Chapter Fifteen

1
Chapter Fifteen

• Acids, Bases, And
• Acid-Base Equilibria

2
The Arrhenius Theory (section 2.8)

• acid a substance that produces H ion when
  dissolved in water
• base - a substance that produces OH- ion when
  dissolved in water
• strong acid ionizes essentially completely into
  H and an anion
• strong base dissociates nearly completely into
  OH- and a cation
• weak acid and base ionize reversibly
• Limitation of Arrhenius Theory

3
The Brønsted-Lowry Theory

• Acid proton donor
• Base proton acceptor
• Conjugate acid and base, HA/A-, differ by one
  proton.
• The conjugate acid of a base is the base plus
  the attached proton and the conjugate base of an
  acid
• is the acid minus the proton.
• A substance that can act either as an acid or a
  base
• is amphiprotic.
• For weak acids and bases, equations can be
  written
• to describe equilibrium conditions.

4
Ionization of Ammonia
5
Strengths Of Conjugate Acid-Base Pairs

• The stronger an acid, the weaker is its conjugate
  base.
• The stronger a base, the weaker is its conjugate
  acid.
• An acid-base reaction is favored in the direction
  from the stronger member to the weaker member of
  each conjugate acid-base pair.
• Ka and Kb values are used to compare the
  strengths of weak acids and bases.
• Water has a leveling effect when the strong
  acids are dissolved in water, they all completely
  ionize to the hydronium ion.

6
Relative Acid-Base Pair Strength
7
Strong acids Weak conjugate bases
Leveling effect between the dotted lines also
weak acids and bases
Strong bases Weak conjugate acids
8
Acid And Base Ionization Constants

• weak acid CH3COOH H2O º H3O CH3COO-
• H3OCH3COO-
• Acid ionization constant Ka
• CH3COOH
• weak base NH3 H2O º NH4 OH-
• NH4OH-
• Base ionization constant Kb
• NH3
• Acid and base ionization constants are the
  measure of the strengths of acids and bases.

9
Relative Strengths Of Binary Acids

• H X
• The greater the tendency for the transfer of a
  proton from HX to H2O, the more the forward
  reaction is favored and the stronger the acid.
• in a periodic group
• Bond-dissociation energy is inversely
  proportional to acid strength. The weaker the
  bond, the stronger the acid.
• Anion radius is directly proportional to acid
  strength. The larger the resultant anions
  radius, the stronger is the acid.
• The strengths of binary acids increase from top
  to bottom in a group of the periodic table.

10
Relative Strengths Of Binary Acids

• H X
• in a periodic group
• Bond dissociation energy the weaker the bond,
  the stronger the acid.
• Bond dissociation energy 569 gt 431 gt
  368 gt 297
• (kJ/mol) HF HCl HBr HI
• Acid strength Ka 6.6x10-4 lt 106 lt 108 lt
  109
• Anion radius the larger the anions radius, the
  stronger the acid.
• Anion radius (ppm) 136 lt 181 lt 195 lt
  216
• (kJ/mol) HF HCl HBr HI
• Acid strength Ka 6.6x10-4 lt 106 lt 108 lt
  109
• The strength of binary acids increase from top to
  bottom in a
• group of the periodic table.

11
Relative Strengths Of Binary Acids

• H X
• in a period
• The larger the electronegativity difference
  between H and X, the more easily the proton is
  removed and the stronger is the acid.
• ? EN 0.4 lt 0.9 lt 1.4 lt 1.9
• Acid strength CH4 NH3 H2O HF
• The strengths of binary acids increase from left
  to right across a period of the periodic table.

12
Representative Trends In Strengths of Binary
Acids
13
Strengths Of Oxoacids

• H O - E
• Two factors
• - electronegativity of the central atom (E)
• - number of terminal oxygen atoms
• As the electronegativity of the central atom (E)
  increases and as the number of terminal oxygen
  atoms increases, the acid strength also increases.

14
Strengths Of Oxoacids

• As the electronegativity of the central atom (E)
• increases the acid strength increases.
• Electronegativity 2.5 lt 2.8 lt 3.0
• HOI HOBr HOCl
• Acid strength Ka 2.3x10-11 lt 2.5x10-9
  lt 2.9x10-8
• As the number of terminal oxygen atoms increases,
  
• the acid strength also increases.
• of terminal 0 1 2 3
• O atoms O O
• H-O-Cl H-O-Cl-O H-O-Cl-O H-O-Cl-O
• O
• Acid strength 2.9x10-8 lt 1.1x10-2 lt 1000 lt
  108

15
Strengths Of Carboxylic Acids

• O
• R C O - H
• Carboxylic acids all have the -COOH group in
  common therefore, differences in acid strength
  must come from differences in the R group
  attached to the carboxyl group.
• In general, the more that electronegative atoms
  are attached in the R group, the stronger the
  acid.

