1. No category

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
acids 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
 First, assume that the ionization of the
acid and base in the buffer is negligible.
 Use the initial concentrations of the
acid and base in the H-H equation.
pH = pKa + log [base]
[acid]
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.
Buffers
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
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 BronstedLowry 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