1. Acids and Bases and Oxidation-Reduction
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Chapter 8
Acids and Bases and Oxidation-Reduction
Denniston Topping Caret
6th Edition

2. 8.1 Acids and Bases
Acids: Taste sour, dissolve some metals, cause plant dye to change color
Bases: Taste bitter, are slippery, are corrosive
Two theories that help us to understand the chemistry of acids and bases
Arrhenius Theory
Brønsted-Lowry Theory

3. Arrhenius Theory of Acids and Bases
Acid - a substance, when dissolved in water, dissociates to produce hydrogen ions
Hydrogen ion: H+ also called “protons”
HCl is an acid:
HCl(aq)  H+(aq) + Cl-(aq)

4. Arrhenius Theory of Acids and Bases
Base - a substance, when dissolved in water, dissociates to produce hydroxide ions
NaOH is a base
NaOH(aq)  Na+(aq) + OH-(aq)

5. Arrhenius Theory of Acids and Bases
Where does NH3 fit?
When it dissolves in water it has basic properties but it does not have OH- ions in it
The next acid-base theory gives us a broader view of acids and bases

6. Brønsted-Lowry Theory of Acids and Bases
Acid - proton donor
Base - proton acceptor
Notice that acid and base are not defined using water
When writing the reactions, both accepting and donation are evident

7. Brønsted-Lowry Theory of Acids and Bases
HCl(aq) + H2O(l)  Cl-(aq) + H3O+(aq)
What donated the proton? HCl
Is it an acid or base? Acid
What accepted the proton? H2O
Is it an acid or base? Base
acid
base

8. Brønsted-Lowry Theory of Acids and Bases.
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
base
acid
Now let us look at NH3 and see why it is a base
Did NH3 donate or accept a proton? Accept
Is it an acid or base? Base
What is water in this reaction? Acid

9. Acid-Base Properties of Water
Water possesses both acid and base properties
Amphiprotic – a substance possessing both acid and base properties
Water is the most commonly used solvent for both acids and bases
Solute-solvent interactions between water and both acids and bases promote solubility and dissociation

10. Acid and Base Strength Acid and base strength – degree of dissociation
Not a measure of concentration
Strong acids and bases – reaction with water is virtually 100% (Strong electrolytes)

11. Strong Acids and Bases Strong Acids: Strong Bases:
HCl, HBr, HI Hydrochloric Acid, etc.
HNO3 Nitric Acid
H2SO4 Sulfuric Acid
HClO4 Perchloric Acid
Strong Bases:
NaOH, KOH, Ba(OH)2
All metal hydroxides

12. Weak Acids
Weak acids and bases – only a small percent dissociates (Weak electrolytes)
Weak acid examples:
Acetic acid:
Carbonic Acid:
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
H2CO3(aq) + H2O(l) HCO3-(aq) + H3O+(aq)

13. Weak Bases Weak base examples: Ammonia: Pyridine: Aniline:
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
C5H5NH2(aq) + H2O(l) C5H5NH3+(aq) + OH-(aq)
C6H5NH2(aq) + H2O(l) C6H5NH3+(aq) + OH-(aq)

14. Conjugate Acids and Bases
The acid base reaction can be written in the general form:
Notice the reversible arrows
The products are also an acid and base called the conjugate acid and base
HA + B A- + HB+
acid base

15. Conjugate Acid – what the base becomes after it accepts a proton.
acid base
Conjugate Acid – what the base becomes after it accepts a proton.
Conjugate Base – what the acid becomes after it donates its proton
Conjugate Acid-Base Pair – the acid and base on the opposite sides of the equation
HA + B A- + HB+
base acid

16. Acid-Base Dissociation
HA + B A- + HB+
The reversible arrow isn’t always written
Some acids or bases essentially dissociate 100%
One way arrow is used
HCl + H2O  Cl- + H3O+
All of the HCl is converted to Cl-
HCl is called a strong acid – an acid that dissociates 100%
Weak acid - one which does not dissociate 100%

17. Conjugate Acid-Base Pairs
Which acid is stronger:
HF or HCN? HF
Which base is stronger:
CN- or H2O? CN -

18. Acid-Base Practice
Write the chemical reaction for the following acids or bases in water.
Identify the conjugate acid base pairs.
HF (a weak acid)
H2S (a weak acid)
HNO3 (a strong acid)
CH3NH2 (a weak base)
Note: The degree of dissociation also defines weak and strong bases

