1. 5.1 The Nature of Aqueous Solutions
Water
Inexpensive
Can dissolve a vast number of substances
Many substances dissociate into ions
Aqueous solutions are found everywhere
Seawater
Living systems
General Chemistry: Chapter 5
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2. General Chemistry: Chapter 5
Electrolytes
Some solutes can dissociate into ions.
Electric charge can be carried.
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3. General Chemistry: Chapter 5
Chemistry 140 Fall 2002
Types of Electrolytes
Strong electrolyte dissociates completely.
Good electrical conduction.
Weak electrolyte partially dissociates.
Fair conductor of electricity.
Non-electrolyte does not dissociate.
Poor conductor of electricity.
A generalization is helpful:
Essentially all soluble ionic compounds are strong electrolytes.
Most molecular compounds are weak electrolytes or non-electrolytes
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4. Representation of Electrolytes using Chemical Equations
Chemistry 140 Fall 2002
Representation of Electrolytes using Chemical Equations
A strong electrolyte:
MgCl2(s) → Mg2+(aq) + 2 Cl-(aq)
A weak electrolyte:
CH3CO2H(aq) ← CH3CO2-(aq) + H+(aq) →
Strong – complete dissociation
Weak – reversible
CH3OH(aq)
A non-electrolyte:
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5. Three Types of Electrolytes
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6. Solvation: The Hydrated Proton
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7. Relative Concentrations in Solution
MgCl2(s) → Mg2+(aq) + 2 Cl-(aq)
MgCl2(s) → Mg2+(aq) + 2 Cl-(aq)
In M MgCl2:
Stoichiometry is important.
[Mg2+] = M [Cl-] = M [MgCl2] = 0 M
[Mg2+] = M [Cl-] = M [MgCl2] = 0 M
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8. General Chemistry: Chapter 5
EXAMPLE 5-1
Calculating Ion concentrations in a Solution of a Strong Electolyte. What are the aluminum and sulfate ion concentrations in M Al2(SO4)3?.
Write a Balanced Chemical Equation:
Al2(SO4)3 (s) → 2 Al3+(aq) SO42-(aq)
Identify the Stoichiometric Factors :
2 mol Al3+
1 mol Al2(SO4)3
3 mol SO42-
1 mol Al2(SO4)3
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9. General Chemistry: Chapter 5
EXAMPLE 7-3
Aluminum Concentration:
mol Al2(SO4)3
2 mol Al3+
mol Al3+
[Al] =  =
0.0330
1 L
1 L
1 mol Al2(SO4)3
[SO42-] =  =
1 mol Al2(SO4)3
Sulfate Concentration:
1 L
3 mol SO42-
mol Al2(SO4)3
M SO42-
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10. 5-2 Precipitation Reactions
Soluble ions can combine to form an insoluble compound.
Precipitation occurs.
A test for the presence of chloride ion in water.
Ag+(aq) + Cl-(aq) → AgCl(s)
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11. Silver Nitrate and Sodium Iodide
AgNO3(aq)
NaI(aq)
AgI(s)
Na+(aq) NO3-(aq)
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12. General Chemistry: Chapter 5
Net Ionic Equation
Overall Precipitation Reaction:
AgNO3(aq) +NaI(aq) → AgI(s) + NaNO3(aq)
Complete ionic equation:
Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) →
AgI(s) + Na+(aq) + NO3-(aq)
Spectator ions
Ag+(aq) + NO3-(aq) + Na+(aq) + I-(aq) →
AgI(s) + Na+(aq) + NO3-(aq)
Net ionic equation:
Ag+(aq) + I-(aq) → AgI(s)
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13. General Chemistry: Chapter 5
Solubility Rules
Compounds that are soluble:
Alkali metal ion and ammonium ion salts
Nitrates, perchlorates and acetates
Li+, Na+, K+, Rb+, Cs+ NH4+
NO ClO CH3CO2-
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14. General Chemistry: Chapter 5
Solubility Rules
Compounds that are mostly soluble:
Chlorides, bromides and iodides Cl- Br- I-
Except those of Pb2+, Ag+, and Hg22+.
Sulfates SO42-
Except those of Sr2+, Ba2+, Pb2+ and Hg22+.
Ca(SO4) is slightly soluble.
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15. General Chemistry: Chapter 5
5-3 Acid-Base Reactions
Latin acidus (sour)
Sour taste
Arabic al-qali (ashes of certain plants)
Bitter taste
Svante Arrhenius 1884 Acid-Base theory.
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16. Acids and Bases

17. Some Definitions Arrhenius
An acid is a substance that, when dissolved in water, ionizes and increases the concentration of hydrogen ions, H+.
HCl → H+ + Cl-
A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions, OH-. But, Not all bases contain OH-
NaOH → Na+ + OH-

18. Brønsted-Lowry An acid is a proton donor. A base is a proton acceptor.
e.g. ammonia, NH3, is a base.

19. A Brønsted-Lowry acid…
…must have an ionizable or removable (acidic) proton.
A Brønsted-Lowry base…
…must have a pair of nonbonding electrons.

