Transcript Acid and Base Equilibria - South Kingstown High School

Acid and Base
Equilibria
– strong acids
and bases
Memorize
Definitions
Arrhenius: acids produce H+ ions in water
and bases produce OH- ions.
 Bronsted-Lowry: an acid is a proton donor
and a base is a proton acceptor.
HC2H3O2 + H2O  C2H3O2- + H3O+

acid
base
base
NH3 + H2O  NH4+ + OHbase
acid
acid
base
acid
The reaction of HCl and H2O. HCl is the acid because it donates
a proton. Water is the base because it accepts a proton.

Species that differ by a proton, like H2O
and H3O+, are called conjugate acidbase pairs
(a) Formic acid transfers a proton to a water molecule.
HCHO2 is the acid and H2O is the base. (b) When a
hydronium ion transfers a proton to the CHO2- ion, H3O+ is
the acid and formate ion is the base.
Conjugate Acid/Base pairs
Acid/Base
HC2H3O2/C2H3O2NH4+/NH3
H3O+/H2O
H2O/OH
H2O is amphoteric because it can act as
either an acid or a base.

An amphoteric substances can act as
either an acid or base
 For
example, the hydrogen carbonate ion:
As an acid :

3

2
3
HCO (aq)  OH (aq)  CO (aq)  H 2O
As a base :

3

HCO (aq)  H 3O (aq)  H 2CO3 (aq)  H 2O
Strong and Weak Acids and Bases
The strength of an acid is a measure of
its ability to transfer a proton
 Acids that dissociate completely with
water (like HCl and HNO3) are classified
as strong
 Acids that are less than completely
ionized are called weak acids
 Bases can be classified in a similar
fashion.

Acetic acid (HC2H3O2) is a weak acid
 It ionizes only slightly in water


HC2 H 3O2 (aq)  H 2O 
H
O
(
aq
)

C
H
O

3
2 3 2 (aq)

weaker acid
we
aker base
stronger acid
stronger base
 The
hydronium ion is a better proton donor
than acetic acid (it is a stronger acid)
 The acetate ion is a better proton acceptor
than water (it is a stronger base)

The position of an acid-base equilibrium
favors the weaker acid and base (the
reactants are favored in this example)
The strengths of the binary acids

increases from left to right within the
same period
 For example, PH3 < H2S < HCl

increase from top to bottom within a
group
 For example, HCl < HBr < HI
Strength of All Acids
1. The POLARITY of the X-H bond

The greater the polarity of the X-H bond the
greater the strength of the acid. Polarity is
measured by the difference in
electronegativity between the bonded atoms.

When an acid dissociates in water, the X-H
bond is broken. The greater the polarity of the
bond, the easier it is to break and produce H+
ions, and thus the stronger the acid.
Explain the difference in the Ka Values
Formula
 HF
 H 2O
 NH3
 CH4

Ka value
7.2 x 10-4
1.8 x 10-16
1 x 10-33
1 x 10-49
DEN
1.8
1.2
0.8
0.4
2. The CHARGE on the acid or base
Compounds becomes less acidic and
more basic as the negative charge
increases.
 It is easier to remove a positive ion (H+)
from a neutral atom or molecule than a
negatively charged one.
 Bases (H+ acceptors) become stronger as
their negative charge increases because
they have a stronger force of attraction for
pulling in extra H+ ions.

Compounds becomes less acidic
and more basic as the negative
charge increases.
Formula
 H3PO4
 H2PO4 HPO42 PO43-
pH
1.5
4.4
9.3
12.0
Oxyacids
When the polarity, size, and charge of two
compounds are all the same (e.g.
oxyacids of the same element) we must
find another way to measure the relative
strengths of these acids.

Trends in oxoacids (acids of hydrogen,
oxygen, and one other element)
Oxyacid

Oxyacid - An acid in which the acid
hydrogen atoms are attached to an
oxygen atom

Examples of oxyacids of the same
element:
H2SO4 and H2SO3
HNO3 and HNO2
3. Oxidation State



As the oxidation state of an atom increases, its
tendency to draw electrons increases.
In an oxyacid, the central atom pulls electrons
away from the oxygen, consequently making the
oxygen more electronegative.
The O-H bond, therefore becomes more polar,
making it easier to form ions and thus increasing
the strength of the acid.
As oxidation state increases so
does the acidity of the oxyacid.
Oxyacid
Ka value
HClO
HClO2
HClO3
HClO4
2.9 x 10-8
1.1 x 10-2
5.0 x 102
1 x 103
Oxidation # of Cl
+1
+3
+5
+7

When the central atom holds the same
number of oxygen atoms, the trend is
the same as for binary acids across a
period, but the reverse for down a
column.
Acid strength: HClO4 > HBrO4 > HIO4
 Acid strength: HClO4 > H2SO4 > H3PO4


For a given central atom, the acid
strength of an oxoacid increases with
the number of oxygens held by the
central atom

Acid strength: H2SO4 > H2SO3
There is a third definition for acid
and bases
It is a further generalization, or
broadening, of the species that can be
classified as either an acid or base
 The definitions are based on electron
pairs and are called Lewis acids and
bases

Definitions
Lewis acid - accepts a pair of electrons to
form a coordinate covalent bond.
 Lewis base – donate a pair of electrons

Cl
H
B 
N: +
H
H
H
Cl
Cl
Cl
N B
H
H
Cl
Cl
NH3 (a Lewis base) forms a coordinate covalent bond with
BF3 (a Lewis acid) during neutralization. NH3BF3 is called an
addition compound because it was made by joining two
smaller molecules.
Carbon dioxide (Lewis acid) reacts with hydroxide
ion (Lewis base) in solution to form the
bicarbonate ion.
The electrons in the coordinate covalent bond
come from the oxygen atom in the hydroxide
ion.

