18-1
Chapter 18
Acid-Base Equilibria
18-2
Acid-Base Equilibria
18.1 Acids and bases in water
18.2 Auto-ionization 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 relationships 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 donations and the Lewis acid-base definition
18-3 Figure 18.1
Etching with acids
The inside surfaces of these light bulbs are etched with HF.
Acids are used to wash away oxides of silicon and metals during the production of computer chips.
18-4
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18-5
For reaction between a strong acid and strong base:
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
H+(aq) + OH-(aq) H2O(l)
For acid dissociation:
HA(g or l) + H2O(l) A-(aq) + H3O+(aq)
In aqueous solution, the H+ ions bind covalently towater to form solvated hydronium ions.
18-6
Figure 18.2
The nature of the hydrated proton
18-7
The Classical (Arrhenius) Definition of Acids and Bases
An acid is a substance that has H in its formula and dissociates in water
to yield H3O+.
A base is a substance that has OH in its formula and dissociates in water
to yield OH-.
Arrhenius acids contain covalently bonded H atoms that ionize in water.
Neutralization occurs when the H+ ion from the acid and the OH- ionfrom the base combine to form water.
H+(aq) + OH- (aq) H2O(l) ∆Horxn = -55.9 kJ
18-8
Defining the acid dissociation constant
Weak acids dissociate very slightly into ions in water.
Strong acids dissociate completely into ions in water.
HA(g or l) + H2O(l) H3O+ (aq) + A- (aq)
HA(aq) + H2O(l) H3O+ (aq) + A- (aq)
Kc >> 1
Kc << 1
Kc = [H3O+][A-]
[H2O][HA]
Kc[H2O] = Ka =[H3O+][A-]
[HA]weaker acid, lower [H3O
+],
smaller Ka
stronger acid, higher [H3O+],
larger Ka
18-9
Figure 18.3
strong acid: HA(g or l) + H2O(l) H3O+(aq) + A-(aq)
The extent of dissociation of strong acids
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18-10
Figure 18.3
The extent of dissociation of weak acids
weak acid: HA(aq) + H2O(l) H3O+(aq) + A-(aq)
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18-11
Figure 18.4
Reaction of zinc with a strong and a weak acid
1 M HCl(aq) 1 M CH3COOH(aq)
Zn(s) + 2H3O+(aq) Zn+2(aq)
+ 2H2O(l) + H2(g)
18-12
Acid strength decreases down
the table (smallerKa values)
18-13
SAMPLE PROBLEM 18.1
SOLUTION:
Classifying acid and base strength from the chemical formula
PROBLEM: Classify each of the following compounds as a strong acid, weak acid, strong base or weak base.
(a) H2SeO4 (b) (CH3)2CHCOOH (c) KOH (d) (CH3)2CHNH2
PLAN: Pay attention to the text definitions of acids and bases. Look at O for acids and for the -COOH group; watch for amine groups and cations in bases.
(a) strong acid - H2SeO4 - the number of oxygen atoms
exceeds the number of ionizable protons by a factor of 2.(b) weak acid - (CH3)2CHCOOH is an organic acid having a -COOH group
(a carboxylic acid).
(c) strong base - KOH is a Group 1A hydroxide.
(d) weak base - (CH3)2CHNH2 has a lone pair of electrons on the
nitrogen and is an amine.
18-14
The auto-ionization of water and the pH scale
H2O(l) H2O(l)
H3O+(aq) OH-(aq)
+
+
18-15
Kc = [H3O+][OH-]
[H2O]2
Kc[H2O]2 = [H3O+][OH-]
Defining the ion-product constant of water
Kw =
A change in [H3O+] causes an inverse change in [OH-].
