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Chapter 16: ACID-BASE EQUILIBRIA

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Chapter 16: ACID-BASE EQUILIBRIA 16.1 Acids and Bases: A Brief Review 16.2 Br nsted-Lowry Acids and Bases 16.3 The Autoionization of Water 16.4 The pH Scale – PowerPoint PPT presentation

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Title: Chapter 16: ACID-BASE EQUILIBRIA


1
Chapter 16 ACID-BASE EQUILIBRIA
16.1 Acids and Bases A Brief Review 16.2
Brønsted-Lowry Acids and Bases 16.3 The
Autoionization of Water 16.4 The pH Scale 16.5
Strong Acids and Bases 16.6 Weak Acids 16.7 Weak
Bases 16.8 Relationship between Ka and Kb 16.9
Acid-Base Properties of Salt Solutions 16.10
Acid-Base Behavior and Chemical Structure 16.11
Lewis Acids and Bases
2
16.1 Acids and Bases A Brief Review
  • ACIDS BASES
  • Taste sour Taste bitter
  • Turn litmus red Feels slippery
  • H Turn litmus blue
  • OH-
  • Arrhenius
  • when dissolved in when dissolved in
  • water, H water, OH-

3
16.2 Brønsted-Lowry Acids and Bases
  • BrønstedLowry
  • Acid Proton donor
  • Base Proton acceptor
  • H is a proton with no valence electron.
  • H H3O

4
  • A BrønstedLowry acid
  • must have a removable (acidic) proton.
  • A BrønstedLowry base
  • must have a pair of nonbonding electrons.
  • HCl (g) H2O (l) ? Cl- (aq) H3O (aq)
  • HCl (g) NH3 (g) ? Cl- (g) NH4 (g)
  • NH3 (aq) H2O (l) ? NH4(aq) OH- (aq)

5
If it can be either
  • ...it is amphiprotic.
  • HCO3-
  • HSO4-
  • H2O

6
What Happens When an Acid Dissolves in Water?
  • Water acts as a BrønstedLowry base and abstracts
    a proton (H) from the acid.
  • As a result, the conjugate base of the acid and a
    hydronium ion are formed.

7
Conjugate Acids and Bases
  • From the Latin word conjugare, meaning to join
    together.
  • Reactions between acids and bases always yield
    their conjugate bases and acids.

8
Sample Exercise 16.1 Identifying Conjugate Acids
and Bases
  • What is the conjugate base of each of the
    following acids HClO4, H2S, PH4, HCO3?
  • What is the conjugate acid of each of the
    following bases CN, SO42, H2O, HCO3 ?
  • c) Write the formula for the conjugate acid of
    each of the following HSO3, F , PO43, CO.

9
Sample Exercise 16.2 Writing Equations for
Proton-Transfer Reactions
  • The hydrogen sulfite ion (HSO3) is amphiprotic.
  • Write an equation for the reaction of HSO3 with
    water, in which the ion acts as an acid and the
    reaction in which the ion acts as a base. In both
    cases identify the conjugate acidbase pairs.
  • When lithium oxide (Li2O) is dissolved in water,
    the solution turns basic from the reaction of the
    oxide ion (O2) with water. Write the reaction
    that occurs, and identify the conjugate acidbase
    pairs.

10
Acid and Base Strength
  • Strong acids are completely dissociated in water.
  • Their conjugate bases are quite weak with no
    tendency to be protonated.
  • Weak acids only dissociate partially in water.
  • Their conjugate bases are weak bases.
  • Substances with negligible acidity do not
    dissociate in water.
  • Their conjugate bases are exceedingly strong.

11
Acid and Base Strength
  • In any acid-base reaction, the equilibrium will
    favor the reaction that moves the proton to the
    stronger base.

HCl(aq) H2O(l) ??? H3O(aq) Cl-(aq)
H2O is a much stronger base than Cl-, so the
equilibrium lies very far to the right (Kgtgt1).
Acetate is a stronger base than H2O, so the
equilibrium favors the left side (Klt1).
12
Sample Exercise 16.3 Predicting the Position of a
Proton-Transfer Equilibrium
13
16.3 The Autoionization of Water
  • As we have seen, water is amphoteric.
  • In pure water, a few molecules act as bases and a
    few act as acids.
  • This is referred to as autoionization.

