1. A subset of chemical reactions
Reactions in solution
A subset of chemical reactions

2. Learning objectives Define solution, solute, and solvent
Distinguish among strong, weak and non-electrolyte
Identify strong acids and strong bases
Apply solubility rules to prediction of precipitate formation
Classify types of chemical reaction
Predict course of reaction based on activity series
Define oxidation and reduction
Identify oxidizing and reducing agent in reactions
Determine oxidation numbers in ions and compounds

3. Solution A homogeneous mixture of two or more substances
May or may not contain electrolytes
Electrolytes- substances that conduct electricity when dissolved (acids, bases, and ionic compounds)

4. Electrolytes and ionic compounds
All ionic compounds are electrolytes when dissolved in water
Not all ionic compounds are soluble
How do we tell?
Rules to predict solubility
Covalent molecular compounds* are non-electrolytes – no ions produced
*Except acids and bases

5. Dissociation and ionization: same or different?
Ionic compounds dissociate in water
Ions already exist in the solid
Acids or bases* ionize in water
A pure acid or base contains no ions
*Except strong bases like NaOH, Ca(OH)2 are ionic

6. When the weak are made strong
Strong electrolytes are characterized by their nearly complete dissociation in water
Weak electrolytes dissociate to a much smaller extent.

7. Strong, weak or non electrolyte?
All soluble (ionic) salts are strong electrolytes
Strong acids and bases are strong electrolytes
Weak acids and bases are weak electrolytes
Insoluble compounds are non-electrolytes
Molecular compounds are non-electrolytes (except acids/bases)

8. Classifying chemical reactions
Single displacement reactions
Double displacement (metathesis)/ (partner exchange) reactions (in solution)
Acid-base reactions
Oxidation-reduction reactions
Combination reactions
Decomposition reactions

9. Single replacement (displacement)
Element displaces another element from compound (redox)

10. Metathesis (double displacement) reactions involve changing partners
AX + BY = AY + BX
Driven by removal of ions from solution
One example: Precipitation reactions
Two solutions react to form an insoluble solid (precipitate)

11. Precipitation reactions
Does one of the possible cation-anion combinations produce an insoluble salt?
Initial compounds are all soluble
Use solubility rules to investigate
If yes, a precipitate is produced

12. Writing Equations for Rxns in Solution
Molecular equation: normal equation
2KCl (aq) + Pb(NO3)2 (aq) 2KNO3 (aq) + PbCl2 (s)
Ionic equation: shows all ions separately
2K+ (aq) + 2Cl- (aq) + Pb2+ (aq) + 2NO3– (aq)  2K+ (aq) + 2NO3 – (aq) + PbCl2 (s)
Net ionic equation: does not show spectator ions
2Cl- (aq) + Pb2+ (aq)  PbCl2 (s)
Spectator ion – any ion that is present both before and after the reaction
ALL EQUATIONS STILL NEED TO BE BALANCED!

13. Writing Equations for Rxns in Solution
Molecular equation: normal equation
2KCl (aq) + Pb(NO3)2 (aq) 2KNO3 (aq) + PbCl2 (s)
Ionic equation: shows all ions separately
2K+ (aq) + 2Cl- (aq) + Pb2+ (aq) + 2NO3– (aq)  2K+ (aq) + 2NO3 – (aq) + PbCl2 (s)
Net ionic equation: does not show spectator ions
2Cl- (aq) + Pb2+ (aq)  PbCl2 (s)
Spectator ion – any ion that is present both before and after the reaction
ALL EQUATIONS STILL NEED TO BE BALANCED!

14. Solubility rules– you will have to memorize these for the AP exam!

15. For your worksheet Write all answers on another sheet of paper.
For Part 1: show ionic and net ionic equations
For Part 2:
Step 1: Finish the equation by switching the anions
K3PO4(aq) + Al(NO3)3(aq)  KNO3 + AlPO4
make sure you write the
new compounds with the correct subscripts! (use charges)
Step 2: Figure out which product is soluble (aq),
and which is insoluble (s) by using solubility rules
Step 3: Write the ionic and net ionic equations

16. Precipitation Lab
You must show me your pre-lab questions before you may begin.
IMPORTANT: We are using nickel nitrate (Ni(NO3)2) instead of lead II nitrate (Pb(NO3)2)

17. Day 4 – Acid-Base Reactions

18. Know your acids The six strong acids All other acids are weak
HCl, HBr, HI (but not HF)
HNO3 (but not HNO2)
H2SO4 (but not H2SO3)
HClO4 (maybe HClO3)
All other acids are weak