= 1.0 x 10-14 at 25 oC
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
In an acidic solution, [H3O+] > [OH-]
In a basic solution, [H3O+] < [OH-]
In a neutral solution, [H3O+] = [OH-]
18-16
Figure 18.5
The relationship between [H3O+] and [OH-] and the relative acidity of solutions
[H3O+] [OH-]divide into Kw
acidic solution
basic solution
[H3O+] > [OH-] [H3O+] = [OH-] [H3O+] < [OH-]
neutral solution
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18-17
SAMPLE PROBLEM 18.2 Calculating [H3O+] and [OH-] in an aqueous
solution
PROBLEM: A research chemist adds a measured amount of HCl gas to pure water at 25 oC and obtains a solution with [H3O
+] = 3.0 x 10-4 M.
Calculate [OH-]. Is the solution neutral, acidic or basic?
SOLUTION:
PLAN: Use the Kw at 25 oC and the [H3O+] to find the corresponding [OH-].
Kw = 1.0 x 10-14 = [H3O+] [OH-]
[OH-] = Kw/ [H3O+] = 1.0 x 10-14/3.0 x 10-4 = 3.3 x 10-11 M
[H3O+] > [OH-]; the solution is acidic.
18-18
Figure 18.6
The pH values of some familiar
aqueous solutions
pH = -log [H3O+]
pOH = -log [OH-]
pK = -log K
Related Expressions
18-19
A low pK corresponds to a high K.
18-20
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Relationshipsbetween [H3O+],
pH, [OH-] and pOH
Figure 18.7
18-21
SAMPLE PROBLEM 18.3 Calculating [H3O+], pH, [OH-], and pOH
PROBLEM: In a restoration project, a conservator prepares copper-plate etching solutions by diluting concentrated nitric acid, HNO3, to 2.0 M, 0.30 M,
and 0.0063 M HNO3. Calculate [H3O+], pH, [OH-], and pOH of the
three solutions at 25 oC.
SOLUTION:
PLAN: HNO3 is a strong acid so [H3O+] = [HNO3]. Use Kw to find the [OH-]
and then convert to pH and pOH.
For 2.0 M HNO3, [H3O+] = 2.0 M and -log [H3O
+] = -0.30 = pH
[OH-] = Kw/ [H3O+] = 1.0 x 10-14/2.0 = 5.0 x 10-15 M; pOH = 14.30
[OH-] = Kw/ [H3O+] = 1.0 x 10-14/0.30 = 3.3 x 10-14 M; pOH = 13.48
For 0.3 M HNO3, [H3O+] = 0.30 M and -log [H3O
+] = 0.52 = pH
[OH-] = Kw/ [H3O+] = 1.0 x 10-14/6.3 x 10-3 = 1.6 x 10-12 M; pOH = 11.80
For 0.0063 M HNO3, [H3O+] = 0.0063 M and -log [H3O
+] = 2.20 = pH
18-22
Figure 18.8
Methods for measuring the pH of an aqueous solution
pH meterpH (indicator) paper
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18-23
The Brønsted-Lowry Definition of Acids and Bases
An acid is a proton donor, that is, any species that donates an H+ ion.All Arrhenius acids are Brønsted-Lowry acids.
A base is a proton acceptor, that is, any species
that accepts an H+ ion.
In the Brønsted-Lowry definition, an acid-base reaction is a protontransfer process.
18-24
Figure 18.9
Proton transfer is the essential feature of a Brønsted-Lowry acid-base reaction
(acid, H+ donor) (base, H+ acceptor)
HCl H2O
+
Cl- H3O+
+
lone pair
binds H+
(base, H+ acceptor) (acid, H+ donor)
NH3 H2O
+
NH4+ OH-
+
lone pair
binds H+
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18-25
An acid reactant produces a base product and the two constitute an acid-base conjugate pair.
The Conjugate Acid-Base Pair
H2S + NH3 HS- + NH4+
H2S and HS- are a conjugate acid-base pair.
HS- is the conjugate base of the acid H2S.
NH3 and NH4+ are a conjugate acid-base pair.
NH4+ is the conjugate acid of the base NH3.
Every acid has a conjugate base, and every base has a conjugate acid.
