1. Acid and Base Equilibrium
Chapter 16 Brown LeMay

2. Basic Concepts
Acids – sour or tart taste, electrolytes, described by Arrhenius as sub that inc. the H+ conc.
Bases – bitter to the taste, slippery, electrolytes, describes by Arrhenius as sub that inc. the OH- conc.

3. Dissociation of Water
Pure water exists almost entirely of water molecules. It is essentially a non-electrolyte.
Water ionizes to a small extent – auto-ionization
The equilibrium expression is
H20(l) <-> H+(aq) + OH-(aq)
Kw = [H+] [OH-]

4. Since the H+ ion does not exist alone in water Kw is often expressed
Kw = [H3O+] [OH-]
because water conc is constant it does not appear in the expression
The proton in water

5. Values for Kw and [H+] and [OH-]
Kw = [H3O+] [OH-] = 1.0 x 10-14
1.0 x = [X] [X]
1.0 x = X2
X = 1.0 X 10-7 = [H3O+] = [OH-]

6. The Bronsted – Lowry definition
holds true for situations not involving water
Acids donate protons
Bases accept protons

7. HCl(g) + NH3(g)  NH4+(g) + Cl-(aq)\
donates H+ accepts H+ conj acid conj base
BL-Acid BL-Base
notice that the reaction doesn’t happen in water and that the OH- concentration has not increased

8. Conjugate Acids & Bases and Amphoteric Substances
HNO2(aq) + H2O(l)  NO2-(aq)+H3O+(aq)
conjugate
acid base base acid
donates - accepts protons
note HNO2 is considered a Arrenius acid H+ inc in H20 also a BL because donates a Proton
all Arrenius acids and bases are also BL acids and bases – however all BL acids and bases may or may not be arrenius acids or bases water is not an Arrenius base in this example

9. every acid has a conjugate base formed from the removal of a proton from that acid
every base has a conjugate acid formed from the addition of a proton to that base

10. Amphotheric Substances
NH3(aq)+H20(l)  NH4+(aq) + OH-(aq)
Base Acid Conj Acid Conj Base
p-acc p-donar
note water is acting as an acid in this reaction and a base in the previous one that makes it a amphortic substance

11. Strengths of Acid, Bases
The stronger the acid is the weaker its conjugate base (weaker acid  stronger conj base)
The stronger the base the weaker its conjugate acid (weaker base  stronger conj acid)
Stronger acids and bases ionize to a greater extent than do weak acids and bases.

12. Strong Acids – dissociate completely into ions
HNO3(aq)  H+(aq) + NO3-(aq)
the production of H+ ions from the acid dominates – ignore the H+ donated from the water it is insignificant
The  equilibrium lies so far to the right because HNO3 doesn’t reform
The neg log of the H+ from the acid determines pH.
Strong bases also dissociate completely and the conc of OH- from the base is the only factor considered when calculating pH.

13. Acids and Bases Strong Acids HClO4 prechloric H2SO4 sulfuric
HI hydroiodic
HCl hydochloric
HBr hydrobromic
HNO3 nitric
Strong Bases
GI hydroxides ex. NaOH,KOH
G2 Hydroxides Sr(OH)2
GI Oxides ex.
Na2O, K2O
GI,II Amides ex. KNH2,Ca(NH2)2

14. The pH scale
pH – is defined as the neg log (base-10) of the H+ ion concentration
pH = -log [H+]
What is the pH of a neutral solution
[H+] = 1.0 x 10^-7
pH = -log [1.0 x 10^-7
pH = 7

15. Strong acids and the pH scale
An acidic solution must have a [H+] conc greater than 1.0 X ex 1.0 X 10-6
-log [1.0 X 10-6] = pH 6
What is the pH of a basic solution?
A basic solution is one in which the [OH-]
is greater than 1X10-7.

16. Calc. pH of strong basic solutions
Calculate the pH of a sol that has a [OH-] con. Of 1.0 X10-5
Kw = 1x10-14 = [H+] [OH-]
Kw = 1x10-14 = [H+] [1.0X10-5]
[H+] = 1x = 1.0 X 10-9
1.0X10-5
pH = -log 1.0 x pH = 9

17. The “p” Scale The negative log of a quantity is labeled p (quantity)
Ex: we could reference the quantity of
OH- directly: pOH = -log[OH-]
From the definition of Kw
-log Kw =(-log [H+]) (-log [OH-])= -log 1x10-4
Kw = pH + pOH = 14

18. Calc. pH using the p scale
Ex. OH- conc = 1.0X10-5
-log 1.0X10-5 = 5 = pOH
pH + pOH = 14
pH + 5 = 14
pH = 9

19. Weak Acids partially ionize in aqueous solution
mixture of ions and un-ionized acid in sol.
WA are in equilibrium (H20 is left out because it a pure liquid)
HA(aq) + H20(l)  H30+(aq) + A-(aq)
Ka is the acid dissociation constant
Ka = [H30] [A-] = [H+] [A-]
[HA] [HA]

20. Acid Dissociation Constant
The larger the Ka value the stronger the acid is – more product is in solution

21. Weak Bases
Weak bases in water react to release a hydroxide (OH-) ion and their conjugate acid:
Weak Base(aq) + H2O(l) Conjugate Acid(aq) + OH-(aq)

