Acids and Bases
Acid/Base Reactions
 Common examples of acid/base
reactions are taking an antacid for an
upset stomach, Using hair conditioner
and spreading lime on soil
 A study of acid/base reactions allows
us understand what goes on in these
everyday events
We can tell how acidic or basic a substance
is by measuring something called pH
 Some acids are dangerous and must
be handled with care Eg. sulphuric
acid
 Others are harmless and can even
form part of our diet Eg. Ethanoic
acid in vinegar and citric acid in
oranges and lemons
 The word acid comes from the Latin word
acidus meaning sour
 Acids turn moist blue litmus paper red
 Acids react with metals to release Hydrogen
gas
Arrhenius’ Theory of Acids and
Bases
 Svante Arrhenius was a Swedish
chemist who put forward a theory
about acids and bases in 1884
 Picture here
 He suggested that some substances
dissociate into ions when added to
water
 Defined an acid as:
An acid is a substance that
dissociates in water to produce
H+ ions
 When Hydrogen chloride (HCl) is
added to water it breaks up
(dissociates) into H+ and Cl- ions
HCl
H+ + Cl-
 The dissociation of Nitric acid and
Sulphuric acid may be presented as
HNO3
H + + NO3-
H2SO4
2H+ + SO42-
And
 Acids like hydrochloric acid and Nitric
acid which donate one H+ are called
monobasic
 Acids like Sulphuric acid which donate
2 H+ are called Dibasic
 An acid like phosphoric acid H3PO4
which donates 3 H+ is called Tribasic
Strong and Weak acids
 Acids such as Hydrochloric acid, Sulphuric
acid and nitric acid are strong acids
because they will fully dissociate in water –
this means almost every molecule will
break up to give H+
 Acids such as ethanoic acids are called
weak acids as they do not fully dissociate
in water
Hydronium ion
 As chemists learned more about
atomic structure it was realised that
H+ could not exist on its own as this
is just a single proton
 What is formed in reality is a
Hydronium ion H3O+ which occurs
when the H+ joins to a water
molecule
The formation of a Hydronium Ion
 Fig 12.3
 The hydronium ion is very common and is
found in every acidic solution
 It is more accurate to represent the
dissociation of acids in water as
HA
H3O+ + AExample
HCL
H3O+ +
Cl-
Importance of water
 If you add substances such as pure
ethanoic acid or pure hydrogen
chloride to solvents such as benzene
or methylbenzene then they can’t
form an acid they will only form an
acid when mixed with water
Bases
 A group of substances that behave
oppositely to acids are called bases
 They turn red litmus paper blue
Sodium hydroxide
 NaOH is a common base found in the
laboratory
 Most bases have a soapy feel as they
turn the oil on your skin to soap
 Arrhenius defined a base as:
A base is a substance that
dissociates in water to produce
OH- ions
 The OH- ion is commonly called the
hydroxide ion
 When NaOH is added to water it
dissociates or breaks up into Na+ and
OH- ions
NaOH
Na+ and OH-
Strong and Weak bases
 A strong base dissociates fully Eg.
NaOH
 A weak base doesn’t dissociate fully
Eg. Calcium hydroxide Ca(OH)2 and
Magnesium hydroxide Mg(OH)2
 NB A base that dissolves in water is
called an alkali
Problems with Arrhenius’ theory
1. It is hydronium ions rather than bare
H+ ions that exist
2. Definitions are only in aqueous
solutions and don’t take account of
solvents such as ammonia, benzene
etc.