18-26
The conjugate base of the pair has one fewer H and one more negative charge than the acid.
The conjugate acid of the pair has one more H and one less negative charge than the base.
A Bronsted-Lowry acid-base reaction occurs when an acid and a basereact to form their conjugate base and conjugate acid, respectively.
acid1 + base2 base1 + acid2
18-27
Table 18.4 Conjugate Pairs in Some Acid-Base Reactions
base1 acid2+acid1 base2+
conjugate pair
conjugate pair
reaction 4 H2PO4- OH-+
reaction 5 H2SO4 N2H5++
reaction 6 HPO42- SO3
2-+
reaction 1 HF H2O+ F- H3O++
reaction 3 NH4+ CO3
2-+
reaction 2 HCOOH CN-+ HCOO- HCN+
NH3 HCO3-+
HPO42- H2O+
HSO4- N2H6
2++
PO43- HSO3
-+
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18-28
SAMPLE PROBLEM 18.4 Identifying conjugate acid-base pairs
PROBLEM: The following reactions are important environmental processes. Identify the conjugate acid-base pairs.
(a) H2PO4-(aq) + CO3
2-(aq) HPO42-(aq) + HCO3
-(aq)
(b) H2O(l) + SO32-(aq) OH-(aq) + HSO3
-(aq)
SOLUTION:
PLAN: Identify proton donors (acids) and proton acceptors (bases).
(a) H2PO4-(aq) + CO3
2-(aq) HPO42-(aq) + HCO3
-(aq)proton donor
proton acceptor
proton acceptor
proton donor
conjugate pair1conjugate pair2
(b) H2O(l) + SO32-(aq) OH-(aq) + HSO3
-(aq)
conjugate pair2conjugate pair1
proton donor
proton acceptor
proton acceptor
proton donor
18-29
Relative Acid-Base Strength and Reaction Direction
General Rule: An acid-base reaction proceeds to the greater extent in the direction in which a stronger acid and stronger base form aweaker acid and a weaker base.
H2S + NH3 HS- + NH4+ Kc > 1
A competition for the proton between the two bases!
HF + H2O F- + H3O+ Kc < 1
a b b a
a b b a
18-30
Figure 18.10
Strengths of conjugate acid-base
pairs
An acid-base reaction proceeds to the right if the acid reacts with a
base that is lower on the list because this combination produces a weaker
conjugate base and a weaker conjugate acid.
A weaker acid has a strongerconjugate base.
18-31
SAMPLE PROBLEM 18.5 Predicting the net direction of an acid-base reaction
PROBLEM: Predict the net direction and whether Ka is greater or less than 1 for each of the following reactions (assume equal initial concentrations of all species):
(b) H2O(l) + HS-(aq) OH-(aq) + H2S(aq)
(a) H2PO4-(aq) + NH3(aq) HPO4
2-(aq) + NH4+(aq)
SOLUTION:
PLAN: Identify the conjugate acid-base pairs and then consult Figure 18.10 to determine the relative strength of each. The stronger the species, the more preponderant will be its conjugate.
(a) H2PO4-(aq) + NH3(aq) HPO4
2-(aq) + NH4+(aq)
stronger acid weaker acidstronger base weaker base
Net direction is to the right; Kc > 1.
(b) H2O(l) + HS-(aq) OH-(aq) + H2S(aq)stronger baseweaker base stronger acidweaker acid
Net direction is to the left; Kc < 1.
18-32
Weak Acid Equilibrium Problems
Follow the general procedures described in Chapter 17
Two Key Assumptions
[H3O+] from the auto-ionization of water is negligible.
[HA]eq = [HA]init - [HA]dissoc ≈ [HA]init
(because a weak acid has a small Ka)
[HA]dissoc is very small
18-33
SAMPLE PROBLEM 18.6Find the Ka of a weak acid from the pH of its
solution
PROBLEM: Phenylacetic acid (C6H5CH2COOH, denoted as HPAc) builds up in
the blood of people afflicted with phenylketonuria, an inherited genetic disorder that, if left untreated, causes mental retardation and death. A study of the acid shows that the pH of a 0.12 M solution of HPAc is 2.60. What is the Ka of phenylacetic acid?