20. In an acid – base reaction, the acid donates a proton (H+) to the base.

21. If a substance can be either…
…it is amphiprotic.
H2O
HCO3-

22. What Happens When an Acid Dissolves in Water?
Water acts as a Brønsted-Lowry base and abstracts a proton (H+) from the acid.
As a result, the conjugate base of the acid and a hydronium ion are formed.

23. Self-ionization of water
H2O + H2O ⇄ H3O+ (aq) + OH- (aq)

24. Autoionization of Water
As we have seen, water is amphoteric.
In pure water, a few molecules act as bases and a few act as acids.
This is referred to as autoionization.
H2O (l) + H2O (l)
H3O+ (aq) + OH- (aq)

25. For ALL aqueous solutions:
Kc = [H3O+] [OH-]
At 25C, Kw = [H3O+]·[OH-] = 1.0· = Kw (at 25 ºC)
[H3O+] = [OH-] the solution is….
[H3O+] > [OH-] the solution is….
[H3O+] < [OH-] the solution is….

26. General Chemistry: Chapter 5
Chemistry 140 Fall 2002
Acids
Acids provide H+ in aqueous solution.
Strong acids completely ionize:
Weak acid ionization is not complete:
HCl(aq)
H+(aq) + Cl-(aq) →
Strong acids completely ionize in solution
Weak acids partially ionize in solution
Compare to the electrolyte strengths. → ←
CH3CO2H(aq)
H+(aq) + CH3CO2-(aq)
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27. Acids
Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water.
Brønsted and Lowry defined them as proton donors.

28. General Chemistry: Chapter 5
Chemistry 140 Fall 2002
Bases
Bases provide OH- in aqueous solution.
Strong bases:
Weak bases:
NaOH(aq)
Na+(aq) + OH-(aq) →
H2O
Strong bases completely dissociate (or nearly so)
Primarily hydroxides of group 1 and some group 2 metals
Certain substances produce ions by reacting with water, not just dissolving in it.
NH4 is a weak base because the eaciton does not go to completion.
MOST bases are weak bases → ←
NH3(aq) + H2O(l)
OH-(aq) + NH4+(aq)
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29. Acids There are only seven strong acids: Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Chloric (HClO3)
Perchloric (HClO4)

30. Acid-Base Reactions
In an acid-base reaction, the acid donates a proton (H+) to the base.

31. Recognizing Acids and Bases.
Chemistry 140 Fall 2002
Recognizing Acids and Bases.
Acids have ionizable hydrogen ions.
CH3CO2H or HC2H3O2
Bases have OH- combined with a metal ion.
KOH
or can be identified by chemical equations
Na2CO3(s) + H2O(l)→ HCO3-(aq) + 2 Na+(aq) + OH-(aq)
Ionizable protons are usually inidicated separately in a formula.
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32. More Acid-Base Reactions
Chemistry 140 Fall 2002
More Acid-Base Reactions
Milk of magnesia Mg(OH)2
Mg(OH)2(s) + 2 H+(aq) → Mg2+(aq) + 2 H2O(l)
Mg(OH)2(s) + 2 CH3CO2H(aq) →
Mg2+(aq) + 2 CH3CO2-(aq) + 2 H2O(l)
Equation 1 is reaction with strong acid
Equation 2 is reaction with weak acid
Note the characteristic formation of water.
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33. More Acid-Base Reactions
Chemistry 140 Fall 2002
More Acid-Base Reactions
Limestone and marble.
CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2CO3(aq)
CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2O(l) + CO2(g)
CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2CO3(aq)
But: H2CO3(aq) → H2O(l) + CO2(g)
CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2O(l) + CO2(g)
This equation shows CO32- acting as a base instead of OH-.
Our definition must be expanded from simple Arrhenius theory (Bronsted-Lowry and Lewis)
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34. Electrochemical Reactions
In electrochemical reactions, electrons are transferred from one species to another.
Mg (s) + HCl (aq) →

35. Oxidation Numbers
In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.

36. Oxidation and Reduction
A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.

37. Oxidation and Reduction
A species is reduced when it gains electrons.
Here, each of the H+ gains an electron, and they combine to form H2.

38. Oxidation and Reduction
What is reduced is the oxidizing agent.
H+ oxidizes Zn by taking electrons from it.
What is oxidized is the reducing agent.
Zn reduces H+ by giving it electrons.

39. Assigning Oxidation Numbers, ON
Elements in their elemental form have an oxidation number of 0.
Mg Cl2 Pb
The oxidation number of a monatomic ion is the same as its charge.
Zn2+ F- Na+

40. HF NaH Nonmetals tend to have negative oxidation numbers
In combination with other elements:
Nonmetals tend to have negative oxidation numbers
Oxygen has an oxidation number of −2, except in the peroxide ion, which has an oxidation number of −1.
MgO Na2O
Hydrogen is −1 when bonded to a metal and +1 when bonded to a nonmetal.
HF NaH

41. Assigning Oxidation Numbers
Fluorine always has an oxidation number of −1.
HF NaF
The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.
ZnBr CuCl2

42. Assigning Oxidation Numbers
The sum of the oxidation numbers in a neutral compound is 0.
HCl H2O K2O
5. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.
CO ClO4-

43. Examples
Determine the oxidation number of sulfur in the following compounds:
H2S
SO42- S8
S2O32-

44. Electrochemical Reactions

45. In a spontaneous Oxidation-Reduction (redox) Reaction electrons are transferred and energy is released
as the reaction proceeds:
Zn strip dissolves…
The blue color due to Cu2+ fades.
Copper metal is deposited.