19. The Dissociation of Water
Pure water is virtually 100% molecular
Very small number of molecules dissociate
Dissociation of acids and bases is often called ionization
Called autoionization
Very weak electrolyte
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

20. Hydronium Ion H3O+ is called the hydronium ion
In pure water at room temperature:
[H3O+] = 1 x 10-7 M
[OH-] = 1 x 10-7 M
What is the equilibrium expression for:
Remember, liquids are not included in equilibrium expressions
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

21. Ion Product of Water
This constant is called the ion product for water and has the symbol Kw
Since [H3O+] = [OH-] = 1.0 x 10-7 M, what is the value for Kw?
1.0 x 10-14
It is unitless

22. 8.2 pH: A Measurement Scale for Acids and Bases
pH scale – a scale that indicates the acidity or basicity of a solution
Ranges from 0 (very acidic) to 14 (very basic)
The pH scale is rather similar to the temperature scale assigning relative values of hot and cold
The pH of a solution is defined as:
pH = -log[H3O+]

23. A Definition of pH Use these observations to develop a concept of pH
if know one concentration, can calculate the other
if add an acid, [H3O+]  and [OH-] 
if add a base, [OH-]  and [H3O+] 
[H3O+] = [OH-] when equal amounts of acid and base are present
In each of these cases 1 x = [H3O+][OH-]

24. Measuring pH pH of a solution can be:
Calculated if the concentration of either is known
[H3O+]
[OH-]
Approximated using indicator / pH paper that develops a color related to the solution pH
Measured using a pH meter whose sensor measures an electrical property of the solution that is proportional to pH

25. Calculating pH
How do we calculate the pH of a solution when either the hydronium or hydroxide ion concentration is known?
How do we calculate the hydronium or hydroxide ion concentration when the pH is known?
Use two facts:
pH = -log[H3O+]
1 x = [H3O+][OH-]

26. Calculating pH from Acid Molarity
What is the pH of a 1.0 x 10-4 M HCl solution?
HCl is a strong acid and dissociates in water
If 1 mol HCl is placed in 1 L of aqueous solution it produces 1 mol [H3O+]
1.0 x 10-4 M HCl solution has [H3O+]=1.0x10-4M
= -log [H3O+]
= -log [1.0x10-4]
= -[-4.00] = 4.00
pH = -log[H3O+]

27. Calculating [H3O+] from pH
What is the [H3O+] of a solution with pH = 6.00?
4.00 = -log [H3O+]
Multiply both sides of equation by –1
-4.00 = log [H3O+]
Take the antilog of both sides
Antilog –4.00 = [H3O+]
Antilog is the exponent of 10
1.0 x 10-4 M = [H3O+]
pH = -log[H3O+]

28. Calculating the pH of a Base
What is the pH of a 1.0 x 10-3 M KOH solution?
KOH is a strong base (as are any metal hydroxides)
1 mol KOH dissolved and dissociated in aqueous solution produces 1 mol OH-
1.0 x 10-3 M KOH solution has [OH-] = 1.0 x 10-3 M
Solve equation for [H3O+] = 1 x / [OH-]
[H3O+] = 1 x / 1.0 x 10-3 = 1 x 10-11
pH = -log [1 x 10-11]
= 11.00
1 x = [H3O+][OH-]
pH = -log[H3O+]

29. Calculating pH from Acid Molarity
What is the pH of a 2.5 x 10-4 M HNO3 solution?
We know that as a strong acid HNO3 dissociates to produce 2.5 x 10-4 M [H3O+]
pH = -log [2.5 x 10-4]
= 3.6
pH = -log[H3O+]

30. Calculating [OH-] from pH
What is the [OH-] of a solution with pH = 4.95?
First find [H3O+]
4.95 = -log [H3O+]
[H3O+] =
[H3O+] = x 10-5
Now solve for [OH-]
[OH-] = 1 x / x 10-5
= 8.91 x 10-10
pH = -log[H3O+]
1 x = [H3O+][OH-]

