1. Chapter 8 Acid-Base Equilibria
8-1 Brønsted-Lowry Acids and Bases
8-2 Water and the pH Scale
8-3 The Strengths of Acids and Bases
8-4 Equilibria Involving Weak Acids and Bases
8-5 Buffer Solutions
8-6 Acid-Base Titration Curves
8-7 Polyprotic Acids
8-8 Lewis Acids and Bases
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2. Acid and Base Definitions
Arrhenius (Section 4.3)
Acids are H+ donors
Bases are OH- donors
Broadened Definition (Section 4.3)
Acids are substances that increase [H+]
Bases are substances that increase [OH-]
Brønsted-Lowry (Section 8.1)
Lewis (Section 8.8)
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3. Chapter 8 Acid-Base Equilibria
Brønsted-Lowry
Acids
Bases
Conjugate Base -
Conjugate Acid
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4. Brønsted-Lowry Acids and Bases
Acid-Base Equilibria
Brønsted-Lowry Acids and Bases
A Brønsted-Lowry acid is a substance that can donate a hydrogen ion.
A Brønsted-Lowry base is a substance that can accept a hydrogen ion.
In the Brønsted-Lowry Acid and Base concept, acids and bases occur as conjugate acid-base pairs.
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5. Conjugate Base - subtract an H+ from the acid
Conjugate Acid add H+ to the base
Examples
is the conjugate base of H2O
is the conjugated base of H3O+ (called the hydronium ion)
is the conjugated acid of OH-
(or shown as H+) is the conjugate acid of H2O
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6. Conjugate Base - subtract an H+ from the acid
Conjugate Acid add H+ to the base
H2O
- H+
+ H+
OH-
conjugate base of H2O
conjugate acid of H2O
H3O+
Water is the conjugate acid of OH-
Water is the conjugate base of H3O+
Hydroxide ion
Hydronium ion
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7. CH3CO2H + H2O ↔ H3O+ + CH3CO2- Pairs 1 2 2 1
Acetic Acid
Acetate Ion
Pairs
acid
base
acid
base
Point of View #1
acid
base
Conjugate acid of H2O
Conjugate base of CH3CO2H
CH3CO2H + H2O ↔ H3O+ + CH3CO2-
Point of View #2
acid
base
Conjugate base of H3O+
Conjugate acid of CH3CO2-
CH3CO2H + H2O ↔ H3O+ + CH3CO2-
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8. Autoionization of H2O H2O + H2O ↔ H3O+ + OH- Pairs 1 2 2 1
acid
base
acid
base
Point of View #1
acid
base
Conjugate acid of H2O #2
Conjugate base of H2O #1
H2O + H2O ↔ H3O+ + OH-
# 1
# 2
Point of View #2
acid
base
Conjugate base of H3O+
Conjugate acid of OH-
H2O + H2O ↔ H3O+ + OH-
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9. Exercise 8-1:
Trimethylamine (C3H9N) is a soluble weak base with a foul odor (it contributes to the smell of rotten fish). Write the formula of its conjugate acid.
“… the formula of a conjugate acid is obtained by adding H+ to the formula of the base.”
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10. Nomenclature
When H+ is hydrated it is H3O+ and called a hydronium ion.
Often H3O+ is written in a simpler notation H+ +
111.7°
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11. 1.) Water in NH3 serves as an acid
Amphoterism - an ion or molecule can act as an acid or base depending upon the reaction conditions
1.) Water in NH3 serves as an acid
H2O + NH3 ↔ NH OH-
2.) Water in acetic acid serves as a base
H2O + CH3CO2H ↔ H3O CH3CO2-
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12. H2SO4 + CH3CO2H ↔ CH3CO2H2+ + HSO4-
3.) Acetic Acid is also amphoteric, if in the presence of a strong acid serves as a base
H2SO4 + CH3CO2H ↔ CH3CO2H HSO4-
Amphoterism - an ion or molecule can act as an acid or base depending upon the reaction conditions
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13. Water Autoionization of water: 2 H2O(l)  H3O+(aq) + OH-(aq)
[H3O+][OH-]
[H2O]2
KW = 1.0  (at 25oC)
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14. Strong Acids and Bases
A strong acid is one that reacts essentially completely with water to produce H3O+(aq).
Hydrochloric acid (HCl) is a strong acid:
HCl (aq) + H2O(l)  H3O+(aq) + Cl-(aq) (reaction essentially complete)
Dissolving 0.10 mol of HCl in enough water to make 1.0 L of solution gives a final concentration of 0.10 M for H3O+(aq).
[H3O+][OH-] = KW
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15. A strong base is one that reacts essentially completely with water to produce OH-(aq) ions.