PLAN: Write the dissociation equation. Use pH and solution concentration to find Ka.
Ka = [H3O+][PAc-]
[HPAc]
Assumptions: At a pH of 2.60, [H3O+]HPAc >> [H3O
+]water
[PAc-]eq ≈ [H3O+]eq, and since HPAc is weak, [HPAc]eq ≈
[HPAc]initial = [HPAc]initial - [HPAc]dissociation SOLUTION: HPAc(aq) + H2O(l) H3O
+(aq) + PAc-(aq)
18-34
SAMPLE PROBLEM 18.6 (continued)
concentration (M) HPAc(aq) + H2O(l) H3O+(aq) + PAc-(aq)
initial 0.12 - 1 x 10-7 0
change --x +x +x
equilibrium -0.12 - x xx +(<1 x 10-7)
[H3O+] = 10-pH = 2.5 x 10-3 M, which is >> 10-7 ([H3O
+] from water)
x ≈ 2.5 x 10-3 M ≈ [H3O+] ≈ [PAc-] [HPAc]eq = 0.12 - x ≈ 0.12 M
Ka =(2.5 x 10-3) (2.5 x 10-3)
0.12= 5.2 x 10-5
Testing the assumptions:
= 4 x 10-3 %
x 100[HPAc]dissn: 2.5 x 10-3 M0.12 M
[H3O+]water:
1 x 10-7 M
2.5 x 10-3 Mx100
= 2.1 %
(the 5% rule is not violated in either case)
18-35
SAMPLE PROBLEM 18.7 Determining concentrations from Ka and initial [HA]
PROBLEM: Propanoic acid (CH3CH2COOH, simplified as HPr) is an organic
acid whose salts are used to retard mold growth in foods. What is
the [H3O+] of a 0.10 M aqueous solution of HPr (Ka = 1.3 x 10-5)?
SOLUTION:
PLAN: Write the dissociation equation and Ka expression; make assumptions about concentration that are valid; substitute.
x = [HPr]diss = [H3O+]from HPr = [Pr-]
Assumptions: For: HPr(aq) + H2O(l) H3O+(aq) + Pr-(aq)
Ka = [H3O+][Pr-]
[HPr]
HPr(aq) + H2O(l) H3O+(aq) + Pr-(aq)concentration (M)
initial 0.10 - 0 0
change --x +x +x
equilibrium -0.10 - x xx
Since Ka is small, we assume that x << 0.10
18-36
SAMPLE PROBLEM 18.7 (continued)
(x)2
0.101.3 x 10-5 =
[H3O+][Pr-]
[HPr]=
€
x= (0.10)(1.3x10−5) = 1.1 x 10-3 M = [H3O+]
Checking assumptions:
[HPr]diss: 1.1 x 10-3 M / 0.10 M x 100 = 1.1%
Ka =
18-37
percent HA dissociation = [HA]dissociated
[HA]initial
x 100
As the initial concentration of a weak acid decreases, the percent dissociation of the acid increases!
In the prior problem: for 0.10 M HPr, 1.1% dissociation for 0.010 M HPr, 3.6% dissociation
Rationale: larger solution volume accommodates more ions(analogous to pressure effects on gas equilibria when ∆ngas is non-zero)
18-38
Polyprotic acids
Acids that contain more than one ionizable proton
H3PO4(aq) + H2O(l) H2PO4-(aq) + H3O
+(aq)
H2PO4-(aq) + H2O(l) HPO4
2-(aq) + H3O+(aq)
HPO42-(aq) + H2O(l) PO4
3-(aq) + H3O+(aq)
Ka1 =[H3O
+] [H2PO4-]
[H3PO4]
Ka2 =[H3O
+] [HPO42-]
[H2PO4-]
Ka3 =[H3O
+] [PO43-]
[HPO42-]
Ka1 > Ka2 > Ka3
= 7.2 x 10-3
= 6.3 x 10-8
= 4.2 x 10-13
phosphoric acid, H3PO4
18-39
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18-40
Since successive acid dissociation constants typically differby several orders of magnitude, pH calculations for aqueous
solutions of the free acids can be simplified
by neglecting H3O+ generated by subsequent dissociations.
