE Acids & Alkalis revision notes sub–index:
Index of all pH, Acids, Alkalis, Salts Notes 1.
Examples of everyday acids, alkalis, salts, pH of
solution, hazard warning signs : 2.
pH scale, indicators, ionic theory of acids–alkali neutralisation : 4.
Reactions of acids with
metals/oxides/hydroxides/carbonates, neutralisation reactions : 5.
Reactions of bases–alkalis
like ammonia & sodium hydroxide : 6. Four methods
of making salts : 7. Changes in pH in a
neutralisation, choice and use of indicators : 8. Important formulae
of compounds, salt solubility and water of crystallisation :
10.
More on Acid–Base Theory and Weak and Strong Acids
See also
Advanced Level Chemistry Students Acid–Base Revision
Notes – use index
-
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
[WEBSITE SEARCH
BOX]
|
|
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
 |
| Website content © Dr
Phil Brown 2000+. All copyrights reserved on revision notes, images,
quizzes, worksheets etc. Copying of website material is NOT
permitted. Exam revision summaries & references to science course specifications
are unofficial. |
|
sign), so the longer half–arrow to
the left tells you that most water remains as water molecules!]
Acids, Bases, pH page
section 2.
