ch18-Acid Base Equilibrium-3-26-18-WO Prob_lecture_6e_final.ppt

Acid base equilibrium
Acids and Bases in Water
Autoionization of Water and the pH Scale
18.3 Proton Transfer and the Brønsted-Lowry Acid-Base Definition
18.4 Solving Problems Involving Weak-Acid Equilibria

1. 18-1
Acid-Base Equilibria
18.1 Acids and Bases in Water
18.2 Autoionization of Water and the pH Scale
18.3 Proton Transfer and the Brønsted-Lowry Acid-Base Definition
18.4 Solving Problems Involving Weak-Acid Equilibria
18.5 Weak Bases and Their Relation to Weak Acids
18.6 Molecular Properties and Acid Strength
18.7 Acid-Base Properties of Salt Solutions
18.8 Generalizing the Brønsted-Lowry Concept: The Leveling
Effect
18.9 Electron-Pair Donation and the Lewis Acid-Base Definition

2. 18-2
Table 18.1Some Common Acids and Bases and their
Household Uses.

3. 18-3
Arrhenius Acid-Base Definition
An acidis a substance with H in its formula that dissociates
to yield H
3O
+
.
This is the earliest acid-base definition, which classifies
these substances in terms of their behavior in water.
A base is a substance with OH in its formula that
dissociates to yield OH
-
.
When an acid reacts with a base, they undergo
neutralization:
H
+
(aq) + OH
-
(aq) → H
2O(l) DH°
rxn= -55.9 kJ

4. 18-4
Brønsted-Lowry Acid-Base Definition
An acid is a proton donor, any species that donates an
H
+
ion.
•An acid must contain H in its formula.
A base is a proton acceptor, any species that accepts
an H
+
ion.
•A base must contain a lone pair of electrons to bond
to H
+
.
An acid-base reaction is a proton-transfer process.

5. 18-5
Figure 18.7Dissolving of an acid or base in water as a Brønsted-
Lowry acid-base reaction.
(acid, H
+
donor)
(base, H
+
acceptor)
Lone pair
binds H
+
(base, H
+
acceptor)(acid, H
+
donor)
Lone pair
binds H
+

6. 18-6
Conjugate Acid-Base Pairs
H
2S+ NH
3 HS
-
+ NH
4
+
NH
3acceptsa H
+
to form NH
4
+
.
In the forwardreaction:
In the reversereaction:
H
2S+ NH
3 HS
-
+ NH
4
+
HS
-
acceptsa H
+
to form H
2S.
H
2S donatesa H
+
to form HS
-
.
NH
4
+
donatesa H
+
to form NH
3.

7. 18-7
Conjugate Acid-Base Pairs
H
2S+ NH
3 HS
-
+ NH
4
+
H
2S and HS
-
are a conjugate acid-base pair:
HS
-
is the conjugate baseof the acid H
2S.
A Brønsted-Lowry acid-base reaction occurs when an
acid and a base react to form their conjugate base
and conjugate acid, respectively.
NH
3and NH
4
+
are a conjugate acid-base pair:
NH
4
+
is the conjugate acidof the base NH
3.
acid
1+ base
2 base
1+ acid
2

8. 18-8
Base Acid+Acid Base+
Conjugate Pair
Conjugate Pair
Table 18.4 The Conjugate Pairs in some Acid-Base Reactions
Reaction 4 H
2PO
4
-
OH
-
+
Reaction 5 H
2SO
4 N
2H
5
+
+
Reaction 6 HPO
4
2-
SO
3
2-
+
Reaction 1 HF H
2O+ F
-
H
3O
+
+
Reaction 3 NH
4
+
CO
3
2-
+
Reaction 2 HCOOH CN
-
+ HCOO
-
HCN+
NH
3 HCO
3
-
+
HPO
4
2-
H
2O+
HSO
4
-
N
2H
6
2+
+
PO
4
3-
HSO
3
-
+

9. 18-9
Net Direction of Reaction
The netdirection of an acid-base reaction depends on
the relativestrength of the acids and bases involved.
A reaction will favor the formation of the weakeracid
and base.
H
2S + NH
3 HS
-
+ NH
4
+
stronger base
weaker basestronger acid
weaker acid
This reaction favors the formation of the products.