19. Recognizing acids Mineral acids: HCl, HNO3 etc.
Conventionally H appears first in the formula
All strong acids are mineral
May be strong or weak
Organic acids: CH3COOH etc
Harder to spot
Sometimes written with H in front – HCH3CO2
Always weak
Presence of –OH (-SH): necessary but not sufficient
Not all –OH are acidic (CH3OH is not an acid)

20. Recognizing bases
Mineral bases usually distinguished by OH groups – all strong
NaOH, Ca(OH)2
Ammonia, NH3, is an exception – is weak
Organic bases do not contain –(OH) – all weak

21. Neutralization Combine acid with base: ACID + BASE = SALT + WATER
HCl(aq) + NaOH(aq) = H2O(l) + NaCl(aq)
Mg(OH)2(s) + 2HCl(aq) = MgCl2(aq) + 2H2O(l)
Salt contains anion of acid and cation of base:
HCl + NaOH = NaCl + H2O
HCl + KOH = KCl + H2O
HNO3 + KOH = KNO3 + H2O
2HCl + Ca(OH)2 = CaCl2 + 2H2O
HCN + NaOH = NaCN + H2O

22. Acid-base reaction with gas formation
Tums...
HCl(aq) + NaHCO3(aq) = NaCl(aq) + H2CO3(aq)
H2CO3 is unstable:
H2CO3(aq) = H2O(l) + CO2(g)
Bad egg gas:
2HCl + Na2S = H2S(g) + 2NaCl(aq)

23. Oxidation - reduction Oxidation is loss of electrons
Reduction is gain of electrons
Oxidation is always accompanied by reduction
The total number of electrons is kept constant
Oxidizing agents oxidize and are themselves reduced
Reducing agents reduce and are themselves oxidized

25. Oxidation numbers
Oxidation number is the number of electrons gained or lost by the element in making a compound
Metals are typically considered more 'cation-like' and would possess positive oxidation numbers, while nonmetals are considered more 'anion-like' and would possess negative oxidation numbers.

26. Predicting oxidation numbers
Oxidation number of atoms in element is zero
Oxidation number of element in monatomic ion equals charge
Sum of oxidation numbers in compound is zero
Sum of oxidation numbers in polyatomic ion equals charge
F has ON –1
H has ON +1; except in metal hydrides where it is –1
Oxygen is usually –2. Exceptions:
O is –1 in hydrogen peroxide, and other peroxides
O is –1/2 in superoxides KO2
In OF2 O is +2

27. Position of element in periodic table determines oxidation number
G1A is +1
G2A is +2
G3A is +3 (some rare exceptions)
G5A are –3 in compounds with metals, H or with NH4+
Exceptions are compounds with elements to right (e.g. NO2, PF5); in which case use rules 3 and 4.
G6A below O (S, Se etc.) are –2 in binary compounds with metals, H or NH4+
When combined with O or lighter halogen (e.g. SeO2, SF6) use rules 3 and 4.
G7A elements are –1 in binary compounds with metals, H or NH4+ or with a heavier halogen (e.g. Cl in BrCl3)
When combined with O or a lighter halogen, use rules 3 and 4 (e.g. Br in BrCl3 or Cl in ClO4-).

28. Identifying reagents
Those elements that tend to give up electrons (metals) are typically categorized as reducing agents and those that tend to accept electrons (nonmetals) are referred to as oxidizing agents.

29. Identify redox by change in oxidation numbers
Reducing agent increases its oxidation number (Na)
Oxidizing agent decreases its oxidation number (H in H2O)

30. Nuggets of redox processes
Where there is oxidation there is always reduction
Oxidizing agent
Reducing agent
Is itself reduced
Is itself oxidized
Gains electrons
Loses electrons
Causes oxidation
Causes reduction

31. Iron reduces Cu2+ to Cu Iron reduces Cu2+ ions to Cu
Cu does not reduce Fe2+

32. Applying activity series to metals in acids
Mg is higher than H in activity series – forms H2
Cu is lower than H in activity series – no H2 produced

33. Element can be oxidizer and reducer depending on relative positions in activity series
Fe reduces Cu2+
Cu reduces Ag+ (lower activity)
Fe2+ is reduced by Zn (higher activity)

34. Combination reactions
Element + element  compound (redox)
Metal + nonmetal  binary ionic compound
Nonmetal + nonmetal  binary covalent compound
Compound + element  compound (redox)
Compound + compound  compound

35. Decomposition reactions
Compound  element + element (redox)
Compound  element + compound (redox)
Compound  compound + compound

36. Production of a gas
If product is a gas that has low solubility in water, reaction produces gas
Any carbonate with an acid for example