Chapter 29 – Acids, Bases and pH
Section A – introduction to acids and bases
Properties of ACIDS
Properties of BASES
Have a pH of greater than 7
Have a pH of less than 7
Can have a sour taste
Can have a soapy feel
Turn blue litmus red
Turn red litmus blue
Ex of household acids: orange juice, vinegar etc
Ex of household bases: toothpaste, oven cleaner
etceyc
Neutralisation reactions
Acid + Base
water + salt
Examples from everyday life:


A basic wasp sting is neutralised by vinegar
Acidic soil is neutralised by lime
Section B – The Arrhenius theory of acids and bases
Arrhenius acid is a substance that dissociates in water, forming hydrogen ions
HA
H+ + A-
A strong acid fully
dissociates in water to
form hydroxide ions
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A weak acid only partially
dissociates in water to
form hydroxide ions
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Chapter 29 – Acids, Bases and pH
Arrhenius base is a substance that dissociates in water, forming hydroxide ions
A weak base only partially
dissociates in water to
form hydroxide ions
A strong base fully
dissociates in water to
form hydroxide ions
Section C – Bronsted-Lowry theory of acids and bases
An acid is a proton ( H+) donor, and a base is a proton ( H+) acceptor.
Acid
Base
Section D – Reversible reactions and Conjugate base pairs
A reversible reaction is one which happens in ______________ directions, this is indicated by the
arrow shown below:
NH4+ + H20
NH3 +H30+
In the forward direction NH4+ is the acid and it turns into NH3 after it donates the proton. The base is
H20 and it turns into H30+ after it accepts the proton.
In the reverse reaction H30+ is the acid and it turns into H20 after it donates the proton. The base is
NH3 and it turns into NH4+ after it accepts the proton.
*This means the Bronsted acid turns into the conjugate base, and the Bronsted base turns into the
conjugate base.
Conjugate acid- base pair an acid and a base that differ in structure by a proton
This example contains two conjugate base pairs:
NH4+ and NH3
2. H20 and H30+
1.
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Chapter 29 – Acids, Bases and pH
Section E – Self ionisation of water
Water can dissociate itself (break up ) in the following way:
The ionic product of water :
H20
H+ + OH-
Kw = [OH-][H+]
(Square Brackets mean concentration in moles per litre)
*A note about Kw and temperature - At room temperature 25OC Kw = 1 x 10-14 but this will increase
as temperature increases because as temperature increases more collisions happen between water
molecules causing more ions to be generated.
in water:
Kw = [OH-][H+]
= [OH-] [H+] at 25OC
[H+] = [OH-], because for
every one water
molecule that dissociates
one H+ ion is made and
one OH- ion is made
every one water
molecule that dissociates
_______H+ ion is made
and ______ OH- ion is
made
Kw = [H+][H+]
Kw = [H+]2
1 x 10-14 =[H+]2
√1 𝑥 10-14 = [H+]
1 x 10-7 = [H+]
at 25OC
******Also
1 x 10-7 = [OH-] at 25OC
 Any neutral solution – like pure water will have [H+] = [OH-],
 If extra H+ ions are added into a neutral solution then [H+] > [OH-], and the solution will be acidic
 If extra OH- ions are added or generated in a neutral solution then [H+] < [OH-] and the solution
will be basic
Section F – The pH scale
pH=
pH= - log10 [H+],
the square brackets mean concentration in moles per litre
**Using this formula provides a way to discuss the acidity of a solution without having to express the
concentration
in moles per litre
concentration of the
[H+] square
in mol/L brackets
which can mean
often involve
very small numbers
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Chapter 29 – Acids, Bases and pH
Example 1 - A solution has a concentration of 7x10-7 moles of H+ per litre. What is the pH?
pH= - log10 [H+],
Answer:
pH= - log10 [7x10-7]
pH = 6.15
The pH scale
This is a scale from 0 – 14 which tells you how acidic or basic a solution is.
The pH of a solution can be measured by using:
1)universal indicator paper and a chart 2) pH meter
The pH scale can be very useful but also has its limitations: it can only be used at a temperature of
25OC and the substance you are measuring must be in a dilute aqueous solution.
*An interesting example is the case of water at temperature over 25OC – we already mentioned that
the Kw (ionic product) of water increases with increasing temperatures.
At
100 OC is Kw = 52.3 x 10-14
[H+] = √52.3 𝑥 10−14
[H+] = .000000723
pH = 6.1407
So you might assume that the water is now acidic , but it is not because even though the
concentration of [H+] is higher than normal we still know that in pure water [H+] = [OH-] and this
means the solution is still neutral!
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Chapter 29 – Acids, Bases and pH
Section G – Calculating pH values
Strong acids
Strong acids fully dissociate in water to produce H+ ions.
Example 1
Calculate the pH of a 0.5 M HF solution
Answer
H+ + F-
HF
0.5 M
0.5 M
pH= - log10 [H+]
pH= - log10 [0.5]
pH= 0.3010
Example 2
Find the pH of a 0.4M solution of sulphuric acid
Answer
H2SO4
2H+ + SO4-2
0.4M
0.8M
pH= - log10 [H+]
pH= - log10 [0.8]
pH= 0.0969
Strong bases
Strong bases fully dissociate in water to produce hydroxide (OH- ) ions.
