Chemistry 40S
Unit 5: Acids & Bases
Lesson 1
Learning Outcomes
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C12-5-01: Outline the historical development
of acid-base theories. Include: the Arrhenius,
Brønsted-Lowry, and Lewis theories
C12-5-02: Write balanced acid-base
chemical equations. Include: conjugate acidbase pairs and amphoteric behaviour
C12-5-03: Describe the relationship between
the hydronium and hydroxide ion
concentrations in water. Include: the ion
product of water, Kw
Arrhenius Acid & Base Theory

Arrhenius proposed that electrolytes break up into charged
particles in water.
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Arrhenius later discovered that all acidic and basic
solutions he tested were electrolytes
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Eventually as the knowledge of atomic structure increased,
Arrhenius' theory was accepted.
Determined that acids and bases must ionize or dissociate in
water.
Arrhenius Acid – a substance which releases hydrogen ions
or increases the hydrogen ion concentration when
dissolved in water.
Arrhenius Base - a substance which releases hydroxide ions
or increases the hydroxide ion concentration when
dissolved in water
Drawbacks to Arrhenius’
Theory

The Arrhenius definition says that acids and bases
can only occur in water solutions. E.g. hydrogen
chloride gas & ammonia gas will react together
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
E.g. baking soda in water turns litmus blue, but has no
apparent hydroxide ion & metal ions turn litmus red,
but have no hydrogen ions.
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There are many substances which are acidic or basic
but do not have a hydrogen ion or a hydroxide ion
The Arrhenius definition does not account for the acidity
and basicity of these examples.
Arrhenius Theory---not all bad. It was important in
establishing the concept of dissociation and
explaining the process of neutralization
BrØnsted-Lowry Theory
 BrØnsted
& Lowry independently
developed a more general definition of
acids and bases within months of each
other

Their theories accounted for the acidity of
metal ions AND the basicity of baking soda.
 BrØnsted-Lowry
Acid – proton (H+) donor
 BrØnsted-Lowry Base – proton (H+)
acceptor
BrØnsted-Lowry Theory
 Acids
don’t just throw off their protons in
solution
 Bases however have a strong affinity for
protons
 If the attraction between the base and
the proton is greater than the bond
between the proton and the rest of the
acid molecule then the proton will be
donated
Amphoteric Substances
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Amphoteric Substances – act like acids and
bases substances like water can act as BOTH
an acid and a base.
E.g. water, hydrogen carbonate ion (HCO3)
Hydrogen carbonate can act as an acid by
losing it’s hydrogen in water:
HCO3– + H2O → CO32– + H3O+
Or as a base by stealing a proton from water:
HCO3– + H2O → H2CO3 + OH–
Conjugate Pairs
 The
general form of a Brønsted-Lowry
acid-base reaction is:
Acid + Base ⇌ Conjugate Acid + Conjugate Base
 The
conjugate acid is what remains after
a base has accepted a proton and the
conjugate base is what remains after the
acid has donated its proton.
Conjugate Pairs
Drawbacks of BrØnsted-Lowry
Theory

The Brønsted-Lowry Theory expands the
definition of an acids and bases to occur
without water


BUT there is still a requirement for the presence
of a solvent
The Brønsted-Lowry Theory explains why
substances can act as an acid and a base
and how substances without hydroxides can
act as bases

BUT it does not explain how substances without
protons can act as acids.
Lewis Theory


Aqueous metal ions with charges of 2+ and 3+,
except group 2 metals, are acidic in water
Metal ions act as Lewis acids in water:
Fe3+ + 6 H2O → Fe(H2O)63+

Because of the large positive charge and the
polar nature of water, the oxygen side of the
water (δ-) is attracted to the highly positive metal
ion and forms a covalent bond with 6 water
molecules.
Lewis Theory
 Lewis

E.g. Fe3+
 Lewis

Acid – accepts an electron pair
Base – donates an electron pair
E.g. H2O
Drawbacks of Lewis Theory

Lewis definition is so general that any reaction in
which a pair of electrons are transferred becomes an
acid-base reaction


Includes many reactions that would not be included
with the Brønsted-Lowry definition
The Lewis Theory works with the Brønsted-Lowry
Theory because any Brønsted-Lowry base must have
a pair of non-bonding electrons in order to accept a
proton
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BUT Lewis Theory vastly expands the number of acids
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
Any compound that has one or more valence shell
orbitals can now act as an acid
Reactions can be difficult to deal with according to
Lewis Theory
Ionization of Water


Water is amphoteric (can act as both an acid and a base)
HA + H2O(l) ⇌ H3O+(aq) + A–(aq)
or
B + H2O(l) ⇌ BH+(aq) + OH–(aq)
Pure water actually dissociates into ions, or ionizes, slightly
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Self-ionization or autoionization.
The equation for self-ionization is written as:
H2O(l) + H2O(l) ⇌ H3O+(aq) + OH–(aq)
or
+
H2O(l) ⇌ H (aq) + OH–(aq)
An equilibrium is established between hydronium and
hydroxide ions.
Ion Product of Water
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If an equilibrium is established between hydronium
ions, hydroxide ions and water molecules, an
equilibrium law can be written:
+
−
𝑯𝟑𝑶 [𝑶𝑯 ]
𝑲𝒄 =
𝑯𝟐𝑶 [𝑯𝟐𝑶]
Since water is a liquid, the product of KC and
water results in the ion product for water, Kw. The
equilibrium law for water becomes:
Kw = [H3O+][OH–]
At 25°C, the concentration of the hydronium and
hydroxide ions are equal at 1.0 x 10–7 mol/L.

Therefore, at 25°C, the value of Kw is 1.0 x 10–14.
Adding Acid to Water
Le Chatelier's Principle can be used to predict
the effect of dissolving an acid or base on
hydronium and hydroxide ion concentrations
Adding Acid
 Acid produces a large amount of H3O+ ions
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[H3O+] increases  equilibrium will shift to the left
to use up some of the added hydronium and
maintain Kw at 1.0 x 10–14.
Since equilibrium shifts left, the [OH-] is reduced
Adding a strong acid to water increases the
[H3O+] and reduces the [OH-]concentration.
Adding Base to Water
Adding Base
 When a base is dissolved in water, the hydroxide
ion concentration increases
 Le Chatelier's Principle  the equilibrium shifts left
to use up some of the added hydroxide and
maintain Kw at 1.0 x 10–14.
 Equilibrium shifts left  [OH-] is reduced
 Adding a strong base to water increases
[OH-] and reduces [H3O+]
 NOTE: hydronium ions AND hydroxide ions are
BOTH present in any solution - whether they are
acidic or basic.
Calculating Hydroxide
Concentration
Example 1: If 2.5 moles of hydrochloric acid
is dissolved in 5.0 L of water, what is the
concentration of the hydroxide ions?
Assume the volume remains unchanged.
Calculating Hydronium
Concentration
Example 2: 0.40 g of NaOH is dissolved in
water to make a solution with a volume of
1.0 L. What is the hydronium ion
concentration in this solution?