14
Ion-Product Constant
  • The equilibrium expression for this process is
  • Kc H3O OH-
  • This special equilibrium constant is referred to
    as the ion-product constant for water, Kw.
  • At 25C, Kw 1.0 ? 10-14

15
Sample Exercise 16.4 Calculating H for Pure
Water
  • Calculate the values of H and OH- in a
    neutral solution at 25 ºC.
  • Indicate whether solutions with each of the
    following ion concentrations are neutral, acidic,
    or basic
  • H 4 109 M
  • H 4 109 M
  • OH 7 1013 M .

16
Sample Exercise 16.5 Calculating H from OH-
  • Calculate the concentration of H(aq) in
  • a solution in which OH is 0.010 M
  • a solution in which OH is 1.8 109 M
  • Calculate the concentration of OH(aq) in a
    solution in which
  • H 2 106 M
  • H OH
  • H 100 OH

17
16.4 The pH Scale
  • pH is defined as the negative base-10 logarithm
    of the hydronium ion concentration.
  • pH -log H3O
  • In pure water, Kw H3O OH- 1.0 ? 10-14
  • Because in pure water H3O OH-,
  • H3O (1.0 ? 10-14)1/2 1.0 ? 10-7
  • Therefore, in pure water,
  • pH -log (1.0 ? 10-7) 7.00

18
pH
  • An acid has a higher H3O than pure water, so
    its pH is lt7
  • pH decreases as acidity H increases
  • A base has a lower H3O than pure water, so its
    pH is gt7.
  • pH decreases as basicity OH- increases

19
pH
  • These are the pH values for several common
    substances.

20
Sample Exercise 16.6 Calculating pH from H
1. Calculate the pH values for the following
solutions H 1.0 1012 M H 5.6 106
M 2. In a sample of lemon juice H is 3.8
104 M. What is the pH? 3. A commonly
available window-cleaning solution has OH
1.9 106 M . What is the pH?
21
Sample Exercise 16.7 Calculating H from pH
A sample of freshly pressed apple juice has a pH
of 3.76. Calculate H. A solution formed by
dissolving an antacid tablet has a pH of 9.18.
Calculate H .

22
Other p Scales
  • The p in pH tells us to take the negative log
    of the quantity (in this case, hydrogen ions).
  • Some similar examples are
  • pOH -log OH-
  • pKw -log Kw
  • Because H3O OH- Kw 1.0 ? 10-14,
  • we know that
  • -log H3O -log OH- -log Kw 14.00
  • or, in other words, pH pOH pKw 14.00

23
How Do We Measure pH?
  • For less accurate measurements, one can use
  • Litmus paper
  • Red paper turns blue above pH 8
  • Blue paper turns red below pH 5
  • An indicator

24
How Do We Measure pH?
  • For more accurate measurements, one uses a pH
    meter, which measures the voltage in the solution.

25
16.5 Strong Acids and Bases
  • Strong Acids
  • You will recall that the seven strong acids are
    HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.
  • These are, by definition, strong electrolytes and
    exist totally as ions in aqueous solution.
  • For the monoprotic strong acids,
  • H3O acid.

26
Sample Exercise 16.8 Calculating the pH of a
Strong Acid
What is the pH of a 0.040 M solution of
HClO4? An aqueous solution of HNO3 has a pH of
2.34. What is the concentration of the acid?
27
  • Strong Bases
  • Strong bases are the soluble hydroxides, which
    are the alkali metal and heavier alkaline earth
    metal hydroxides (Ca2, Sr2, and Ba2).
  • Again, these substances dissociate completely in
    aqueous solution.

28
Sample Exercise 16.9 Calculating the pH of a
Strong Base
  • What is the pH of
  • a 0.028 M solution of NaOH
  • a 0.0011 M solution of Ca(OH)2
  • What is the concentration of a solution of
  • KOH for which the pH is 11.89
  • Ca(OH)2 for which the pH is 11.68

29
16.6 Weak Acids
  • Weak acids (weak ? dilute) are only partially
    ionized
  • For a generalized acid dissociation,
  • the equilibrium expression would be
  • This equilibrium constant is called the
    acid-dissociation constant, Ka.