Lewis acids:
 Molecules
or ions with incomplete valence
shells (for example BF3 or H+)
 Molecules or ions with complete valence
shells, but with multiple bonds that can be
shifted to make room for more electrons
(for example CO2)
 Molecules or ions that have central atoms
capable of holding additional electrons
(usually, atoms of elements in Period 3 and
below, for example SO2)

Lewis bases:
 Molecules
or ions that have unshared pairs
of electrons and that have complete shells
(for example O2- or NH3)

All Brønsted acids and bases are Lewis
acids and bases, just like all Arrhenius
acids and bases are Brønsted acids and
bases
Neutral solutions: [H+] = [OH-]
 Acidic solutions: [H+] > [OH-]
 Basic solutions: [H+] < [OH-]
 To make the comparison of small values
of [H+] easier, the pH was defined:




pH  oflog[
In terms
theH
pH:] or [ H ]  10
Neutral solutions: pH = 7.00
 Acidic solutions:
pH < 7.00
 Basic solutions:
pH > 7.00

-pH
The pH of some
common solutions.
[H+] decreases,
while [OH-]
increases, from top
to bottom.
The pH of a solution can be measured
with a pH meter or estimated using a
visual acid-base indicator
 An acid-base indicator is a species that
changes color based on the pH
 Calculating the pH of a strong acid or
base is “easy” because they are 100%
dissociated in aqueous


For example, the pH of 0.10 M HCl is 1.00 and
the pH of 0.10 M NaOH is 13.00

In the last example it was assumed that
the total concentration of [H+] was due to
the strong acid (HCl) and [OH-] was due
to the strong base (NaOH)

This assumption is valid because the
autoionization of water is suppressed in
strongly acidic or strongly basic solutions
 This
assumption fails for very dilute solutions
of acids or bases (less than 10-6 M)
Equilibrium Constant Expression

The equilibrium or ionization constants
for weak acids and bases and water
are given the labels of Ka, Kb and Kw.
1)
A weak acid:
HA + H2O  H3O+ + A-
Ka = [H+][A-]
[HA]
Equilibrium Constant Expressions
2.
A weak base:
B + H2O  BH+ + OHKb = [BH+][OH-]
[B]
3.
Water dissociation:
H20  H+ + OHKw = [H+][OH-]
(1 x 10-7)(1x 10-7)= 1 x 10-14
Equilibrium Constant
Relationships
The product of the Ka and Kb for an acid
and its conjugate base is the Kw
KaKb = Kw = 1 x 10-14
 The greater the Ka or Kb the greater the
dissociation of the acid or base
• Ka and Kb values are usually very small
since they are weak acids and bases.
• pKa and pKb are used to show the
equilibrium concentrations.

For theweak acid :
HA  H 2O H 3O   A
[ H 3O  ][ A ]
Ka 
[ HA]
for theconjugatebase :
A  H 2O HA  OH 
[ HA][OH  ]
Kb 
[ A ]
the product is
[ H 3O  ][ A ] [ HA][OH  ]
K a  Kb 

[ HA]
[ A ]
 [ H 3O  ][OH  ]  K w
Ionization Constant
Relationships
pKb = -log Kb
 pKa = -log Ka
 pKw = -log Kw = -log(1 x 10-14)
pKw = 14
So, pKa + pKb = 14
And, pH + pOH = 14
 If you know ka for a weak acid you can always
find kb for its conjugate base and visa-versa.


Example: Morphine is very effective at
relieving intense pain and is a weak base.
What is the Kb, pKb, and percentage
ionization of morphine if a 0.010 M solution
has a pH of 10.10?



B(aq)  H 2O BH (aq)  OH
I
0.010
C
-x
E 0.010- x

0
0
x
x
x
x
[ BH  ][OH  ]
Kb 
[ B]
x2

0.010- x
At equilibrium, [OH-] = x = 10-pOH
SOLUTION: Use pOH = 14.00 – pH, substituting:
[OH  ]  10 (14.00 10.10 )
 1.3 10 4 M , then
x2
(1.3 10 4 ) 2
Kb 

4
0.010 x (0.010 1.3  10 )
 1.6 106 so pK b  5.80,and
x
% ionization
 100%
0.010
 1. 3%
Ionization Constant Calculations
1.
A monoprotic acid solution has a concentration
of 0.100 M and the pH is 2.44 @ 25oC.
Calculate the Ka and pKa.
Ka = 1.36 x 10-4, pKa = 3.86
2.
Calculate the pH of a 0.010 M solution of HCl.
pH = 2.0
3.
Hydrazine, N2H4 has a concentration of 0.025
M. Calculate the pH and % ionization. The Kb
= 1.7 x 10-5
pH = 10.81, % ionization = 2.6 %