Relative Strengths of Acids
 The relative strength of various acids can be
compared using Ka.
 The larger the value of Ka, the stronger
the acid
 The structure of an acid plays an important
role in determining the strength of an acid.
Relative Strengths of Acids
 For an acid with the general formula, HX,
the strength of the acid depends on:
 the polarity of the H - X bond
H
X
 the strength of the H - X bond
 the stability of the conjugate base, X-
Relative Strengths of Acids
 Using information about these properties,
several trends related to the relative
strengths of acids can be identified for
binary acids and oxyacids.
 Binary acid:
 an acid with the general formula, HX,
which is generally composed of two
elements (one of which is hydrogen)
 Oxyacid:
 an acid that contains oxygen
Relative Strengths of Acids
Relative Strengths of Binary Acids:
 Within a group, the strength of an acid
increases moving down the group
 HCl is stronger than HF
 Within the same period, the strength
increases as the electronegativity of the
element X increases (i.e. left to right)
 HCl is stronger than H2S
Relative Strengths of Acids
Periodic
Table
Increasing acid strength
Increasing acid strength
For binary acids:
Relative Strengths of Acids
 Relative Strengths of Oxyacids:
 For oxyacids with the same central atom,
acid strength increases as the number of
oxygen atoms attached to the central
atom increases.
 HClO3 is stronger than HClO
Relative Strengths of Acids
 Relative Strengths of Oxyacids:
 For oxyacids with the same number of O
atoms, acid strength increases as the
electronegativity of the central atom
increases
 HClO is stronger than HBrO
In general, electronegativity increases toward
the top within a group and from the left toward
the halogens within a period.
Relative Strengths of Acids
Increasing acid
strength
Increasing acid
strength
Periodic
Table
Halogens
For oxyacids with the same # of oxygens:
Relative Strengths of Acids
Example: Identify the stronger acid in each
of the following pairs:
HClO2 vs. HIO2
H2SeO3 vs. H2SeO4
H2SeO3 vs. HBrO3
H2O vs. HF
Buffers
 Before measuring the pH of an aqueous
solution using a pH meter, chemists must
standardize the pH meter.
 Adjust the reading of the pH meter to the
correct value using standard buffers.
 Buffer:
 A solution that resists a change in pH when
small amounts of acid or base are added
Buffers
 Buffers resist changes in pH because they
contain both:
 An acid
 Neutralizes any OH- ions added
 A base
 Neutralizes any H+ ions added
 At the same time, the acidic and basic
components of a buffer must not react with
each other.
 Generally use a weak conjugate acid-base
pair
Buffers
 Examples of buffers:
 NaC2H3O2 + HC2H3O2
 NH4Cl + NH3
 Lactic acid + sodium lactate
 Two important properties of a buffer:
 pH
 Buffer capacity
Buffers
 Buffer capacity
 the amount of acid or base a buffer can
neutralize before the pH begins to change
appreciably
 The buffer capacity depends on the amount
of acid and base from which the buffer is
prepared.
Buffers
 The pH of a buffer depends on the pKa of
the acid and on the relative concentrations
of the acid and base in the buffer.
pH = pKa + log10
 pKa = - log Ka
[base]
[acid]
HendersonHasselbalch
Equation
Buffers
 The Henderson-Hasselbalch equation can be
used to:
 determine the pH of a particular buffer
OR
 determine the ratio of base to acid
needed to prepare a buffer with a specific
pH.
Buffers
Example: What is the pH of a buffer
containing 0.12 M benzoic acid and 0.20 M
sodium benzoate? Ka = 6.3 x 10-5.
Buffers
Example: What base to acid ratio is needed
to make a pH 4.95 buffer using benzoic acid
and sodium benzoate? Ka = 6.3 x 10-5.
Acid-Base Titrations
 An acid-base titration can be used to
determine the concentration of an acid or
base solution
 titration:
 a technique for determining the
concentration of an unknown solution using
a standard solution
 a solution with a known concentration
Acid-Base Titrations
 The equivalence point in a titration can be
determined using either a pH indicator (ex.
phenolphthalein) or a pH meter.
 Equivalence point:
 the point in the titration where
stoichiometrically equivalent amounts of
base have been added to the acid (or vice
versa)
 the base added has completely reacted
with all available protons (H+)
Acid-Base Titrations
 The general shape for a titration curve for a
monoprotic acid:
pH
pKa
Half
neutralization
point
Equivalence point
Region of high
buffer capacity
mL base added
Acid-Base Titrations
 General Shape for a Diprotic Acid Titration
Curve:
Acid-Base Titrations
 General Shape for a Triprotic Acid Titration
Curve:
Titration of Polyprotic Acids
Example: The titration curve on the next slide
was obtained by titrating 0.250 g of a
polyprotic acid with 0.100 M NaOH. Answer
the following questions.
How many equivalence points are there?
Is the acid monoprotic, diprotic, or triprotic?
What are the values for each pKa?
What volume of NaOH was needed to
neutralize the acid?
 What is the molar mass of the unknown acid?




Titration of Polyprotic acids
Titration of Polyprotic Acid
16
14
12
pH
10
8
6
4
2
0
0
10
20
30
mL 0.100 M NaOH
40
50
Lewis Acids and Bases
 In order for a substance to be a Bronsted-
Lowry base (i.e. a proton acceptor), it must
have an unshared pair of electrons to bind
the proton.
H
H
N
H
H
+
H
Cl
H
N
H, Cl
H
 Similarly, a Bronsted-Lowry acid (i.e. a H+
ion) always gains a pair of electrons during an
acid-base reaction.
Lewis Acids and Bases
 Lewis noticed this trend and proposed a new
definition of acids and bases:
 Lewis Acid:
 An electron pair acceptor
 Lewis Base:
 An electron pair donor
Lewis Acids and Bases
 Examples of Lewis Acids:
 H+
 BF3
 Fe3+
 CO2
 Examples of Lewis Bases:
 OH CH3NH2
Lewis Acids and Bases
Example: Identify the Lewis acid and the
Lewis base in the following reactions.
CH3CH2O
+ CH3Br
CH3CH2OCH3 + Br
O
OH
+
O C O
HO C O