46. Oxidation and Reduction Half-Reactions
A reaction represented by two half-reactions.
Oxidation:
Zn(s) → Zn2+(aq) + 2 e-
Reduction:
Cu2+(aq) + 2 e- → Cu(s)
Overall:
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)
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47. Corrosion and…

48. Figure: 20-20
Title:
Protection of iron against electrolytic corrosion
Caption:
In the anodic reaction, the metal that is more easily oxidized loses electrons to produce metal ions: in (a), this is iron; in (b), it is zinc. In the cathodic reaction, oxygen gas, which is dissolved in a thin film of water on the metal, is reduced to OH-. Rusting of iron occurs in (a), but not in (b). When iron corrodes, Fe2+ and OH- ions from the half-reactions initiate these further reactions:
Fe2+ + 2OH--> Fe(OH)2(s)
4Fe(OH)2(s) + O2 + 2H2O --> 4Fe(OH)3(s)
2Fe(OH)3(s) --> Fe2O3 • H2O(s) + 2H2O(l)

49. Oxidation and Reduction
O.S. of some element increases in the reaction.
Electrons are on the right of the equation
Reduction
O.S. of some element decreases in the reaction.
Electrons are on the left of the equation.
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50. Oxidation State Changes
Assign oxidation states:
Fe2O3(s) + 3 CO(g) → 2 Fe(l) + 3 CO2(g) D
Fe3+ is to metallic iron.
CO(g) is to carbon dioxide.
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51. Balancing Oxidation-Reduction Equations
Few can be balanced by inspection.
Systematic approach required.
The Half-Reaction (Ion-Electron) Method
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52. General Chemistry: Chapter 5
Balancing in Acid
Write the equations for the half-reactions.
Balance all atoms except H and O.
Balance oxygen using H2O.
Balance hydrogen using H+.
Balance charge using e-.
Equalize the number of electrons.
Add the half reactions.
Check the balance.
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53. EXAMPLE 5-6
Balancing the Equation for a Redox Reaction in Acidic Solution. The reaction described below is used to determine the sulfite ion concentration present in wastewater from a papermaking plant. Write the balanced equation for this reaction in acidic solution..
SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq)
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54. EXAMPLE 5-6 Determine the oxidation states: 4+ 6+ 7+ 2+
SO32-(aq) + MnO4-(aq) → SO42-(aq) + Mn2+(aq)
Write the half-reactions:
SO32-(aq) → SO42-(aq) + 2 e-(aq)
5 e-(aq) +MnO4-(aq) → Mn2+(aq)
Balance atoms other than H and O:
Already balanced for elements.
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55. EXAMPLE 5-6 Balance O by adding H2O:
H2O(l) + SO32-(aq) → SO42-(aq) + 2 e-(aq)
5 e-(aq) +MnO4-(aq) → Mn2+(aq) + 4 H2O(l)
Balance hydrogen by adding H+:
H2O(l) + SO32-(aq) → SO42-(aq) + 2 e-(aq) + 2 H+(aq)
8 H+(aq) + 5 e-(aq) +MnO4-(aq) → Mn2+(aq) + 4 H2O(l)
Check that the charges are balanced: Add e- if necessary.
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56. EXAMPLE 5-6 Multiply the half-reactions to balance all e-:
5 H2O(l) + 5 SO32-(aq) → 5 SO42-(aq) + 10 e-(aq) + 10 H+(aq)
16 H+(aq) + 10 e-(aq) + 2 MnO4-(aq) → 2 Mn2+(aq) + 8 H2O(l) 6 3
Add both equations and simplify:
5 SO32-(aq) + 2 MnO4-(aq) + 6H+(aq) →
5 SO42-(aq) + 2 Mn2+(aq) + 3 H2O(l)
Check the balance!
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57. General Chemistry: Chapter 5
Chemistry 140 Fall 2002
EXAMPLE 7-3 2+ 3+
H2O2(aq) + 2 Fe2+(aq) + 2 H+ → 2 H2O(l) + 2 Fe3+(aq)
Iron is oxidized and peroxide is reduced. 7+
5 H2O2(aq) + 2 MnO4-(aq) + 6 H+ →
8 H2O(l) + 2 Mn2+(aq) + 5 O2(g) 2+
Fe(2) is oxidized to Fe(3). Therefore peroxide is an oxidizing agent. Peroxide is reduced to water.
Mn(7) is reduced to Mn(2). Therefore peroxide is a reducing agent. Peroxide is oxidized to molecular oxygenl.
Manganese is reduced and peroxide is oxidized.
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