Strong Acids and Bases
Sodium hydroxide (NaOH) is a strong base:
Others are NH2- (amide ion) and H- (Hydride ion)
NaOH(s)  Na+(aq) + OH-(aq) (reaction essentially complete)
Dissolving 0.10 mol of NaOH in enough water to make 1.0 L of solution gives a final concentration of 0.10 M for OH-(aq).
[H3O+][OH-] = KW
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16. The pH Function pH = -log10[H3O+] pH < 7 acidic solution
[H3O+] > [OH-]
pH = 7 neutral solution
[H3O+] = [OH-]
pH > 7 basic solution
[H3O+] < [OH-]
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17. The pH Function EXAMPLE 8-3
Calculate the pH (at 25oC) of an aqueous solution that has an OH-(aq) concentration of 1.2  10-6 M.
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18. If converting from pH to [H+]
pH = -log10[H3O+]
If converting from pH to [H+]
[H+] = 10 -pH = 10 -X
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19. Recall [H3O+][OH-] = KW = 10 -14 pH + pOH = pKw = 14 11/24/2018
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20. The pH Function Exercise page 8-4:
Calculate [H3O+] and [OH-] in saliva that has a pH of 6.60 at 25oC.
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21. Strengths of Acids and Bases
Strong acids (completely ionized in water, i.e., 100%) HI
HBr
HClO4 (perchloric acid)
HCl
HClO3 (chloric acid)
H2SO4
HNO3
HA + H2O ↔ H3O+ + A –
HA = generic acid
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22. Table 8-2 Scale of pKa -11 To +14 Known Scale of pKa -11 To +50
E.g., Acetone = 19, methane = 48, ethane = 50
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24. Strengths of Acids and Bases
B + H2O ↔ BH+ + OH-
B = Base
Kb is the Basicity constant
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25. B + H2O ↔ BH+ + OH- BH+ + H2O ↔ B + H3O+ Kb is the Basicity constant
B = Base
Kb is the Basicity constant
BH+ + H2O ↔ B + H3O+
B = Base
Ka is the Acidity constant
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26. B + H2O ↔ BH+ + OH- BH+ + H2O ↔ B + H3O+ B = Base B = Base 11/24/2018
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27. Chapter 8 Acid-Base Equilibria
Base Strength
strong acids have weak conjugate bases
weak acids have strong conjugate bases
The strength of a base is inversely related to the strength of its conjugate acid; the weaker the acid, the stronger its conjugate base, and vice versa
pKa + pKb = pKw
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28. HIn(aq) + H2O(l)  H3O+(aq) + In-(aq)
Indicators (denoted by In)
A soluble compound, generally an organic dye, that changes its color noticeably over a fairly short range of pH.
Typically, a weak organic acid that has a different color than its conjugate base.
HIn(aq) + H2O(l)  H3O+(aq) + In-(aq)
[H3O+][In-]
[HIn]
= Ka
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30. Methyl Red
Bromothymol blue
Phenolphtalein
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32. Equilibria Involving Weak
Acids and Bases
Weak acids Ka < 1
i.e., pKa > 0
H3O+ (hydronium ion)
Ka = 1
pKa =0
HA + H2O ↔ H3O+ + A-
HA is a weak acid,
Ka < 1 or pKa > 1
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35. Problem: (a) Calculate pH and (b) the fraction of CH3CO2H ionized at equilibrium. Assume 1.0M CH3CO2H initially
CH3CO2H + H2O ↔ H3O+ + CH3CO2-
Init. conc.
∆ conc.
Equil. conc.
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36. Assume that y is small (less than 5% of the initial conc.)
A trick to solving:
Assume that y is small (less than 5% of the initial conc.)
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37. (b) Fraction CH3CO2H ionized at equilibrium
Problem: (a) Calculate pH and (b) the fraction of CH3CO2H ionized at equilibrium. Assume 1.0M CH3CO2H initially
CH3CO2H + H2O ↔ H3O+ + CH3CO2-
(b) Fraction CH3CO2H ionized at equilibrium
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38. 1 Weak Bases Similarly NH3 acts as a weak base in H2O Kb = 1.8 x 10-5
H20 + NH3 ↔ NH4+ + OH-
acid1 base2 acid2 base1 1
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39. As before, you can calculate a pH at equilibrium. Assume 0
As before, you can calculate a pH at equilibrium. Assume 0.01M NH3 initially. Calculate the pH of the resulting solution
H20 + NH3 ↔ NH OH-
Init. conc.
∆ conc.
Equil. conc.