10. 18-10
Figure 18.8Strengths of conjugate acid-base pairs.
The stronger the acid is, the
weaker its conjugate base.
When an acid reacts with a
base that is farther down the
list, the reaction proceeds to
the right(K
c> 1).

11. 18-11
Strong and Weak Acids
A strongacid dissociates completelyinto ions in water:
HA(gor l) + H
2O(l) → H
3O
+
(aq) + A
-
(aq)
A dilute solution of a strongacid contains no HA molecules.
A weakacid dissociates slightlyto form ions in water:
HA(aq) + H
2O(l) H
3O
+
(aq) + A
-
(aq)
In a dilute solution of a weakacid, most HA molecules are
undissociated.
[H
3O
+
][A
-
]
[HA][H
2O]
K
c= has a very smallvalue.

12. 18-12
Strong acid: HA(gor l) + H
2O(l) → H
3O
+
(aq) + A
-
(aq)
Figure 18.1AThe extent of dissociation for strong acids.
There are no HA molecules in solution.

13. 18-13
Figure 18.1BThe extent of dissociation for weak acids.
Weak acid: HA(aq) + H
2O(l) H
3O
+
(aq) + A
-
(aq)
Most HA molecules are undissociated.

14. 18-14
Figure 18.2 Reaction of zinc with a strong acid (left) and a
weak acid (right).
1 MHCl(aq) 1 MCH
3COOH(aq)
Zinc reacts rapidly with
the strong acid, since
[H
3O
+
] is much higher.

15. 18-15
The Acid Dissociation Constant, K
a
[H
3O
+
][A
-
]
[HA][H
2O]
K
c= K
c[H
2O] = K
a=
[H
3O
+
][A
-
]
[HA]
The value of K
ais an indication of acid strength.
Stronger acid larger K
ahigher [H
3O
+
]
Weaker acid smaller K
alower % dissociation of HA
HA(aq) + H
2O(l) H
3O
+
(aq) + A
-
(aq)

16. 18-16
Table 18.2 K
aValues for some Monoprotic Acids at 25°C

17. 18-17
Classifying the Relative Strengths of Acids
•Strongacidsinclude
–the hydrohalic acids (HCl, HBr, and HI) and
–oxoacids in which the number of O atoms exceeds the number
of ionizable protons by two or more (eg., HNO
3, H
2SO
4, HClO
4.)
•Weak acidsinclude
–the hydrohalic acid HF,
–acids in which H is not bonded to O or to a halogen (eg., HCN),
–oxoacids in which the number of O atoms equals or exceeds
the number of ionizable protons by one (eg., HClO, HNO
2), and
–carboxylic acids, which have the general formula RCOOH(eg.,
CH
3COOHand C
6H
5COOH.)

18. 18-18
Classifying the Relative Strengths of Bases
•Strongbasesinclude
–water-soluble compounds containing O
2-
or OH
-
ions.
–The cationsare usually those of the most active metals:
•M
2O or MOH, where M = Group 1A(1) metal (Li, Na, K, Rb, Cs)
•MO or M(OH)
2where M = group 2A(2) metal (Ca, Sr, Ba).
•Weak bases include
–ammonia (NH
3),
–amines, which have the general formula
–The common structural feature is an N atom with a lone
electron pair.

19. 18-19
Autoionization of Water
Water dissociates very slightly into ions in an equilibrium
process known as autoionizationor self-ionization.
2H
2O (l) H
3O
+
(aq) +OH
-
(aq)

20. 18-20
The Ion-Product Constant for Water (K
w)
2H
2O (l) H
3O
+
(aq) +OH
-
(aq)
[H
3O
+
][A
-
]
[H
2O]
2
K
c=
K
c[H
2O]
2
= K
w= [H
3O
+
][OH
-
] = 1.0x10
-14
(at 25°C)
= 1.0x10
-7
(at 25°C)
In pure water,
[H
3O
+
] = [OH
-
] =
Both ionsare present in all aqueous systems.