16
Strengths Of Carboxylic Acids

• In general, the more that electronegative atoms
  are
• attached in the R group, the stronger the acid.
• I-CH2CH2COOH Cl-CH2CH2COOH CH3-CHClCOOH
  CH3CCl2COOH
• Ka 8.3x10-5 lt 1.0x10-4 lt
  1.4x10-3 lt 8.7x10-3

17
Strengths Of Amines As Bases

• Aromatic amines are much weaker bases than
  aliphatic amines.
• This is due in part to the fact that the p
  electrons in the benzene ring of an aromatic
  molecule are delocalized and can involve the Ns
  lone-pair electrons in the resonance hybrid.
• As a result, the lone-pair electrons are much
  less likely to accept a proton.
• Electron-withdrawing groups on the ring further
  diminish the basicity of aromatic amines relative
  to aniline.

18
Strengths Of Amines As Bases

• BrNH2 NH3 C6H5NH2
• Kb 2.5x10-8 1.8x10-5 7.4x10-10

19
Amine bases
ammonia
R CH3, CH2CH3 aliphatic amine bases
N
H
aromatic amine base
H
20
Self-Ionization Of Water

• Even the purest of water conducts electricity.
  This is due to the fact that water self-ionizes,
  that is, it creates a small amount of H3O and
  OH-.
• H2O H2O º H3O OH-
• Kw H3OOH-
• Kw - ion product of water
• Kw 1.0 x 10-14 at 25 oC
• This equilibrium constant is very important
  because it applies to all aqueous solutions -
  acids, bases, salts, and non-electrolytes - not
  just to pure water.

21
Self ionization reaction of water

-
O
O

O
H

H
H
H
H
22
pH and pOH

• pH - logH3O H3O 10-pH
• pH - logOH- OH- 10-pOH
• pKw pH pOH 14.00
• neutral solution H3O OH- 10 7 M pH
  7.0
• acidic solution H3O gt 10-7 M pH lt
  7.0
• basic solution H3O lt 10-7 M pH gt 7.0

23
The pH Scale
24
An Example

• The pH of milk of magnesia, a suspension of
  solid magnesium hydroxide in its saturated
  aqueous solution, is measured to be 10.52. What
  is the molarity of Mg(OH)2 in its saturated
  aqueous solution?

25
Equilibrium In Solutions Of Weak Acids And Weak
Bases

• weak acid HA H2O º H3O A-
• H3OA-
• Ka
• HA
• weak base B H2O º HB OH-
• HBOH-
• Kb
• B
• You need to be able to write acid and base
  ionization equations!!!

26
pKa and pKb

pKa -logKa pKb -logKb larger Ka gt smaller
pKa gt stronger acid larger Kb gt smaller pKb gt
stronger base
27
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28
Some Acid-Base Equilibrium Calculations

• These calculations are similar to the equilibrium
  calculations performed in Chapter 14.
• An equation is written for the reversible
  reaction, data are organized under this equation,
  the changes that occur in establishing
  equilibrium are assessed, and finally
  calculations of equilibrium concentrations are
  done.
• When Macid/Ka gt 100 or Mbase/Kb gt 100,
• the calculations can be simplified.

29
An Example

• 1.Determine the concentrations of H3O, CH3COOH
  and CH3COO-, and the pH of 1.00 M CH3COOH
  solution. Ka 1.8 x 10-5.
• 2. What is the pH of a solution that is 0.200 M
  in methylamine, CH3NH2? Kb 4.2 x 10-4.

30
Are Salts Neutral, Acidic or Basic?

• Salts are ionic compounds formed in the reaction
• between an acid and a base.
• 1. NaCl
• Na is from NaOH , a strong base
• Cl- is from HCl, a strong acid
• H2O
• NaCl (s) gt Na (aq) Cl- (aq)
• Na and Cl- ions do not react with water.
• The solution is neutral.

31
Are Salts Neutral, Acidic or Basic?

• 2. KCN
• K is from KOH , a strong base
• CN- is from HCN, a weak acid
• H2O
• KCN (s) gt K (aq) CN- (aq)
• K ions do not react with water, but CN- ions
  do.
• CN- H2O º HCN OH- hydrolysis
• The OH- ions are produced, so the solution is
  basic.