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40. As before, you can calculate a pH at equilibrium. Assume 0
As before, you can calculate a pH at equilibrium. Assume 0.01M NH3 initially. Calculate the pH of the resulting solution
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41. Other salts behave similarly, NH4Cl and AlCl3 give acid solutions.
Hydrolysis is a term applied to reactions of aquated ions that change the pH from 7
When NaCl is placed in water, the resulting solution is observed to be neutral (pH = 7)
However when sodium acetate (NaC2H3O2) is dissolved in water the resulting solution is basic
Other salts behave similarly, NH4Cl and AlCl3 give acid solutions.
These interactions between salts and water are called hydrolysis
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42. It is still a Brønsted-Lowry Acid and Base Reaction
Despite the special term, hydrolysis, there is no reason to treat hydrolysis in a special manner.
It is still a Brønsted-Lowry Acid and Base Reaction
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43. Example problem:
Suppose a 0.1 mole solution sodium acetate is dissolved in 1 liter of water. What is the pH of the solution?
Init. conc.
∆ conc.
Equil. conc.
Find Kb
Find [OH-]
Find [H+]
Find pH
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44. Ka x Kb = Kw Example problem: What is the pH of the solution?
CH3CO2- + H2O ↔ CH3CO2H + OH-
Init. conc. 0.1M 0 ~0
∆ conc. - y + y y
Equil. conc. 0.1– y y y
Ka x Kb = Kw
Find Kb
Find [OH-]
Find [H+]
Find pH
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45. Anions Cations Raise pH Lower pH Non-Hyrolyzed Ions (a few)
Hydrolysis of
Result
Anions
Cations
Raise pH
Lower pH
Non-Hyrolyzed Ions (a few)
7 Anions, not hydrolyzed
Cl -, Br -, I -, HSO4-, NO3-, ClO3-, ClO4-
10 Cations, not hydrolyzed
Li+, Na+, K+, Rb+, Sc+, Mg++, Ca++, Sr++, Ba++, Ag+
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46. Anions Cations Raise pH Lower pH
Hydrolysis of
Result
Anions
Cations
Raise pH
Lower pH
Can predict pH of some salts (relative pH)
Na3PO4 is basic
(a non hydrolyzed cation and a hydrolyzed anion)
FeCl3 is acidic
(a hydrolyzed cation and a non hydrolyzed anion)
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47. Chapter 8 Acid-Base Equilibria
Buffer Solutions: important in biochemical and physiological processes
Organisms (and humans) have built-in buffers to protect them against large changes in pH.
Buffers any solutions that maintain an approximately constant pH despite small additions of acids or bases
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48. Death = 7.0 < pH > 7.8 = Death
Buffers any solutions that maintain an approximately constant pH despite small additions of acids or bases
Human blood (pH=7.4) is maintained by a combination of CO3-2, PO4-3 and protein buffers, which accept H+
Death = 7.0 < pH > 7.8 = Death
How Do Buffers Work?
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49. [H+] depends on Ka and the ratio of acid to salt.
How Do Buffers Work?
HA + H2O ↔ H3O+ + A –
HA = generic acid
[H+] depends on Ka and the ratio of acid to salt.
Thus if both conc. HA and A- are large then small addictions of acid or base don’t change the ratio much
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50. H+ depends on Ka and the ratio of acid to salt.
HA + H2O ↔ H3O+ + A –
H+ depends on Ka and the ratio of acid to salt.
Thus if both conc. HA and A- are large then small addictions of acid or base don’t change the ratio much
Try Example 8-11
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53. Extremely important for Organic Chemistry
8-8 Lewis Acids
Extremely important for Organic Chemistry
Recall from Chapter 3 that main group elements in a molecule are surrounded by 8 electrons (the Octet Rule).
Hydrogen by 2 electrons
Carbon is 6C or 1s2 2s2 2p2
Hydrogen is 1H or 1s1
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54. Lewis Structures
Covalent bonds
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55. Lewis Structures
Lone pair
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56. Lewis Structures
Electron deficient
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57. Chapter 8 Acid-Base Equilibria
Lewis Acids-Bases
Lewis Acid is any species that accepts electron pairs
Called electrophiles (electron seeking)
e.g., Boron, Al, Group III elements
BF3, Mg2+, H+, H3O+ BH3, SnCl4, CH3+, AlCl3, HCl
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58. Chapter 8 Acid-Base Equilibria
Lewis Acids-Bases
Lewis Base is any species that donates electrons through coordination to its lone pair
Called Nucleophiles (nucleus seeking)
e.g., Group V, VI, VII elements
Br -1, Cl -1, F -1, H -1, HO -1, :NH3, H2O::
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