21. 18-21
Higher [H
3O
+
] lower [OH
-
]
Higher [OH
-
] lower [H
3O
+
]
A change in [H
3O
+
] causes an inverse change in [OH
-
],
and vice versa.
We can define the terms “acidic” and “basic” in terms of
the relative concentrations of H
3O
+
and OH
-
ions:
In an acidic solution,[H
3O
+
] > [OH
-
]
In a neutral solution,[H
3O
+
] = [OH
-
]
In basic solution,[H
3O
+
] < [OH
-
]

22. 18-22
Figure 18.3The relationship between [H
3O
+
] and [OH
-
] and the
relative acidity of solutions.

23. 18-23
The pH Scale
pH = -log[H
3O
+
]
The pH of a solution indicates its relative acidity:
In an acidic solution,pH < 7.00
In a neutral solution,pH = 7.00
In basic solution,pH > 7.00
The higherthe pH, the lowerthe [H
3O
+
] and the less
acidicthe solution.

24. 18-24
Figure 18.4
The pH values of some
familiar aqueous
solutions.
pH = -log [H
3O
+
]

25. 18-25
Table 18.3 The Relationship between K
aand pK
a
Acid Name (Formula) K
aat 25°C pK
a
Hydrogen sulfate ion (HSO
4
-
) 1.0x10
-2
1.99
Nitrous acid (HNO
2) 7.1x10
-4
3.15
Acetic acid (CH
3COOH) 1.8x10
-5
4.75
Hypobromousacid (HBrO) 2.3x10
-9
8.64
Phenol (C
6H
5OH) 1.0x10
-10
10.00
pK
a= -logK
a
A low pK
acorresponds to a high K
a.

26. 18-26
pH, pOH, and pK
w
pH = -log[H
3O
+
]
pOH = -log[OH
-
]
pH + pOH = pK
w for any aqueous solution at any temperature.
pK
w= pH + pOH = 14.00 at 25°C
K
w= [H
3O
+
][OH
-
] = 1.0x10
-14
at 25°C
Since K
wis a constant, the values of pH, pOH, [H
3O
+
],
and [OH
-
] are interrelated:
•If [H
3O
+
] increases, [OH
-
] decreases (and vice versa).
•If pH increases, pOH decreases (and vice versa).

27. 18-27
Figure 18.5The relations among [H
3O
+
], pH, [OH
-
], and pOH.

28. 18-28
Figure 18.6Methods for measuring the pH of an aqueous solution.
pH paper
pH meter

29. 18-29
Sample Problem 18.6 Using Molecular Scenes to Predict the Net
Direction of an Acid-Base Reaction
PROBLEM:Given that 0.10 Mof HX (blue and green) has a pH of 2.88,
and 0.10 MHY (blue and orange) has a pH 3.52, which
scene best represents the final mixture after equimolar
solutions of HX and Y
-
are mixed?
PLAN:A stronger acid and base yield a weaker acid and base, so we
have to determine the relative acid strengths of HX and HY to
choose the correct molecular scene. The concentrations of the
acid solutions are equal, so we can recognize the stronger acid
by comparing the pH values of the two solutions.

30. 18-30
Sample Problem 18.6
SOLUTION:
The HX solution has a lower pH than the HY solution, so HX is the
stronger acid and Y
-
is the stronger base. The reaction of HX and Y
-
has a K
c> 1, which means the equilibrium mixture will contain more
HY than HX.
Scene 1 has equal numbers of HX and HY, which could occur if the
acids were of equal strength. Scene 2 shows fewer HY than HX,
which would occur if HY were the stronger acid.
Scene 3 is consistent with the relative acid strengths, because it
contains more HY than HX.

