You must become familiar with: 1.) the precipitation diagram (fig. 4.3) before you can write precipitation reactions.

1
Chapter 4

• You must become familiar with 1.) the
  precipitation diagram (fig. 4.3) before you can
  write precipitation reactions.

2
Chapter 4

• 2.) The charges of the transition metal ions in
  solution (figure 4.2).
• 3.) Strong acids and bases (figure 4.1) if you
  are to write acid-base equations.

3
Chapter 4

• Only molecular weak acids and bases are
  considered in this chapter acidic and basic ions
  are covered in Chapter 13.

4
Chapter 4

• The half-equation method is the one we will use
  for balancing redox equations. Advantage trains
  you to break down reactions into reduction and
  oxidation useful in Ch. 18

5
Chapter 4

• Note (even though it may be taught in Chem 1)
  it isnt a good idea to construct net equations
  from molecular equation. It is a waste of time
  and isnt what really happens in solution.

6
Solutions

• Dissociation
• Note that ionic compounds (at least to some
  extent) break up in water to produce cations and
  anions.

7
Solutions

• Ex NaOH(s)
• Ex K2CrO4(s)
• (Ionization Equations)

8
Solutions

• Solution Stoichiometry
• Molarity- mols of solute per liter of solution
• Ex. 1 What volume of 12 M HCl must be taken to
  obtain 0.10 mol of HCl?

9
Solutions

• Ex 2 What mass of NaOH is contained in 125 ml of
  6.00 M NaOH?

10
Solutions

• Ex 3 What are the molarities of the aluminum and
  sulfate ions in 0.100 M aluminum sulfate? (use
  dissociation equation)

11
Precip Reactions

• Solubility Rules (fig 4.3) - used in predicting
  the results of precipitation reactions.

12
Precip Reactions

• Ex. 1. Mix solutions of barium nitrate and
  sodium carbonate. What happens? What are the ions
  present? Possible precipitates?

13
Precip Reactions

• Ex. 1. Mix solutions of barium chloride and
  sodium hydroxide. What happens? What are the ions
  present? Possible precipitates?

14
Acid/Base Reactions

• Arrhenius acid and base definitions.
• Strong vs weak acids and bases.
• Ex. HCl vs HF, Calcium hydroxide vs ammonia.
  (draw the dissociation / ionization equations)

15
Net Ionic Equations

• Some ions in a neutralization reaction are
  considered spectator ions. That is, they are
  unchanged after the reaction.

16
Net Ionic Equations

• Chemists often write net ionic equations to show
  only those ions that actually take place in the
  reaction. Ex. Cons. Of mass lab

17
Net Ionic Equations

• The binary acids HCl, HBr, and HI are strong
  acids, the rest are weak.
• Ternary acids with 2 or more oxygens versus
  hydrogens are usually strong.

18
Net Ionic Equations

• Organic acids are weak.
• Polyprotic acids second step always results in a
  weak acid (see next slide)
• Group 1 and 2 bases are strong (usually).

19
Polyprotic Acids

• An acid that is capable of producing more than
  one mole of H.
• Dissociation equations for sulfuric acid and
  phosphoric acid.

20
Net Ionic Equations

• Polyprotic acids - go through two steps of
  ionization
• Ex. Sulfuric acid

21
Net Ionic Equations

• Molecules and weak acids and bases are not
  written in ionic form.
• Soluble salts are written in ionic form, oxides
  and gases are written as molecules.

22
Acid/Base Reactions

• Strong acid and strong base (ex. Hydrochloric
  acid and sodium hydroxide)
• Strong base and weak acid (ex. Hydrofluoric acid
  and calcium hydroxide)

23
Acid/Base Reactions

• Strong acid and weak base (ex. Hydrochloric acid
  and ammonia)

24
Titrations

• Suppose 22.0ml of 0.150 M HCl is reacted with
  32.0ml of Ba(OH)2. What is the molarity of the
  barium hydroxide?

25
Titrations

• It is found that 25.00ml of 0.0800 M Ca(OH)2 is
  required to react with 10.00ml of HCl. What is
  the molarity of the HCl?
• OH- H --gt H2O HCl .400 M

26
Solutions

• Ex 4 What volume of 0.200 M copper (II) sulfate
  solution is required to react with 50.0ml of
  0.100 M NaOH? (use net ionic)

27
Redox Reactions

• Oxidation numbers pseudocharge assigned
  according to arbitrary rules accounts for the
  movement of electrons.

28
Redox Reactions

• Rules
• 1. Elements in the uncombined form are assigned a
  number of zero.

29
Redox Reactions

• 2. Monatomic ions equal their charge number.
• 3. Group one elements in a compound 1, group 2
  2, fluorine -1, hydrogen normally 1,
  oxygen is normally -2.

30
Redox Reactions

• 4. The sum of the oxidation numbers in a molecule
  0. In polyatomic ions, it equals the charge
  number.

31
Redox Reactions

• Ex. What is the oxidation number for sulfur in
  sulfuric acid? Chromium in the dichromate ion?

32
Redox Reactions

• Oxidation an increase in oxidation number.
• Reduction a decrease in oxidation number.
• Ex. Copper and zinc chloride (1/2 reactions)

33
Redox Reactions

• The oxidizing agent is reduced. The reducing
  agent is oxidized. No means yes. Yes means no.

34
Redox Reactions

• Balancing Redox Equations
• Ex. In solution the chlorate ion reacts with the
  iodide ion to produce the chloride ion and solid
  iodine.

35
Redox Reactions

• 1. Split into two half reactions (oxidation and
  reduction).

36
Redox Reactions

• 2. Balance half-equations separately in the
  following order
• A. Balance atoms of element being oxidized or
  reduced.

37
Redox Reactions

• B. Balance oxidation numbers by adding electrons.
• C. Balance charge by adding acid (H) or base
  (OH-). (depends)

38
Redox Reactions

• D. Balance oxygen by adding water molecules.
• E. Combine the half reactions so that the
  electrons cancel.

39
Solutions

• Balance Ex 5 I-(aq) ClO3-(aq) --gt I2(s)
  Cl-(aq)

40
Solutions

• Ex 5 6I-(aq) ClO3-(aq) 6H(aq) --gt 3 I2(s)
  Cl-(aq) 3 H2O

41
Solutions

• Suppose 22.0ml of 0.150 M potassium chlorate is
  required to react with a sample weighing 5.00g.
  What is the percent of the iodide ion in the
  sample?

42
Solutions

• Stoich Practice 1 Fe2(aq) Cr2O72-(aq) --gt
  Fe2(aq) Cr3(aq)

43
Solutions

• An iron ore sample with a mass of 0.9132g is
  dissolved in HCl(aq). The iron obtained is Fe2.
  This solution is titrated with 28.72ml of 0.05051
  M K2Cr2O7. What is the Fe in the original ore?

44
Solutions

• Problem 2
• C2O42-(aq) MnO4- --gt Mn2(aq) CO2

45
Solutions

• 50.0ml of a saturated sodium oxalate solution
  requires 25.8ml of 0.02140 M KMnO4 in acid
  solution. What mass of sodium oxalate would be
  present in 1 L of this saturated solution?

46
Solutions

• Practice Problem 3
• As2O3(s) MnO4-(aq) --gt H3AsO4(aq) Mn2(aq)

47
Solutions

• A KMnO4 solution is standardized by titration
  with solid As2O3. A 0.1078g sample of As2O3
  requires 22.15ml of KMnO4 to complete the
  titration. What is the molarity of the KMnO4
  solution?