Why are the Ka values successively smaller as more H+ ions
dissociate from a polyprotic acid like phosphoric acid?
18-41
SAMPLE PROBLEM 18.8 Calculating equilibrium concentrations for a polyprotic acid
PROBLEM: Ascorbic acid (H2C6H6O6; abbreviated H2Asc), known as vitamin
C, is a diprotic acid (Ka1 = 1.0 x 10-5 and Ka2 = 5 x 10-12) found in
citrus fruit. Calculate [H2Asc], [HAsc-], [Asc2-], and the pH of a
0.050 M aqueous solution of H2Asc.
SOLUTION:
PLAN: Write out expressions for both dissociations and make assumptions.
Ka1 >> Ka2 so the first dissociation produces virtually all of the H3O+.
Ka1 is small, thus [H2Asc]initial ≈ [H2Asc]eq
After finding concentrations of the various species for the first dissociation, use them as initial concentrations for the second dissociation.
H2Asc(aq) + H2O(l) HAsc-(aq) + H3O+(aq)
Ka1 = [HAsc-][H3O
+]
[H2Asc]= 1.0 x 10-5
HAsc-(aq) + H2O(l) Asc2-(aq) + H3O+(aq) Ka2 =
[Asc2-][H3O+]
[HAsc-]= 5 x 10-12
18-42
SAMPLE PROBLEM 18.8 (continued)
H2Asc(aq) + H2O(l) HAsc-(aq) + H3O+(aq)concentration (M)
initial 0.050 - 0 0
change - x - + x + x
equilibrium 0.050 - x - x x
Ka1 = [HAsc-][H3O+]/[H2Asc] = 1.0 x 10-5 = (x)(x)/0.050 M
pH = -log(7.1 x 10-4) = 3.15
HAsc-(aq) + H2O(l) Asc2-(aq) + H3O+(aq)
7.1 x 10-4 - 0 0
change - x - + x + x
equilibrium 7.1 x 10-4 - x - x x
initial
€
= (0.050)(1.0x10−5)x x = 7.1 x 10-4 M = [HAsc-]
concentration (M)
x = [(7.1 x 10-4)(5 x 10-12)]0.5 = 6 x 10-8 M = [H3O+]
18-43
We can ignore the hydronium ion generated by second ionization.
Also: 7.1 x 10-4/0.050 x 100 = 1.4% This value is < 5%, thus the assumption made in the analysis of the first
dissociation reaction is justified.
[Asc-2] = (Ka2 x HAsc-)/[H3O+] = 5 x 10-12 M
18-44
Weak Bases: Their Relationship to Weak Acids
B(aq) + H2O(l) BH+(aq) + OH-(aq)
[BH+][OH-]
[B][H2O]
Base-dissociation constant, Kb
[BH+][OH-][B]
pKb decreases with increasing Kb (i.e., increasing base strength)
Kc =
Kb =
18-45
Main Classes of Weak Bases: nitrogen-containing molecules (ammoniaand amines) and anions of weak acids.
Ammonia:
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Kb = 1.76 x 10-5 (25 oC)
Amines:
RNH2, R2NH and R3N: all have a lone pair of electrons thatcan bind a proton donated by an acid
18-46
+
Figure 18.11
Abstraction of a proton from water by methylamine
+
lone pair
binds H+
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CH3NH2
methylamineH2O
CH3NH3+
methylammonium ion
OH-
18-47
Base strengthdecreases
downthe table(smaller
Kb values)
18-48
SAMPLE PROBLEM 18.9 Determining pH from Kb and initial [B]
PROBLEM: Dimethylamine, (CH3)2NH, a key intermediate in detergent
manufacture, has a Kb = 5.9 x 10-4. What is the pH of a 1.5 M
aqueous solution of (CH3)2NH?