31. 18-31
Solving Problems Involving
Weak-Acid Equilibria
Problem-solving approach
1.Write a balanced equation.
2.Write an expression for K
a.
3.Define xas the change in concentration that
occurs during the reaction.
4.Construct a reaction table in terms of x.
5.Make assumptions that simplify the calculation.
6.Substitute values into the K
aexpression and
solve for x.
7.Check that the assumptions are justified.

32. 18-32
Solving Problems Involving
Weak-Acid Equilibria
The notation system
•Molar concentrations are indicated by [ ].
•A bracketed formula with no subscript
indicates an equilibrium concentration.
The assumptions
•[H
3O
+
] from the autoionizationof H
2O
is negligible.
•A weak acid has a small K
aand its
dissociation is negligible. [HA] ≈ [HA]
init.

33. 18-33
Concentration and Extent of Dissociation
Percent HA dissociated =
[HA]
dissoc
[HA]
init
x 100
As the initial acid concentration decreases, the percent
dissociation of the acid increases.
HA(aq) + H
2O(l) H
3O
+
(aq) + A
-
(aq)
A decreasein [HA]
initmeans a
decreasein [HA]
dissoc= [H
3O
+
] = [A
-
],
causing a shift towards the products.
The fractionof ions present increases, even though the
actual [HA]
dissocdecreases.

34. 18-34
A polyprotic acidis an acid with more than one ionizable
proton. In solution, each dissociation step has a different
value for K
a:
K
a1> K
a2> K
a3
Polyprotic Acids
K
a1=
[H
3O
+
][H
2PO
4
-
]
[H
3PO
4]
= 7.2x10
-3H
3PO
4(aq) + H
2O(l) H
2PO
4
-
(aq) + H
3O
+
(aq)
K
a2=
[H
3O
+
][HPO
4
2-
]
[H
2PO
4
-
]
= 6.3x10
-8
H
2PO
4
-
(aq) + H
2O(l) HPO
4
2-
(aq) + H
3O
+
(aq)
K
a3=
[H
3O
+
][PO
4
3-
]
[HPO
4
2-
]
= 4.2x10
-13HPO
4
2-
(aq) + H
2O(l) PO
4
3-
(aq) + H
3O
+
(aq)
We usually neglect [H
3O
+
] produced after the first dissociation.

35. 18-35
Table 18.5 Successive K
avalues for Some Polyprotic Acids at
25°C

36. 18-36
Weak Bases
A Brønsted-Lowry base is a species that accepts an H
+
.
For a weak base that dissolves in water:
B(aq) + H
2O(l) BH
+
(aq) + OH
-
(aq)
The base-dissociation or base-ionizationconstantis
given by:
[BH
+
][OH
-
]
[B]
K
b=
Note that no base actually dissociatesin solution, but ionsare
produced when the base reacts with H
2O.

37. 18-37
Figure 18.9Abstraction of a proton from water by the base
methylamine.
Lone pair of N
pair binds H
+

38. 18-38
Table 18.6 K
bValues for Some Molecular (Amine) Bases at 25°C

39. 18-39
Anions of Weak Acids as Weak Bases
The anionsof weak acids often function as weak bases.
A
-
(aq) + H
2O(l) HA(aq) + OH
-
(aq)
K
b=
[HA][OH
-
]
[A
-
]
A solution of HA is acidic, while a solution of A
-
is basic.
HF(aq) + H
2O(l) H
3O
+
(aq) + F
-
(aq)
HF is a weak acid, so this equilibrium lies to the left.
[HF] >> [F
-
], and [H
3O
+
]
from HF>> [OH
-
] ;
the solution is therefore acidic.
from H
2O

40. 18-40
F
-
(aq) + H
2O(l) HF(aq) + OH
-
(aq)
If NaF is dissolved in H
2O, it dissolves completely, and F
-
can act as a weak base:
HF is a weak acid, so this equilibrium also lies to the left.
[F
-
] >> [HF], and [OH
-
] >> [H
3O
+
] ;
the solution is therefore basic.
from H
2Ofrom F
-

41. 18-41
K
aand K
bfor a Conjugate Acid-Base Pair
HA + H
2O H
3O
+
+ A
-
A
-
+ H
2O HA + OH
-
2H
2O H
3O
+
+ OH
-
= [H
3O
+
][OH
-
]
K
cfor the overall equation = K
1x K
2,so
[H
3O
+
][A
-
]
[HA]
[HA][OH
-
]
[A
-
]
x
K
a x K
b = K
w
This relationship is true for any conjugate acid-base pair.