3. Not all acid base reactions require
liquid eg Ammonia gas and hydrogen
chloride gas react
Bronsted/Lowry Theory of Acids
and Bases
 1923 Danish chemist Bronsted +
English chemist Lowry independently
proposed new definitions for an acid
and a base
 These simply state that:
 A base is a proton donor
 An acid is a protons aceptor
Consider This
 When hydrogen chloride is added to
water the following reaction occurs
HCl
+
Because
this
donates a
protono it is
a B/L acid
H2O
Because
this
accepts a
proton it is
a B/L base
H3O+ + Cl-
Consider This
 When ammonia gas dissolves in water
the following reaction occurs
NH3 +
Accep
ts a
proto
n B/L
base
H2O
Donat
es a
proto
n B/L
acid
NH4
+
+ OH-
 Water can act as either an acid or a
base and for this reason is known as
amphoteric or amphiprotic
Bronsted Lowry theory does not
only refer to reactions where water
is the solvent
 In the following reaction between
Hydrogen chloride and Ammonia
+ Cl-
HCl + NH3
NH4+
Advantages of the Bronsted/Lowry
Theory over Arrhenius’ Theory
1. Arrhenius’ theory was confined to
aqueous solutions
2. Bronsted/Lowry’s theory broadens
the range of species that can be
defined as acids and bases
3. Substances that are not classified as
acids or bases under Arrhenius’
theory are classified as acids or bases
in the Bronsted/Lowry theory
Note
 An acid will only donate a proton
when there is something there to
accept it
 A base will only accept a proton if
there is something there to donate it
Conjugate acid/base pairs
 Certain reactions are reversible and
can happen in both directions
 We use the symbol
to
show this
Consider This
• If the following reaction is written as
CH3COOH + H2O
CH3COO- +
H3 O +
Ethanoic acid
Water
This means the reaction
CH3COOH + H2O
H3 O +
CH3COO- +
And the reaction
CH3COO- + H3O+
H2 O
Can take place
CH3COOH +
Identify the B/L acid in the
following Equations
 CH3COOH + H2O
CH3COO- + H3O+
 CH3COO- + H3O+
CH3COOH + H2O
 An acid changes to a conjugate base when it
donates a proton
 A base changes to a conjugate acid when it
accepts a proton
•
•
CH3COO- is the conjugate base of
CH3COOH
CH3COOH is the conjugate acid of
CH3COO-
Every base has a conjugate acid and
every acid has a conjugate base
Since CH3COO- and CH3COOH only differ by a
proton we refer to them as a conjugate
acid base pair
 Why can H2O and H3O+ be called a
conjugate acid base pair?
 NB Study examples 12.1 and 12.2 on
pages 140 and 141 of your text book
Neutralisation
• When acids and bases react with each
other in the right proportions they
cancel each other out
• When this happens both the acid and
the base lose their characteristic
properties
• The solution formed is neutral and has
no effect on litmus paper
• This type of reaction is known as
neutralisation
 Neutralisation is the reaction between an acid
and a base to form salt and water
Reaction between HCl and
NaOH
HCl + NaOH
H2O + NaCl
Acid + Base
Water + Salt
 The word salt is a general term
used to describe the substance
formed when the hydrogen in an
acid is replaced by a metal or an
ammonium ion
Understanding what happens
•
HCl + NaOH
H2O + NaCl
•
In this equation the following ions are all present
in solution
H+ + Cl- + Na+ + OHNa+ + Cl- + H2O
The spectator ions are crossed out as what is
important is the reaction where the H+ (Or more
correctly Hydronium ions )and OH- ions react to
form water
H + + OHH2O
Everyday examples of Neutralisation
1. Medicine – Hydrochloric acid in
stomachs helps digestion. Over
eating and stress can produce too
much acid. Antacids such as Alka
Seltzer, Bisodol and Milk of Magnesia
may be used to try to neutralise
excess acid
2. Agriculture – If soil is too acidic
crop yields can be low, therefore
farmers often spread lime (CaO) on
soil to neutralise the acidity
3. Environmental Protection- In
areas affected by acid rain limestone
is often added to lakes to neutralise
acidity
• Toothpaste is slightly basic to neutralise
acids that cause tooth decay
• Baking soda may be used to neutralise
the acid of bees
• Vinegar is used to neutralise the alkaline
sting of wasps
• Shampoo is slightly basic and can cause
scales on hair to stick out, conditioner is
slightly acidic so neutralises the base
flattening the scales and making hair
more shiny and manageable