IB Chemistry Power Points
Topic 08
Acids and Bases
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Introduction to
Acids and Bases
In aqueous solutions, a proportion of the water
molecules dissociate;
 The ions formed are H+ or positively charged
hydrogen ions and negatively charged hydroxide ions
(OH-)
Technically
+
2 H2O(l)  H3O
+
−
-14
Kw = [H ][OH ] = 1 x 10
(aq)
-
+ OH (aq)
 Some chemical compounds contribute additional H+
to make the solution more acidic. Other compounds
remove H+ ions.
 A compound that increases H+ is called an acid
 Examples: HCl, H2SO4, HNO3, CH3COOH
 A compound that removes H+ ions from an aqueous
solution is called a base. This reaction is called a
neutralization.
 Often this is done by adding OH- ions for example
NaOH, KOH, Ca(OH)2. Soluble bases are called
alkalis.
Neutralization Reactions
 hydroxides
acid + base  water + salt
HCl + NaOH  H2O + NaCl (aq)
• metal oxides
acid + base  water + salt
2 HCl + Cu2O  H2O + CuCl2 (aq)
• ammonia
acid + base  salt
HCl + NH3  NH4Cl (aq)
Three theories of acids
Arrhenius (most common)
Bronsted-Lowry
Lewis
Arrhenius (most common):
+
an acid dissociates to yield H
and a base dissociates to yield OH
Hydrochloric acid
Sodium hydroxide
+
-
H + Cl
+
-
Na + OH
Bronsted-Lowry:
+
an acid is a proton (H ) donor
and a base is a proton acceptor
amphiprotic
Lewis:
An acid is an electron pair acceptor
and a base is an electron pair donor
A dative covalent bond is formed
Example of Lewis Acid
Lewis
Acid
Lewis
Base
This is a common example that is not an obvious acid/base rxn
Boron trifluoride acts as a Lewis Acid. The boron has only
6 electron in valence shell so the lone pair of electrons
forms a dative bond and fills up the valence shell of the
boron
Indicators
Acids and bases are substances with
specific physical and chemical
properties.
We can determine if substances are
acidic or basic by testing their reaction
with indicators.
Indicators are organic substances that
change color in the presence of an acid or
a base.
Some common indicators
in acid
Litmus
red
Phenolphthalein
colorless
Methyl orange
red
in base
blue
pink
yellow
Reactions of acids
 React with active metals (above copper in reactivity
series)
2 HCl + Ca  CaCl2 + H2
 Reaction with carbonates
H2SO4 + Na2CO3 Na2SO4 + CO2 + H2O
 Reaction with bicarbonates
HNO3 + NaHCO3 NaNO3 + CO2 + H2O
Acid/base properties of Period 3 oxides (topic 3)
 Metal oxides Na2O and MgO react with water to form
hydroxides (basic solutions)
Na2O + H2O  2 NaOH (aq)
 Aluminum oxide is amphoteric (will react as a base
with an acid or vice versa)
Al2O3 + 6 HCl  2 AlCl3 + 3 H2O
 Other period 3 oxides (non-metal S, P, Cl oxides) react
with water to form acidic solutions
SO3 + H2O H2SO4 (aq)
see page 15 in study guide
Acid/base properties of Period 3 chlorides (topic 13)
 Chlorides across Period 3 become more acidic across
the period
NaCl (aq) is neutral
MgCl2 (aq) is weakly acidic
Chlorides of Al, Si, P, S and Cl2 react with
water to produce HCl (aq) solutions
see Study guide page 16
Strong Acids vs Weak Acids
The strength of an acid or base depends on how
easily it dissociates in water.
The dissociation of an acid or base is an equilibrium.
+

HA(aq) H (aq) + A (aq)
+

BOH(aq) B (aq) + OH (aq)
Strong acids or bases dissociate (ionize) easily –
the equilibrium favors the ionic products : kc >> 1
Strong vs Weak
When the strength of an acid or base is discussed, it is
very important NOT to confuse “strength” with
“concentration”
3
A 5M acid solution contains 5 mol of acid per dm but its
strength is determined by how much of that acid is
ionized.
Strong acids : HCl, H2SO4, HNO3 (mono vs diprotic)
Strong bases : NaOH, KOH, Ba(OH) 2
Weak acids: CH3COOH, H2CO3, carbonic acid CO2(aq)
Weak bases: NH3, ethylamine CH3CH2NH2
Strong vs Weak
How to tell
Strong acids and bases are mostly ionized and therefore
solutions are good electrolytes (high conductivity). The pH of
the solution can also be measured.
What is the pH scale?
 pH is a measurement of hydrogen ion concentration
 It tells you how acids or basic (or alkaline) something is
 Ranges from 0 (most acidic) to 14 (most basic

pH   log[ H ]
How does scale work?
 The scale is logarithmic. As you go up
or down, the concentration is changed
by a power of ten
 Example pH 3 is 100 times more
concentrated than pH 5
 neutral
 pH 10 is 100 times less concentrated
than pH 8
Strong Acid
example HCl
HCl(aq)
+
H
(aq)
+
Cl (aq)
[H+ ][Cl- ]
k=
>> 1
[HCl]
•
•
•
•
•
completely dissociated
pH of 0.1 M soln = 1
strong electrolyte
reacts vigorously
note simplified “net ionic”
equation
Weak Acid
example CH3COOH
+
CH3COOH (aq)  H
(aq)
+
CH3COO (aq)
[H+ ][CH3COO- ]
k=
<< 1
[CH3COOH]
•
•
•
•
partially dissociated
pH of 0.1 M soln = 2.9
weak electrolyte
reacts slowly