42. 18-42
Acid Strength of Nonmetal Hydrides
For nonmetal hydrides (E-H), acid strength depends on:
•the electronegativity of the central nonmetal (E), and
•the strength of the E-H bond.
Acrossa period, acid strength increases.
Electronegativity increases across a period, so the acidity of E-H
increases.
Downa group, acid strength increases.
The length of the E-H bond increases down a group and its bond
strength therefore decreases.

43. 18-43
6A(16)
H
2O
H
2S
H
2Se
H
2Te
7A(17)
HF
HCl
HBr
HI
Electronegativity
increases, so
acidity increases
Bond strength decreases, so acidity increases
Figure 18.10The effect of atomic and molecular properties on
nonmetal hydride acidity.

44. 18-44
Acid Strength of Oxoacids
All oxoacids have the acidic H bonded to an O atom.
Acid strength of oxoacids depends on:
•the electronegativity of the central nonmetal (E), and
•the number of O atoms around E.
For oxoacids with the samenumber of O atoms, acid
strength increases as the electronegativity of E increases.
For oxoacids with differentnumbers of O atoms, acid
strength increases with the number of O atoms.

45. 18-45
Figure 18.11The relative strengths of oxoacids.
A
Electronegativity increases, so acidity increases.
B
Number of O atoms increases, so acidity increases.

46. 18-46
If a salt that consists of the anionof a weak acidand the
cationof a weak base, the pH of the solution will
depend on the relative acid strength or base strength of
the ions.
NH
4CN
NH
4+ is the cation of a
weak base, NH
3.
CN
-
is the anion of a
weak acid, HCN.
Salts of Weakly Acidic Cations and
Weakly Basic Anions
CN
-
(aq) + H
2O(l) HCN(aq) + OH
-
(aq)
NH
4
+
(aq) + H
2O(l) NH
3(aq) + H
3O
+
(aq)

47. 18-47
K
aof NH
4
+
=
K
w
K
bof NH
3
1.0x10
-14
1.76x10
-5
= = 5.7x10
-10
K
bof CN
-
=
K
w
K
aof HCN
1.0x10
-14
6.2x10
-10
= = 1.6x10
-5
The reaction that proceeds farther to the right determines the
pH of the solution, so we need to compare the K
aof NH
4
+
with the K
bof CN
-
.
Since K
bof CN
-
> K
aof NH
4
+
, CN
-
is a stronger base than
NH
4
+
is an acid. A solution of NH
4CN will be basic.

48. 18-48
Table 18.7 The Acid-Base Behavior of Salts in Water

49. 18-49
14.6: Buffers
•Buffer –A solution containingappreciable
amounts of both a weak acid and its
conjugate base (or a weak base and its
conjugate acid).
–Buffer solutions resist changes in pH upon addition of a
small amount of strong acid or base.
–Examples:

50. 18-50
pH 8
No Buffer
pH 8
With Buffer

51. 18-51
How Buffers Work
•Consider the buffer: HA/A
-
•Buffers are good at getting rid of strong acids and bases.
–Added OH
-
or H
3O
+
are consumed and therefore do not
directly affect the pH.
–Any added OH
-
, is converted into the weak base, A
-
HA (aq) + OH
-
(aq) →A
-
(aq) + H
2O
–Any added H
3O
+
, is converted to the weak acid, HA
A
-
(aq) + H
3O
+
(aq) →HA (aq) + H
2O
Therefore buffers can be used to maintain a nearly
constant pH!