31. The pH Scale

32. The Importance of pH and pH Control
Any change that takes place in aqueous solution generally has at least some pH dependence
Agriculture – crops grow best in soil with proper pH
Physiology – blood pH shift of 1 pH is fatal
Acid Rain – lowers pH of water in aquatic systems causing problems for native fishes
Municipal services – sewage treatment and water purification require optimal pH
Industry – many processes require strict pH control for cost-effective production

33. 8.3 Reactions Between Acids and Bases
Neutralization reaction – the reaction of an acid with a base to produce a salt and water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Acid Base Salt Water
Break apart into ions:
H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O
Net ionic equation
Show only the changed components
Omit any ions appearing the same on both sides of equation = Spectator Ions
H+ + OH-  H2O

34. Net Ionic Neutralization Reaction
The net ionic neutralization reaction is more accurately written:
H3O+(aq) + OH-(aq)  2H2O(l)
This equation applies to any strong acid / strong base neutralization reaction
An analytical technique to determine the concentration of an acid or base is titration
Titration involves the addition of measured amount of a standard solution to neutralize the second, unknown solution
Standard solution – solution of known concentration

35. Acid – Base Titration
Buret – long glass tube calibrated in mL which contains the standard solution
Standard solution is slowly added until the color changes
The equivalence point is when the moles of H3O+ and OH- are equal
Indicator – a substance which changes color as pH changes
Flask contains a solution of unknown concentration plus indicator

36. Determine the Concentration of a Solution of Hydrochloric Acid
Place a known volume of acid whose concentration is not known into a flask
Add an indicator, experience guides selection, here phenol red is good
Known concentration of NaOH is placed in a buret
Drip NaOH into the flask until the indicator changes color

37. Determine the Concentration of a Solution of Hydrochloric Acid
Indicator changes color
equivalence point is reached
mole OH- = mole H3O+ present in the unknown acid
Volume dispensed from buret is determined
Calculate acid concentration:
Volume of Hydrochloric Acid: mL
Volume of NaOH added: mL
Concentration of NaOH: M
Balanced reaction shows that HCl and NaOH react 1:1

38. Determine the Concentration of a Solution of Hydrochloric Acid
35.00 mL NaOH x 1L NaOH x mol NaOH
103 mL NaOH L NaOH
= x 10-3 mol NaOH
3.500 x 10-3 mol NaOH x 1 mol HCl
1 mol NaOH
= x 10-3 mol HCl
this amount of HCl is contained in mL
3.500 x 10-3 mol HCl x 103 mL HCl
25.00 mL HCl L HCl
= x 10-1 mol HCl / L HCl = M HCl

39. Polyprotic Substances
The previous examples have the acid and base at a 1:1 combining ratio
Not all acid-base pairs do this
Polyprotic substance – donates or accepts more than one proton per formula unit
Hydrochloric acid is monoprotic, producing one H+ ion for each unit of HCl
Sulfuric acid is diprotic, each unit of H2SO4 produces 2 H+ ions
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2 H2O(l)

40. Dissociation of Polyprotic Substances
Step 1.
H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq)
Step 2.
HSO4-(aq) + H2O(l) SO42-(aq) + H3O+(aq)
In Step 1 H2SO4 behaves as a strong acid – dissociating completely
In Step 2 HSO4-( behaves as a weak acid – reversibly dissociating, note the double arrow

41. 8.4 Acid-Base Buffers
Buffer solution - solution which resists large changes in pH when either acids or bases are added
These solutions are frequently prepared in laboratories to maintain optimum conditions for chemical reactions
Buffers are also used routinely in commercial products to maintain optimum conditions for product behavior

42. A buffer is LeChatelier’s Principle in action
The Buffer Process
Buffers act to establish an equilibrium between a conjugate acid – base pair
Buffers consist of either
a weak acid and its salt (conjugate base)
a weak base and its salt (conjugate acid)
Acetic acid (CH3COOH) with sodium acetate (CH3COONa)
An equilibrium is established in solution between the acid and the salt anion
A buffer is LeChatelier’s Principle in action
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

43. Addition of Base (OH-) to a Buffer Solution
Adding a basic substance to a buffer causes changes
The OH- will react with the H3O+ producing water
Acid in the buffer system dissociates to replace the H3O+ consumed by the added base
Net result is to maintain the pH close to the initial level
The loss of H3O+ (the stress) is compensated by the dissociation of the acid to produce more H3O+
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