22. A common weak base is ammonia
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)
Since H2O is a pure liquid it is not expressed in the equilibrium Kb expression
Kb = [NH4+][OH-] (base dissociation
[NH3] constant)
Kb always refers to the equilibrium in which a base reacts with H2O to form the conjugate acid and OH-

23. Lewis Acids and Bases
Review - An Arrhenius acid reacts in water to release a proton - base reacts in water to release a hydroxide ion
In the Bronstead-Lowry description of acids and bases: acid reacts to donate a proton - a base accepts a proton

24. G.N.Lewis defination -
Lewis acid is defined as an electron-pair acceptor
Lewis base is defined as an electron-pair donor

25. In the example with ammonia, the ammonia is acting as a Lewis base (donates a pair of electrons), and the proton is a Lewis acid (accepts a pair of electrons)

26. Lewis is consistent with the description by Arrhenius, and with the definition by Bronstead-Lowry. However, the Lewis description, a base is not restricted in donating its electrons to a proton, it can donate them to any molecule that can accept them.

27. Calculating the pH of a Weak Acid
What is the pH of an aq sol that is M pyruvic acid HC2H3P3? Ka = 1.4x10-4 at 25oC
HC2H3P3  H+ + C2H3P3- I
C -X X X
E X +X X

28. Ka = [H+] [C2H3P3-] pluggin in the values
[HC2H3P3] from the table
1.4x10-4 = X2
( X)
1.4x10-4 (0.0030) = X2
4.2x x10-4x = X2
X x10-4x-4.2x10-7 = 0 a quadratic
ignore the neg sol x = 5.82 x10-4
pH = -log 5.82 x10-4 pH = 3.24

29. Learning Check
What is the pH at 25oC of a solution made by dissolving a 5.00 grain tablet of aspirin (acetylsalicylic acid) in liters of water? The tablet contains 0.325g of the acid HC9H7O4. Ka = 3.3x10-4 mm = 180.2g/l
H+ = 9.4x10-4 pH = 3.03

30. Buffers
A solution that resists changes in pH when a limited amount of an acid or base is added to it.
Buffers contain either a weak acid and it’s conj. base or a weak base and it’s conj. acid.

31. Examples Ex. Weak acid and conj. base equal molar amounts
Strong Acid added
H+ + A-  HA
conj. base weak acid
the conj base interacts with the H+ ions
from the strong acid changing them to a
weak acid

32. Strong base added
OH + HA  HOH + A-
weak acid conj base
the weak acid interacts with the OH- ion from the base to form water and the conj. base
If the concentration of A- and HA are large and the amount of H+ or OH- is small the
solution will be buffered or the change in pH will be minimized.

33. Buffering capacity – the amount of acid or base a buffer can react with before a significant change in pH occurs
Ratio of acid to conj base – unless the ratio is close to 1 ( between 1:10 and 10:1) will be too low to be useful.

34. Calculating the pH of a buffer
Note: a solution of 0.10 M acedic acid and its conj base 0.20 M acetate from sodium acetate is a buffer solution pH = 5.07
Ex. Calc. the pH of a buffer by mixing 60.0 ml of M NH3 with 40.0ml of M NH4Cl.

35. 1St cal the conc. of each species
M = moles/liters
mol of NH3 0.10M = X/0.060 l = mol
mol of NH4 0.10M = X/0.040 l = mol
[NH3] = mol/ l = M
[NH4] = mol/ l = M
NH3 + H2O  NH4+ + OH-
I M O
C -X X X
E x x X

36. Kb = [NH4+] [OH-] ( X)X
1.8X [NH3] (0.060-X)
Ignore X 1.8X10-5 = X X = 2.7X10-5
0.060
-log (2.7x10-5) = 4.6 pOH pH = 9.4
Or using Henderson - Hasselbalch equation
pOH = pKb + log [conj acid] = log ( [0.04]
[B] [0.06])
= 4.6 pOH pH= 9.4

37. What is the pH of a buffer prepared by adding 30. 0ml of
What is the pH of a buffer prepared by adding 30.0ml of .15M HC2H3O2 to 70ml of .2M NaC2H3O2?
[HC2H3O2] = .15M = x/.03 = .0045/.01 = .045
[C2H3O2] = .20M =x/.07 = .0140/.01 =
ka = 1.7x10-5
HHeq pH = pKa + log [conj base]
[acid]
pH = log(.140 = 5.3
.045)

38. Adding an acid or a base to a buffer
Calc the ph of 75ml of the buffer solution of(0.1M HC2H3O2 and 0.2M NaC2H3O2)
to which 9.5 ml of 0.10M HCl has been added. Compare the change to that of adding HCl to pure water.
H+ + C2H3O2 -  HC2H3O2
H+ = 0.10M = n/.0095l = moles
C2H3O2 - = 0.2M = n/.075 = moles
HC2H3O2 = 0.10M = n/.075 = moles

39. Neutralization Reaction
C2H3O2 = C2H3O2 moles – H+ moles
0.0150n n = 0.014
HC2H3O2 = Orginial Conc. + Conc Contributed by reaction
0.075 moles = mol
[C2H3O2] = 0.014mol/0.085l = 0.16M
[HC2H3O2] = / = 0.10M
pH = log (.16/.10) = 4.96