32
Are Salts Neutral, Acidic or Basic?

• 3. NH4Cl
• NH4 is from NH3 , a weak base
• Cl- is from HCl, a strong acid
• H2O
• NH4Cl (s) gt NH4 (aq) Cl- (aq)
• Cl- ions do not react with water, but NH4 ions
  do.
• NH4 H2O º H3O NH3 hydrolysis
• The H3O ions are produced, so the solution is
  acdic.

33
Hydrolysis

• The hydrolysis of an ion is the reaction of an
  ion with water to produce the conjugate acid and
  hydroxide ion or the conjugate base and hydrogen
  ion.
• You need to be able to write equation for
  hydrolysis reaction!

34
Ions As Acids And Bases

• Certain ions can cause an aqueous solution to
  become
• acidic or basic due to hydrolysis.
• Salts of strong acids and strong bases form
  neutral solutions.
• Salts of weak acids and strong bases form basic
  solutions.
• Salts of strong acids and weak bases form acidic
  solutions.
• Salts of weak acids and weak bases form solutions
  that are acidic in some cases, neutral or basic
  in others.

35
Strong Acids And Strong Bases

• Table 4.1, p.141
• Strong acids
• HCl, HBr, HI, HNO3, H2SO4, HClO4
• Strong bases
• Group IA and IIA hydroxides - Memorize!!

36
An Example

• Indicate whether the solutions
• (a) Na2S and (b) KClO4 are acidic, basic or
  neutral.

37
The pH Of A Salt Solution

• What is the pH of 1.0 M NaCN solution?
• Hydrolysis of CN- ions CN- H2O º HCN OH-
• CN- is a conjugate base of HCN. Ka of HCN can be
  
• found. What is Kb for CN-?
• Ka x Kb Kw
• so, Kb Kw/Ka

38
Common Ion Effect Illustrated
blue-violet pH gt 4.6
yellow pH lt 3.0
1.00 M CH3COOH 1.00 M CH3COOH 1.00 M
CH3COONa
39
The Common Ion Effect

• If one solution contains a weak acid and another
  contains the same acid and its conjugate base as
  a second solute, the two solutions have different
  pH values.
• The solution containing both the weak acid and
  its conjugate base has a pH much higher than the
  solution containing only the weak acid.
• The conjugate base is referred to as a common ion
  because it is found in both the weak acid and the
  anion.
• The common ion effect is the suppression of the
  ionization of a weak acid or a weak base by the
  presence of a common ion from a strong
  electrolyte.

40
The Common Ion Effect- An Example

• Calculate the pH of 1.00 M CH3COOH-1.00 M
  CH3COONa solution.

41
Depicting Buffer Action
42
Buffer Solutions

• A buffer solution is a solution that changes pH
  only slightly when small amounts of a strong acid
  or a strong base are added.
• A buffer contains
• a weak acid with its salt (conjugate base) or
• a weak base with its salt (conjugate acid)
• CH3COOH/CH3COONa
• NH3/NH4Cl

43
How A Buffer Solution Works

• The acid component of the buffer can neutralize
  small added amounts of OH-, and the basic
  component can neutralize small added amounts of
  H3O.
• Pure water does not buffer at all.

44
Henderson-Hasselbalch Equation For Buff Solutions

• conjugate base
• pH pKa log
• weak acid
• If weak acid conjugate base, pH pKa
• Requirements
• The ratio of conjugate base to weak acid is
  between 0.10 and 10
• conjugate base/Ka gt 100, weak acid/Ka gt 100

45
Calculations in Buffer Solutions

• Example 15.17
• A buffer solution is 0.24 M NH3 and 0.20 M NH4Cl.
• What is the pH of this buffer?
• If 0.0050 mol NaOH is added to 0.500 L of this
  solution, what will be the pH?
• Example 15.18
• What concentration of acetate ion in 0.500 M
• CH3COOH produces a buffer solution with pH
  5.00?

46
Buffer Capacity And Buffer Range

• There is a limit to the capacity of a buffer
  solution to neutralize added acid or base, and
  this limit is reached before all of one of the
  buffer components has been consumed.
• In general, the more concentrated the buffer
  components in a solution, the more added acid or
  base the solution can neutralize.
• As a rule, a buffer is most effective if the
  concentrations of the buffer acid and its
  conjugate base are equal.

47
Acid-Base Indicators

• An acid-base indicator is a weak acid having one
  color and the conjugate base of the acid having a
  different color. One of the colors may be
  colorless.
• HIn H2O º H3O In-
• color 1 color 2
• Acid-base indicators are often used for
  applications in which a precise pH reading isnt
  necessary.
• A common indicator used in introductory chemistry
  laboratories is litmus.