44. Addition of Acid (H3O+) to a Buffer Solution
Adding an acidic substance to a buffer causes changes
The H3O+ from the acid will increase the overall H3O+
Conjugate base in the buffer system reacts with the H3O+ to form more acid
Net result is to maintain the H3O+ concentration and the pH close to the initial level
The gain of H3O+ (the stress) is compensated by the reaction of the conjugate base to produce more acid
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

45. Buffer Capacity
Buffer Capacity – a measure of the ability of a solution to resist large changes in pH when a strong acid or strong base is added
Also described as the amount of strong acid or strong base that a buffer can neutralize without significantly changing pH

46. Preparation of a Buffer Solution
Buffering process is an equilibrium reaction described by an equilibrium-constant expression
In acids, this constant is Ka
If you want to know the pH of the buffer, solve for [H3O+], then calculate pH
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

47. Henderson-Hasselbach Equation
Solution of equilibrium-constant expression and pH can be combined into one operation
Henderson-Hasselbach Equation is this combined expression
Using these two equations:
pKa = -log Ka just as pH = -log[H3O+]
pKa = pH – log ( [CH3COO-] / [CH3COOH] )
Henderson-Hasselbach –
pH = pKa + log( [CH3COO-] / [CH3COOH] )
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

48. Henderson-Hasselbach Equation
pH = pKa + log( [CH3COO-] / [CH3COOH] ) can be rewritten
pH = pKa + log ( [conjugate base] / [weak acid])

49. 8.5 Oxidation-Reduction Processes
Oxidation-reduction processes are responsible for many types of chemical change
Oxidation - defined by one of the following
loss of electrons
loss of hydrogen atoms
gain of oxygen atoms
Example: NaNa+ + e-
Oxidation half reaction

50. Oxidation-Reduction Processes
Reduction - defined by one of the following:
gain of electrons
gain of hydrogen
loss of oxygen
Example: Cl + e-  Cl-
Reduction half reaction
Cannot have oxidation without reduction.

51. Oxidation and Reduction as Complementary Processes
Na  Na+ + e-
Cl + e-  Cl-
Na + Cl Na+ + Cl-
Oxidizing Agent
Is reduced
Gains electrons
Causes oxidation
Reducing Agent
Is oxidized
Loses electrons
Causes reduction

52. Applications of Oxidation and Reduction
Corrosion - the deterioration of metals caused by an oxidation-reduction process
Example: rust (oxidation of iron)
4Fe(s) + 3O2(g)  2Fe2O3(s)
Combustion of Fossil Fuels
Example: natural gas furnaces
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)

53. Applications of Oxidation and Reduction
Bleaching
Most bleaching agents are oxidizing agents
The oxidation of the stains produces compounds that do not have color
Example: Chlorine bleach - sodium hypochlorite (NaOCl)

54. Biological Processes Respiration Metabolism
Electron-transport chain of aerobic respiration uses reversible oxidation and reduction of iron atoms in cytochrome c
Metabolism
Break down of molecules into smaller pieces by enyzmes

55. Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Voltaic Cells
Voltaic cell – electrochemical cell that converts stored chemical energy into electrical energy
Let’s consider the following reaction:
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
Is Zn oxidized or reduced?
Oxidized
Copper is reduced

56. Voltaic Cells Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
If the two reactants are placed in the same flask they cannot produce electrical current
A voltaic cell separates the two half reactions
This makes the electrons flow through a wire to allow the oxidation and reduction to occur

57. Voltaic Cell Generating Electrical Current
Cu2+ + 2e-  Cu
Reduction
cathode – electrode where reduction occurs
Zn  Zn2+ + 2e-
Oxidation
anode – electrode
where oxidation occurs

58. Voltaic Cell Generating Electrical Current

59. Silver Battery
Batteries use the concept of the voltaic cell

60. Electrolysis
Electrolysis reactions – uses electrical energy to cause nonspontaneous oxidation-reduction reactions to occur
These reactions are the reverse of a voltaic cell
Rechargeable battery
When powering a device behaves as voltaic cell
With time the chemical reaction nears completion
Battery appears to “run down”
Cell reaction is reversible when battery attached to charger

61. Comparison of Voltaic and Electrolytic Cells