SOLUTION:
PLAN: Perform this calculation as done for acids. Keep in mind that you are working with Kb and a base.
(CH3)2NH(aq) + H2O(l) (CH3)2NH2+(aq) + OH-(aq)
Assumptions:
[(CH3)2NH2+] = [OH-] = x [(CH3)2NH]eq≈ [(CH3)2NH]initial
Kb >> Kw so [OH-]water is negligible
initial 1.50 0 0-
change - x - + x + x
equilibrium 1.50 - x - x x
(CH3)2NH(aq) + H2O(l) (CH3)2NH2+(aq) + OH-(aq)concentration
18-49
SAMPLE PROBLEM 18.9 (continued)
Kb = 5.9 x 10-4 = [(CH3)2NH2
+][OH-]
[(CH3)2NH]
5.9 x 10-4 = (x) (x)
1.5 Mx = 3.0 x 10-2 M = [OH-]
Check assumption: [3.0 x 10-2 M / 1.5 M] x 100 = 2% (error is < 5%; thus, assumption is justified)
[H3O+] = Kw/[OH-] = 1.0 x 10-14/3.0 x 10-2 = 3.3 x 10-13 M
pH = -log (3.3 x 10-13) = 12.48
18-50
Anions of Weak Acids as Weak Bases
A-(aq) + H2O(l) HA(aq) + OH-(aq)
F-(aq) + H2O(l) HF(aq) + OH-(aq)
Kb = ([HF][OH-]) / [F-]
Rationale for the basicity of A-(aq): e.g., 1 M NaF
Acidity is determined by [OH-] generated from the F- reaction and
[H3O+] generated from the auto-ionization of water; since
[OH-] >> [H3O+], the solution is basic.
Ka x Kb = Kw
For a conjugate acid-base pair:
18-51
SAMPLE PROBLEM 18.10 Determining the pH of a solution of A-
PROBLEM: Sodium acetate (CH3COONa, abbreviated NaAc) has
applications in photographic development and textile dyeing. What is the pH of a 0.25 M aqueous solution of NaAc? Ka of
acetic acid (HAc) is 1.8 x 10-5.
SOLUTION:
PLAN: Sodium salts are soluble in water so [Ac-] = 0.25 M.Use Ka to find Kb.
initial 0.25 - 0 0
change -x +x +x-
equilibrium -0.25 - x x x
Ac-(aq) + H2O(l) HAc(aq) + OH-(aq)concentration
Kb = [HAc][OH-]
[Ac-]=
Kw
Ka
= 5.6 x 10-10Kb = 1.0 x 10-14
1.8 x 10-5
18-52
SAMPLE PROBLEM 18.10 (continued)
Kb = [HAc][OH-]
[Ac-]
[Ac-] = 0.25 M - x ≈ 0.25 M (since Kb is small)
5.6 x 10-10 ≈ x2/0.25 M
x ≈ 1.2 x 10-5 M = [OH-]
Check assumption: [1.2 x 10-5 M / 0.25 M] x 100 = 4.8 x 10-3 %
[H3O+] = Kw/[OH-] = 1.0 x 10-14/1.2 x 10-5 = 8.3 x 10-10 M
pH = -log (8.3 x 10-10 M) = 9.08
18-53
Molecular Properties and Acid Strength
Non-metal hydrides: two factors determine acid strength, namely,the electronegativity of the central non-metal atom E and the strength
of the E-H bond
Non-metal hydride acid strength increases across a period (E electronegativity effect)
Non-metal hydride acid strength increases down a group(E-H bond strength effect)
18-54
Figure 18.12
The effect of atomic and molecular properties on non-metal hydride acidity
6A(16)
H2O
H2S
H2Se
H2Te
7A(17)
HF
HCl
HBr
HI
electronegativity increases, acidity
increases
bon
d s
tre
ngth
de
cre
ase
s,
aci
dity
incr
ea
ses
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18-55
H O I H O Br H O Cl> >
H O Cl
O
O
O<<
Figure 18.13
The relative strengths of oxoacids
H O Cl
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Electronegativity increases, acidity increases
number of O atoms increases,acidity increases
18-56
Acidity of Hydrated Metal Ions
Aqueous solutions of certain metal ions are acidic because the hydrated
metal ion transfers an H+ ion to water.