30
Dissociation Constants
  • The greater the value of Ka, the stronger the
    acid.

31
Calculating Ka from the pH
  • The pH of a 0.10 M solution of formic acid,
    HCOOH, at 25C is 2.38. Calculate Ka for formic
    acid at this temperature.
  • HCOOH (aq) ? H (aq) COOH- (aq)
  • We know that
  • pH -log H3O 2.38
  • log H3O - 2.38
  • H3O 10 -2.38 0.0042M

32
Calculating Ka from the pH
  • HCOOH ? H
    COOH-

Initial 0.1 M 0 0
Change 0.0042M
Equilibrium
1.8 ? 10-4
33
Calculating Percent Ionization
  • Percent Ionization ? 100
  • In this example
  • H3Oeq 4.2 ? 10-3 M
  • HCOOHinitial 0.10 M
  • x 100

4.2
34
Sample Exercise 16.10
  • Niacin, one of the B vitamins, has the following
    molecular structure
  • A 0.020 M solution of niacin has a pH of 3.26.
  • What percentage of the acid is ionized in this
    solution?
  • What is the acid-dissociation constant, Ka, for
    niacin?

35
Calculating pH from Ka
  • Knowing Ka and the initial concentration of the
    weak acid, we can calculate H (and pH).
  • Calculate the pH of a 0.30 M solution of acetic
    acid, HC2H3O2, at 25C.
  • HC2H3O2(aq) H2O(l) H3O(aq)
    C2H3O2-(aq)
  • Ka for acetic acid at 25C is 1.8 ? 10-5.

36
Calculating pH from Ka
HC2H3O2(aq) H2O(l) H3O(aq)
C2H3O2-(aq)
  • The equilibrium constant expression is
  • 1.8 ? 10-5

37
Calculating pH from Ka
We next set up a table
C2H3O2, M H3O, M C2H3O2-, M
Initially 0.30 0 0
Change -x x x
At Equilibrium 0.30 - x ? 0.30 x x
We are assuming that x will be very small
compared to 0.30 and can, therefore, be ignored.
38
Calculating pH from Ka
  • Now,

(1.8 ? 10-5) (0.30) x2 5.4 ? 10-6 x2 2.3 ?
10-3 x
pH -log H3O pH -log (2.3 ? 10-3) pH 2.64
39
Sample Exercise 16.12 Using Ka to Calculate pH
Calculate the pH of a 0.20 M solution of HCN.
(According to Appendix D, Ka 4.9
1010) The Ka for niacin is 1.5 10-5.
What is the pH of a 0.010 M solution of niacin?
40
Sample Exercise 16.13 Using Ka to Calculate
Percent Ionization
  • Calculate the percentage of HF molecules ionized
    in the following. (Ka 6.8 10-4)
  • 0.10 M HF solution
  • 0.010 M HF solution.
  • In Practice Exercise 16.11, we found that the
    percent ionization of niacin (Ka 1.5 10-5) in
    a 0.020 M solution is 2.7. Calculate the
    percentage of niacin molecules ionized in a
    solution that is
  • 0.010 M, (b) 1.0 10-3 M.

41
Polyprotic Acids
  • Have more than one acidic proton.
  • If the difference between the Ka for the first
    dissociation and subsequent Ka values is 103 or
    more, the pH generally depends only on the first
    dissociation.

42
16.7 Weak Bases
  • Bases react with water to produce hydroxide ion.

43
Base-dissociation constant
  • The equilibrium constant expression for this
    reaction is

where Kb is the base-dissociation constant.
44
  • Kb can be used to find OH- and, through it, pH.

45
pH of Basic Solutions
  • What is the pH of a 0.15 M solution of NH3?