48
Several Common Indicators
49
Types Of Calculations In Acid-Base Equilibria

• pH, pOH Kw H3OOH- pKw pH pOH
• Weak acid and weak base
• A salt aqueous solution hydrolysis, KaKb Kw
• Common ion effect
• Buffer solution HendersonHasselbalch equation

50
Neutralization Reactions

• Neutralization is the reaction of an acid and a
  base.
• Titration is a common technique for conducting a
  neutralization.
• At the equivalence point in a titration, the acid
  and base have been brought together in exact
  stoichiometric proportions.
• The point in the titration at which the indicator
  changes color is called the end point.
• The indicator endpoint and the equivalence point
  for a neutralization reaction can be best matched
  by plotting a titration curve, a graph of pH
  versus volume of titrant.
• In a typical titration, 50 mL or less of titrant
  that is 1 M or less is used.

51
Neutralization ReactionStrong Acid - Strong Base

• Example 15.20
• Calculate the pH at the following points in the
• titration of 20.00 mL of 0.500 M HCl with 0.500 M
  NaOH.
• Before the addition of any NaOH (initial pH).
• After the addition of 10.00 mL of 0.500 M NaOH
• (half-neutralization point).
• After the addition of 20.00 mL of 0.500 M NaOH
  (equivalence point).
• After the addition of 20.20 mL of 0.500 M NaOH
• (beyond the equivalence point).

52
Titration Curve ForStrong Acid - Strong Base
53
Features Of Titration Curve ForStrong Acid -
Strong Base

• pH is low at the beginning.
• pH changes slowly until just before equivalence
  point.
• pH changes sharply around equivalence point.
• pH 7.0 at equivalence point.
• Further beyond equivalence point, pH changes
  slowly.
• Any indicator whose color changes in pH range of
  4 10 can be used in titration.

54
Neutralization ReactionWeak Acid - Strong Base

• Example 15.21
• Calculate the pH at the following points in the
  titration of
• 20.00 mL of 0.500 M CH3COOH with 0.500 M NaOH.
• Before the addition of any NaOH (initial pH).
• After the addition of 8.00 mL of 0.500 M NaOH
  (buffer region).
• After the addition of 10.00 mL of 0.500 M NaOH
• (half-neutralization point).
• After the addition of 20.00 mL of 0.500 M NaOH
• (equivalence point).
• After the addition of 21.00 mL of 0.500 M NaOH
• (beyond the equivalence point).

55
Titration Curve ForWeak Acid - Strong Base
56
Features Of Titration Curve ForWeak Acid -
Strong Base

• The initial pH is higher because weak acid is
  partially ionized.
• At the half-neutralization point, pH pKa.
• pH is greater than 7 at equivalence point because
  the anion of the weak acid hydrolyzes.
• The steep portion of titration curve around
  equivalence point has a smaller pH range.
• The choice of indicators for the titration is
  more limited.

57
Types Of Calculations In Acid-Base Equilibria

• pH, pOH Kw H3OOH- pKw pH pOH
• Weak acid and weak base
• A salt aqueous solution hydrolysis, KaKb Kw
• Common ion effect
• Buffer solution HendersonHasselbalch equation
• Neutralization, titration curve

58
Lewis Acids And Bases

• There are reactions in non-aqueous solvents, in
  the gaseous state, and even in the solid state
  that can be considered acid-base reactions in
  which Brønsted-Lowry theory is not adequate to
  explain.
• A Lewis acid is a species that is an
  electron-pair acceptor and a Lewis base is a
  species that is an electron-pair donor.
• In organic chemistry, Lewis acids are often
  called electrophiles and Lewis bases are often
  called nucleophiles.

59
Summary

• In the Brønsted-Lowry theory an acid is a proton
  donor and a base is a proton acceptor.
• If an acid is strong, its conjugate base is weak
  and if a base is strong, its conjugate acid is
  weak.
• Water is amphiprotic it can be either an acid or
  a base. It undergoes limited self-ionization
  producing H3O and OH-.
• pH -logH3O pOH -logOH- pKw -logKw
• The pH in both pure water and in neutral
  solutions is 7. Acidic solutions have a pH less
  than 7 and basic solutions have a pH greater than
  7.

60
Summary (continued)

• In aqueous solutions at 25 oC, pH pOH 14.00.
• Hydrolysis reactions cause certain salt solutions
  to be either acidic or basic.
• A strong electrolyte that produces an ion common
  to the ionization equilibrium of a weak acid or a
  weak base suppresses the ionization of the weak
  electrolyte.
• Acid-base indicators are weak acids for which the
  acid and its conjugate base have different
  colors.
• In Lewis acid-base theory, a Lewis acid accepts
  an electron pair and a Lewis base donates an
  electron pair.