Monoprotic AcidBase Equilibria

1
Chapter 10

• Monoprotic Acid-Base Equilibria

2
Chapter 6 - Review

• Brønsted-Lowry Acids and Bases
• Conjugate Acids and Bases
• pH
• Strengths of Acids
• Acid Dissociation Constant / Base Hydrolysis
  Constant
• Ka, Kb Relationship

3
Chapter 10

• Calculation of pH from Equilibrium expressions
• Buffers

4
Strong Acids and Bases

• Assumptions thus far
• pH or pOH calculated assuming the complete
  dissociation of the acid or base in water
• The dissociation of water does not contribute to
  the pH or pOH

5
Strong Acids and Bases

• Four Possible Strengths
• Concentrations ? concentrated acid or base
• Concentration ? 10-6
• Concentration ? 10-8
• 10-6 ? Concentration ? 10-8

6
Strong Acids and Bases

• Concentrations ? concentrated acid or base
• As the concentration of the strong acid or base
  increases, the degree of dissociation decreases
• Dissociation directly related to ion
  stabilization
• Higher concentration means fewer water molecules
  to stabilize the charge
• This class need not worry about this scenario

7
Strong Acids and Bases

• Concentration ? 10-6
• Complete dissociation
• Concentration of ion due to dissociation of water
  is trivial with respect to that produced by the
  acid or base
• pH or pOH is calculated assuming the H or
  OH- is the same as that of the acid or base
  prior to dissociation

8
Strong Acids and Bases

• Concentration ? 10-8
• Acid or Base dissociates completely
• The concentration of the ion due to the
  dissociation of the acid or base is trivial
  compared to the concentration of the ion that
  results from water dissociation.
• pH ? 7.00 (not necessarily equal to 7.00 because
  of ionic strengths and activity coefficients)

9
Strong Acids and Bases

• 10-6 ? Concentration ? 10-8
• The effect of the acid or base dissociation and
  the dissociation due to water on the
  concentration of H and OH- are comparable.
• pH or pOH must be calculated using the systematic
  treatment of equilibria

10
Strong Acids and Base pH Calculation Examples

• Calculate the pH of 0.10 M HBr
• Repeat the calculation of the pH of 0.10 M HBr
  but include activity coefficients.
• Calculate the pH of a 0.10 M KOH solution

11
Strong Acids and Base pH Calculation Examples

• Calculate the pH of 1.0 x 10-8 KOH.
• Calculate the pH of 5.0 x 10-4 M HNO3 and 5.0 x
  10-4 (CH3)4NOH-
• Calculate the pH of 2.0 x 10-7 M (CH3)4NOH-
• Using activity coefficients correctly, calculate
  the pH of (a) 0.050 M HBr and 0.050 M NaOH

12
Weak Acids and Bases

• Do not dissociate completely in water.
• Constants are given by Ka, Kb, and pK(a or b)
• pKx - log Kx
• pKw - log Kw - log HOH-
• pKa - log Ka - log (A-H / HA)
• pKb - log Kb - log (BHOH- / B)
• Acid Base Correlation
• The weaker the acid, the stronger its conjugate
  base
• The conjugate base will never be a strong base
  however
• Vice versa for weak bases

13
Weak Acids and Bases

• Acid dissociation constants
• Found in Appendix G
• Structure and constants given are for the fully
  protonated forms of the acids
• If the reaction given is the hydrolysis of the
  base, Kb must be calculated using the
  relationship of KaKb Kw
• Compound name is that for the neutral molecule

14
Appendix G
15
Weak Acids and Bases Using Appendix G

• Using Appendix G, write the structures of
  pyridine and pyridinium nitrate. Write the Kb
  reaction for pyridine and find the values of Kb
  and pKb.

16
Weak Acid Equilibria Solving for pH

• The pH of a weak acid equilibria reaction is
  determined using
• the systematic treatment of equilibria
• An assumption

17
Weak Acid Equilibria Solving for pH

• Find the pH of an acid solution that has a formal
  concentration of 0.050 M HA and a
• Ka 1.07 x 10-3.

18
Fraction of Dissociation (?) - Acid

• Determined as the concentration of dissociated
  anion divided by the total concentration of that
  atom or group of atoms present in solution.
• ? (A- / A- HA)
• The fraction of dissociated anion increases as
  the weak acid is diluted

19
Weak Acid - Examples

• Find the pH of 0.100 M trimethylammonium
  chloride.
• Find the pH and fraction of dissociation of a
  0.0100 M solution of the weak acid HA with Ka
  1.00 x 10-4
• Find the pH and concentrations of cyclohexylamine
  (C6H11NH2) and cyclohexylammonium ion (C6H11NH3)
  in a 0.020 M solution of cyclohyxylammonium
  iodide.

20
Weak Base Equilibria Solving for pH

• The pH of a weak base equilibria reaction is
  determined similar to that of a weak acid
  equilibria
• the systematic treatment of equilibria
• An assumption

21
Weak base Equilibria Solving for pH

• Use the systematic approach to solve for the pH
  of a weak base.

22
Fraction of Dissociation (?) - Base

• Determined as the concentration of dissociated
  cation divided by the total concentration of that
  atom or group of atoms present in solution.
• ? (BH / BH B)
• The fraction of dissociated cation increases as
  the weak base is diluted

23
Weak Base Examples

• Find the pH of a solution that is 0.0372 M in B
  if the Kb 2.6 x 10-6
• Find the pH of 0.10 M ammonia.
• Find the pH and concentration of (CH3CH2)2NH and
  (CH3CH2)2NH2 in a 0.030 M solution of
  diethylamine.