Generalized Reactions
M(NO3)n(s) + xH2O(l) M(H2O)xn+(aq) + nNO3
-(aq)
M(H2O)xn+(aq) + H2O(l) M(H2O)x-1OH(n-1)+(aq) + H3O
+(aq)
18-57
Al(H2O)5OH2+Al(H2O)63+
Figure 18.14
The acidic behavior of the hydrated Al3+ ion
H2O H3O+
electron density
drawn toward Al3+solvent H2O acts
as base
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18-58
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are needed to see this picture.
high chargeand small size
enhance acidity
18-59
Acid-Base Properties of Salt Solutions
A. Salts That Yield Neutral Solutions: the anion of a strong acid and the cation of a strong base (the ions do not react with water)
The anion of a strong acid is a much weaker base than water (HNO3).
The cation of a strong base only becomes hydrated (NaOH).
HNO3(l) + H2O(l) NO3-(aq) + H3O
+(aq)
NaOH(s) Na+(aq) + OH-(aq)
18-60
B. Salts That Yield Acidic Solutions: (1) the anion of a strong acidand the cation of a weak base (the cation acts as a weak acid); (2) small, highly charged metal ions; (3) cations of strong bases and anions of polyprotic acids with another ionizable proton.
NH4Cl(s) NH4+ (aq) + Cl-(aq)
NH4+(aq) + H2O(l) NH3(aq) + H3O
+ (aq)Case 1
Fe(NO3)3(s) + 6H2O(l) Fe(H2O)63+(aq) + 3NO3
-(aq)
Fe(H2O)63+(aq) + H2O(l) Fe(H2O)5OH2+(aq) + H3O
+(aq)
Case 2
NaH2PO4(s) Na+(aq) + H2PO4- (aq)
H2PO4-(aq) + H2O(l) HPO4
2-(aq) + H3O+(aq)
Case 3
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C. Salts That Yield Basic Solutions: the anion of a weak acid and thecation of a strong base (the anion acts as a weak base)
CH3COONa(s) Na+(aq) + CH3COO-(aq)
CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH-(aq)
18-62
18-63
SAMPLE PROBLEM 18.11 Predicting the relative acidity of salt solutions
PROBLEM: Predict whether aqueous solutions of the following compounds are acidic, basic, or neutral (write an equation for the reaction of the appropriate ion with water to explain pH effect).
(a) potassium perchlorate, KClO4 (b) sodium benzoate, C6H5COONa
(c) chromium trichloride, CrCl3 (d) sodium hydrogen sulfate, NaHSO4
SOLUTION:
PLAN: Consider the acid-base nature of the anions and cations. Strong acid-strong base combinations produce a neutral solution; strong acid-weak base, acidic; weak acid-strong base, basic.
(a) The ions are K+ and ClO4-, which come from a strong base
(KOH) and a strong acid (HClO4). The salt solution will be neutral.(b) Na+ comes from the strong base NaOH while C6H5COO- is the anion of a
weak organic acid. The salt solution will be basic.(c) Cr3+ is a small cation with a large + charge, so its hydrated form will react
with water to produce H3O+. Cl- comes from the strong acid HCl. The
salt solution will be acidic.(d) Na+ comes from a strong base. HSO4
- can react with water to form H3O+.
The salt solution will be acidic.