52. 18-52
•CH
3COOH/ NaCH
3CO
2Buffer
How Buffers Work
•The change in the amounts of CH
3COOH and CH
3CO
2
-
does result in a small change in pH.

53. 18-53
Buffer Capacity
•Buffer Capacity –Amount of strong acid or
base that can be added to a given volume of a
buffer solution before the pH changes
significantly.
•All buffers have a limited capacity of how
much H
3O
+
or OH
-
they can “soak up”
–Eventually, all the HA reacts with the added OH
-
–Eventually, all the A
-
reacts with the added H
3O
+

54. 18-54
Buffer Capacity
•Buffer capacity depends on the number of
moles of the weak acid and its conjugate base
that are in the mixture.
•More moles of buffer components leads to a
higher buffer capacity.
•Once the buffer components near depletion,
large changes in pH result.

55. 18-55
Determining the [H
3O
+
] and pH of a Buffer Solution
•The pH of a buffer system can be calculated if we know the
concentration of the weak acid [HA] and its conjugate base [A
-
]
HA
(aq)+ H
2O
(l)⇌ H
3O
+
(aq) + A
-
(aq)
55

56. 18-56
Henderson-Hasselbalch Equation
Henderson-Hasselbalchequation can be used to calculate
the pH of a buffer.
56pH=pK
a
+log
[A
-
]
[HA]

57. 18-57
Compare two buffer solutions
•Buffer 1:
–1 L buffer with 1.0 M CH
3CO
2H and 1.0 M
NaCH
3CO
2
•Buffer 2:
•1 L buffer with 0.1 M CH
3CO
2H and 0.1 M NaCH
3CO
2

58. 18-58
Initially [CH
3CO
2H] = [NaCH
3CO
2]

59. 18-59
14.7: Acid-Base Titrations
59
•Recall from Chapter 4 that we can
analyze an acid or base by reacting
it with a known concentration of
the other.
•Two types of titrations
1) Strong acid-strong base
2) Weak acid-strong base
Equivalence Point:

60. 18-60
Strong Acid-
Strong Base
Titration
HCl + NaOH H
2O + NaCl
•The pH is initially
very low (only
strong acid present)
•There is a gradual
rise in pH as base is
added.

61. 18-61
HCl + NaOH H
2O + NaCl
•Eventually the pH rises rapidly.
•Mid-point of the vertical part of
the curve is the equivalence
point.
•The pH at the equivalence point
is 7.
•Excess NaOHcontrols the pH
after the equivalence point.

62. 18-62
Weak Acid-Strong Base Titration
CH
3CO
2H + NaOH H
2O + NaCH
3CO
2
•Many features of this titration curve are
similar to a Strong Acid-Strong Base
titration.
•But this titration curve is somewhat more
complicated.

63. 18-63
CH
3CO
2H + NaOH H
2O + NaCH
3CO
2
•The pHstarts off
low, but not as low
as the strong acid
titration.
•There is a gradual
rise in pH as base is
added.

64. 18-64
CH
3CO
2H + NaOH H
2O + NaCH
3CO
2
•Eventually the pH rises
rapidly.
•Mid-point of the vertical part
of the curve is the equivalence
point.
•The pH at the equivalence
point is 8.72
•Excess NaOHcontrols the pH
after the equivalence point.

65. 18-65
CH
3CO
2H + NaOH H
2O + NaCH
3CO
2
•Half-Equivalence Point:
•Determining pK
afrom a
titration curve.

66. 18-66
Acid-Base Indicators
66
•One or two drops of indicator are often added
to the sample being titrated.
•An acid-base indicator is useful in determining
the equivalence point in a titration.
–The pH at which the indicator changes color is called
the endpoint.
•You want to choose an indicator that has an endpoint
that coincides with the pH at the equivalence point.

67. 18-67
Acid-Base Indicators
•Indicators are weak acids
–The weak acid is a different color than its conjugate base.

68. 18-68

69. 18-69
Indicator Selection: Need an indicator that changes color at a
pH near the equivalence point.