24
Buffers

• Buffer A mixture of an acid and its conjugate
  base.
• Buffered Solution A solution that resists
  changes in pH when acids or bases are added
• Importance of Buffers
• Found in all forms of chemistry and biochemistry
• Controls rates of reaction and survival of
  organisms
• Buffer Region on a pH curve where the curve is
  relatively level with regard to pH

25
Buffers

• Why do buffers work?
• Le Chateliers Principle
• Very little reaction occurs to change either
  concentration
• HA dissociates very little, and adding extra A-
  to the solution will make the HA dissociate even
  less.
• A- does not react with water, and addition of HA
  makes A- react even less.

26
pH Calculations for Buffers

• Henderson-Hasselbalch equation
• Merely a rearrangement of the Ka equilibrium
  expression
• pH pKa log (A- / HA) for acids
• pH pKa log (B / BH) for bases

27
Henderson-Hasselbalch Equation

• Real Henderson-Hasselbalch equation requires that
  activities are taken into account.
• pH pKa log (A-?A-) / (HA ?HA)

28
Henderson-Hasselbalch Equation

• Properties
• All Equilibria must be satisfied simultaneously
  in any solution at equilibrium
• For the pH to change by 1 unit, A- / HA must
  change by a factor of 10.
• If HA increases, the pH must decrease
• If A- increases, the pH must increase

29
Henderson-Hasselbalch Equation - Examples

• Sodium hypochlorite (NaOCl, the active ingredient
  in almost all bleaches) was dissolved in a
  solution buffered to a pH 6.20. Find the
  ration OCl-/HOCl in this solution.
• Find the pH of a solution prepared by dissolving
  12.43 g of tris (FM 121.135) plus 4.67 g of tris
  hydrochloride (FM 157.596) in 1.00 L of water
• Note Volume of solution is irrelevant since B
  and BH are dissolved in the same container and
  hence the same volume.

30
Henderson-Hasselbalch Equation - Examples

• Find the pH of a solution prepared from 2.53 g of
  oxoacetic acid, 5.13 g of potassium oxoacetate
  and 103 g of water.
• Write the Henderson-Hasselbalch equation for a
  solution of methylamine. Calculate the quotient
  CH3NH2 / CH3NH3 at (a) pH 4.00, (b) pH
  10.64, and (c) pH 12.00. pKa 10.64

31
Buffering Action

• What is the effect of adding an acid to a buffer
  on pH?
• A buffer resists change in pH, because, the
  buffer consumes the added acid or base
• The limit of relatively unaltered pH occurs at
  the point where the acid consumes all of the B
  or the base consumes all of the BH
• When pH pKa, the buffer has its maximum
  capacity to resist change

32
Buffering Action - Example

• If we add 12.0 mL of 1.00 M HCl to a solution
  containing 12.53 g tris (FM 121.135) and 4.67 g
  of tris hydrochloride (FM 157.596) in 1.00 L,
  what will be the new pH?
• Calculate the pH of a solution prepared by
  dissolving 10.0 g of tris(hydroxymethyl)aminometha
  ne plus 10.0 g of tris hydrochloride in 0.250 l
  of water. What will the new pH be if 10.5 mL of
  0.500 M NAOH is added? pKa 8.08

33
Preparing a Buffer

• Ideally!
• Calculate the amount of acid or base to be added
  to the buffer compound to obtain the desired pH.
• How many milliliters of 0.500 M NaOH should be
  added to 10.0 g of tris hydrochloride to give a
  pH of 7.60 in a final volume of 250 mL? pKa
  8.08

34
Preparing a Buffer

• Calculation Errors!
• Ignored Activity Coefficients
• The pH of a buffer will vary with ionic strength.
  
• Adding an inert salt or altering the volume of a
  buffer solution will change the pH
• Temperature may not be just right
• Most buffers exhibit a noticeable dependence of
  pKa on termperature

35
Preparing a Buffer

• Calculation Errors!
• The approximation that HA FHA and A- FA-
  could be in error
• HA and A- in a dilute solution or at pH extremes
  are not equal to their formal concentrations
• In acidic solutions H ? OH-, so OH- can be
  ignored.
• Vise Versa for basic solutions
• pKa could be reported incorrectly
• Arithmetic error

36
Choosing a Buffer

• Since a buffer is most efficient when pHpKa, it
  is important to choose a buffer that has a pKa as
  close as possible to the desired pH
• A buffers useful pH range, is its pKa 1.
• Table 10-2 lists pKa values for some common
  buffers that are widely used in biochemistry

37
Preparing a Buffer

• Real Life!
• Weigh out desired quantity of buffer compound and
  dissolve it into a beaker using ? 4/5 of the
  desired quantity of water
• Monitor the pH using a calibrated pH electrode
• Add bas or acid to raise or lower the pH
  respectively until the desired pH is attained
• Transfer the solution to a volumetric thoroughly
  rinsing the beaker
• Dilute to mark and mix

38
Buffer Summary

• The Henderson-Hasselbalch equation (with activity
  coefficients) is always true
• Approximations like HA FHA and A- FA are
  not always true
• Over a reasonable range of concentration, the pH
  of a buffer in nearly independent of concentration

39
Chapter 10 - Homework

• Problems 2, 3, 6, 8, 11, 18, 19, 21, 32, 34,
  36, 38