NH3, M NH4, M OH-, M
Initially 0.15 0 0
At Equilibrium 0.15 - x ? 0.15 x x
46
pH of Basic Solutions
  • (1.8 ? 10-5) (0.15) x2
  • 2.7 ? 10-6 x2
  • 1.6 ? 10-3 x2

Therefore, OH- 1.6 ? 10-3 M pOH -log (1.6
? 10-3) pOH 2.80 pH 14.00 - 2.80 pH 11.20
47
16.8 Relationship between Ka and Kb
  • Ka and Kb are related in this way
  • Ka ? Kb Kw
  • Therefore, if you know one of them, you can
    calculate the other.

48
Sample Exercise 16.17 Calculating Ka or Kb for a
Conjugate Acid-Base Pair
  • Calculate
  • the base-dissociation constant, Kb, for the
    fluoride ion (F)
  • the aciddissociation constant, Ka, for the
    ammonium ion (NH4)

49
16.9 Acid-Base Properties of Salt Solutions
  • Assume all salts are completely ionized.
  • Anions of weak acids can react with water in a
    hydrolysis reaction to form OH- and the conjugate
    acid
  • These anions are bases
  • Anions from strong acids (NO3) are not basic
  • Behavior of amphoteric anions (HSO3-) depend on
    value of Ka and Kb

50
Reactions of Cations with Water
  • Cations with acidic protons (like NH4) will
    lower the pH of a solution.
  • Most metal cations (except alkali metals and
    heavier alkaline earth metals) that are hydrated
    in solution also lower the pH of the solution.

51
pH of a salt solution
  • Salts derived from strong base strong acid
  • neutral (pH 7) neither anion nor cation
    hydrolyzes. NaCl (from NaOH HCl)
  • Salts derived from strong base weak acid
  • anion is a strong conjugate base, phgt7
  • NaClO (from NaOH HClO)
  • Salts derived from weak base strong acid
  • cation is a strong conjugate acid, pHlt7
  • NH4Cl (from NH3 HCl)
  • Salts derived from weak base weak acid
  • both anion cation hydrolyzes. pH depends on
    extent NH4C2H3O2 (from NH3 HC2H3O2)

52
Sample Exercise 16.18 Determining Whether Salt
Solutions Are Acidic, Basic, or Neutral
  • Predict whether the salt Na2HPO4 will form an
    acidic solution or a basic solution on dissolving
    in water.
  • Determine whether aqueous solutions of each of
    the following salts will be acidic, basic, or
    neutral
  • Ba(CH3COO)2
  • NH4Cl
  • CH3NH3Br
  • KNO3
  • Al(ClO4)3

53
16.10 Acid-Base Behavior and Chemical Structure
  • Factors Affecting Acid Strength
  • Polarity
  • H X Proton donated
  • H Na Proton accepted (H has neg charge)
  • H C nonpolar no affect on pH
  • Strength of bond
  • HF is a weak acid
  • Stability of conjugate base

54
Oxyacids
In oxyacids, in which an OH is bonded to a
central atom, Y, Y O-H if Y metal, ionic
compound formed ? OH- Y O-H if Y nonmetal,
covalent bond formed ? the more electronegative Y
is, the more acidic the acid.
55
Oxyacids
  • For a series of oxyacids, acidity increases with
    the number of oxygens.

56
Factors Affecting Acid Strength
  • Resonance in the conjugate bases of carboxylic
    acids stabilizes the base and makes the conjugate
    acid more acidic.

57
Sample Exercise 16.20 Predicting Relative
Acidities from Composition and Structure
Arrange the compounds in each of the following
series in order of increasing acid strength (a)
AsH3, HI, NaH, H2O (b) H2SO4, H2SeO3,
H2SeO4. In each of the following pairs choose
the compound that leads to the more acidic (or
less basic) solution (a) HBr, HF (b) PH3, H2S
(c) HNO2, HNO3 (d) H2SO3, H2SeO3.
58
16.11 Lewis Acids and Bases
  • Lewis acids are defined as electron-pair
    acceptors.
  • Atoms with an empty valence orbital can be Lewis
    acids.

59
Lewis Bases
  • Lewis bases are defined as electron-pair donors.
  • Anything that could be a BrønstedLowry base is a
    Lewis base.
  • Lewis bases can interact with things other than
    protons, however.
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