Calculating the pH of a strong base:
pOH= - log10 [OH-]
pH = 14 - pOH
Example 1 - Calculate the pH of 0.2 M KOH solution
Answer :
For a strong base like KOH
KOH
0.2 M


K + OH0.2 M
To find pOH Use the formula pOH= - log10 [OH- ]
pOH= - log10 [0.2]
pOH= 0.6990
pH = 14 – pOH
pH = 14 – 0.6990 = 13.301
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Chapter 29 – Acids, Bases and pH
Weak acids
A weak acid only partially dissociates in water to produce H+ ions.
For this weak acid which we will call HA:
H+ + A-
HA
Ka is the acid dissociation constant- It will tell you to what extent a weak acid will actually dissociate
in water.
For a weak acid : [H+] = √𝐾𝑎 𝑥 [𝐴𝑐𝑖𝑑]
pH = - log10 [H+]
****Square brackets mean concentration in moles per litre!
Example 1 Calculate the pH of a 0.1 solution of methanoic acid given that the value of the acid
dissociation constant is 1.8 X 10-4
Answer
[H+] = √𝐾𝑎 𝑥 [𝐴𝑐𝑖𝑑]
[H+] = √1.8 X 10-4 𝑥 [0.1]
[H+] = 0.0042
pH= 2.3724
Weak bases
A weak acid only partially dissociate in water to produce OH- ions.
For this weak base like NH3
NH3+ H2O
NH4+ OH-
Kb is the base dissociation constant: It will tell you to what extent a weak acid will actually dissociate
in water.
pH = 14 - pOH
pOH= - log10 [OH-]
[OH-] = √𝐾𝑏 𝑥 [𝐵𝑎𝑠𝑒]
****Square brackets mean concentration in moles per litre!
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Chapter 29 – Acids, Bases and pH
Example 1
Calculate the pH of a 0.01 solution of ammonia NH3 given that the value of the base dissociation
constant is 1.8 X 10-5
Answer
[OH-] = √𝐾𝑏 𝑥 [𝐵𝑎𝑠𝑒]
[OH-] = √1.8 𝑥10-5 𝑥 [0.01]
[OH- ] = .0004
pOH= - log10 [OH-]
pOH= 3.3724
pH = 14 – pOH
pH = 10.6275
Section H – Acid Base Indicators
Since acid/ base indicators are themselves either acids or bases you should not add too much of them
to an acid/base titration as they will themselves react in the reaction and cause an inaccurate titre
result!
Acid/ Base Indicators are substances that change colour depending
on the pH of the solution in which they are placed
Litmus as an example of an indicator
In its undissociated form (HIn)
Litmus is red in colour.
HIn
H+ + In-
In its dissociated form (In-) Litmus
is blue in colour.
If blue Litmus is mixed with an acid then then the acid will donate H+ ions to Litmus IN – and will turn
it into HIn which is red in colour
If red Litmus is mixed with a base then the Indicator HIn will donate H+ ions to the base and it will
change to In- which is blue in colour.
*We will look at the dynamics of these reactions in more detail in ch. Chemical Equilibrium
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Chapter 29 – Acids, Bases and pH
Section G - Titration curves and acid- base indicators
During an acid- base titration the pH will change gradually at first but will make a sudden change
close to the end point of the reaction. Using an indicator can tell us when the end point has been
reached by changing colour at this exact moment.
Strong acid strong base titration
Example: Reaction of hydrochloric acid and sodium hydroxide
At the end point the pH jumps suddenly from 3 to 10 .
A suitable indicator must have one colour at pH 3 and a different
colour at pH10
Suitable indicators: methyl orange, phenolphthalein or Litmus
Strong acid weak base titrations
Example: hydrochloric acid and ammonia
At the end point the pH jumps suddenly from 3 to 7 so the
indicator must have a distinct colour change in this range
Suitable indicator: Methyl orange
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Chapter 29 – Acids, Bases and pH
Weak acid strong base titration
Example ethanoic acid and sodium hydroxide reaction
At the end point the pH jumps suddenly from 7 to 10 so the
indicator must have a distinct colour change in this range
Suitable indicator: phenolphthalein
Weak acid weak base titration
Example ethanoic acid and ammonia reaction
There is NO sharp or sudden pH change during this reaction so there
is NOT a suitable indicator available to detect the end point by means
of a colour change.
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Chapter 29 – Acids, Bases and pH
Self-assessment of
Green = I already know this
Orange = I am not sure – have to check it
Red = I don’t know this yet – have to start learning
Green
Orange
Red
Household acids and bases (two examples of each).
Everyday examples of neutralisation, e.g. use of lime in agriculture,
use of stomach powders for acid indigestion.
Neutralisation – formation of a salt from an acid and a base.
Arrhenius theories of acids and bases
Brønsted-Lowry theories of acids and bases
Conjugate acid-base pairs. ( Definition and you must be able to
identify these pairs in an equation)
HL ONLY Self-ionisation of water; K w; ( Definition and you must be
able to show why [H+] and [OH-] are 1 x 10 -7
pH scale – explain what it is
Use of universal indicator paper or pH meter to measure pH of a
solution
Limitations of the pH scale – usefulness confined to dilute aqueous
solutions.
Calculation of pH of strong acids and bases
HL ONLY Calculation of pH of weak acids and bases.
HL ONLY Theory of acid-base indicators.( what an indicator is, how
they work)
HL ONLY Titration curves.( for each type of acid- base titration you
must be able to draw the curve, and know the end points)
HL ONLY Choice of indicator( for each type of indicator you must
know its colours at various pHs)
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Chapter 29 – Acids, Bases and pH
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