18-64
Salts of Weakly Acidic Cations and Weakly Basic Anions
Both components react with water!
Acidity is determined by the relative acid strength and base strength of the separated ions.
example: NH4HS
NH4+(aq) + H2O(l) NH3(aq) + H3O
+(aq)
HS-(aq) + H2O(l) H2S(aq) + OH-(aq)
Ka(NH4+) = 5.7 x 10-10 Kb(HS-) = 1 x 10-7
Since Kb > Ka, the solution is basic.
18-65
SAMPLE PROBLEM 18.12 Predicting the relative acidity of salt solutions from Ka and Kb of the ions
PROBLEM: Determine whether an aqueous solution of zinc formate, Zn(HCOO)2, is acidic, basic, or neutral.
SOLUTION:
PLAN: Both Zn2+ and HCOO- come from weak conjugates. In order to find the relatively acidity, write the dissociation reactions and use the information in Tables 18.2 and 18.7.
Ka Zn(H2O)62+ = 1 x 10-9
Ka HCOOH = 1.8 x 10-4 ; Kb = Kw/Ka = 1.0 x 10-14/1.8 x 10-4 = 5.6 x 10-11
Ka for Zn(H2O)62+ > Kb HCOO-; the solution is acidic.
Zn(H2O)62+(aq) + H2O(l) Zn(H2O)5OH+(aq) + H3O
+(aq)
HCOO-(aq) + H2O(l) HCOOH(aq) + OH-(aq)
18-66
The Leveling Effect
In water, the strongest acid possible is H3O+ and the
strongest base possible is OH-.
Any acid stronger than H3O+ donates its proton to H2O, and
any base stronger than OH- accepts a proton from H2O; thus,water exerts a leveling effect (levels the strengths of all strong acids and bases).
To rank strong acids: must dissolve in a solvent that is a weakerbase than water (i.e., one that accepts their protons less readily).
HCl(g) + CH3COOH(l) Cl-(acet) + CH3COOH2+(acet)
HBr(g) + CH3COOH(l) Br-(acet) + CH3COOH2+(acet)
HI(g) + CH3COOH(l) I-(acet) + CH3COOH2+(acet)
KHI > KHBr > KHCl
18-67
Three Definitions of Acids and Bases
The Arrhenius Definition
The Brønsted-Lowry Definition
The Lewis Definition
18-68
F
B
F F
H
N
H H
+
F
B
F F
H
N
H H
acid base adduct
An acid is an electron-pair acceptor. A base is an electron-pair donor.
M2+
H2O(l)
M(H2O)42+(aq)
adduct
The Lewis Acid-Base Definition
The adductcontains a
new covalentbond.
18-69
Lewis Acids with Electron-Deficient Atoms
Lewis Acids with Polar Multiple Bonds
Metal Cations as Lewis Acids
ROR’ + AlCl3 R-O-R’
AlCl3base acid adduct
SO2 + H2O H2SO3
M2+ + 4H2O M(H2O)42+
baseacid adduct
adductbaseacid
Metal ions actas Lewis acidswhen dissolved
in water.
Lewis acidscontain (or can
generate) a vacant orbital.
18-70
Figure 18.15
Mg2+ ion as a Lewis acid in the chlorophyll molecule
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
The ion accepts electron pairs fromfour nitrogen atoms comprising part
of the chlorophyll molecule
18-71
SAMPLE PROBLEM 18.13 Identifying Lewis acids and bases
PROBLEM: Identify the Lewis acids and Lewis bases in the following reactions:
(a) H+ + OH- H2O
(b) Cl- + BCl3 BCl4-
(c) K+ + 6H2O K(H2O)6+
SOLUTION:
PLAN: Look for electron pair acceptors (acids) and donors (bases).
(a) H+ + OH- H2O
acceptor
donor
(b) Cl- + BCl3 BCl4-
donor
acceptor
(c) K+ + 6H2O K(H2O)6+
acceptor
donor