AP Chemistry
Unit Two - Chemical Reactions incl. Reactions in Solution
List the five common types of
chemical reactions and the
defining characteristics of
each.
For each equation shown,
label the type of reaction.
1. Ca + 2HCl → CaCl2 + H2
2. Na2CO3 → Na2O + CO2
3. C3H8 + 5 O2 → 3CO2 + 4H2O
4. Pb(NO3)2 + 2KCl → 2KNO3 + PbCl2
Complete the following
sentences regarding
synthesis reactions.
A metal combines with a nonmetal to form ____________.
Metallic oxides and water react to form ____________.
Nonmetallic oxides and water react to form ____________.
Predict the products of the
following synthesis reactions.
Balance the equations.
5. K2O + H2O 
6. Al + O2 
Complete the following
sentences regarding
decomposition reactions.
A binary compound will break down to _____________________.
Metallic chlorates decompose into _______________ and ________________.
Metallic carbonates decompose into _______________ and _______________.
Metallic hydroxides decompose int0 _______________ and _______________.
Predict the products of the
following decomposition
reactions. Balance the
reactions.
7. K2CO3 
8. Al(OH)3 
Explain why and how to use
the activity series for metals.
List the general order of the
activity series for metals.
NOTE: You do NOT have
an activity series provided to
you on the AP Exam, so this
general order must be
memorized.
List the order of the activity
series for halogens.
Use the activity series to
predict whether the following
reactions will or will not occur
as written.
9. AgNO3 + Ca
10. Br2 + NaCl
11. Cu +
Mg(ClO3)2
12. Fe + CuSO4
What are the possible
products that drive a double
replacement reaction to
completion?
Memorize the solubility rules.
Label the following
compounds as soluble or
insoluble.
13. KBr
14. Ag2CO3
15. PbCl2
16. KOH
17. Fe(NO3)3
18. BaSO4
What is the Bronsted-Lowry
definitions of an acid and a
base?
List the six strong acids.
List the eight strong bases.
What is the general form of
an acid-base neutralization
reaction?
Define “salt”
Predict the products of the
following acid-base
reactions.
19. Ca(OH)2 + H2SO4
20. HNO2 + Ba(OH)2
Define conjugate acid and
conjugate base.
Label the acid, base,
conjugate acid and
conjugate base in the
following reactions.
21. H2O + CH3COOH → CH3COO- + H3O+
22. HF + OH- → H2O + F-
Define solute, solvent, and
solution.
Define electrolyte, weak
electrolyte and
nonelectrolyte. Explain what
types of substances fall into
each category.
In what form do strong-,
weak-, and non-electrolytes
exist in aqueous solution?
Because of this, how are
strong-, weak-, and nonelectrolytes written in
complete ionic and net ionic
equations.
Complete, then memorize
the following chemical facts
for use when predicting
products and writing net ionic
equations.
H2CO3 decomposes to:
H2SO3 decomposes to:
NH4OH decomposes to:
H2SO4 is a strong acid that completely dissociates to:
For each of the following
reactions, predict the
products and write a
balanced net ionic equation.
23. Solutions of silver nitrate and sodium sulfate are mixed.
24. A piece of aluminum foil is dropped into an iron (II) nitrate solution.
25. Solid calcium phosphate is added to hydrochloric acid.
26. The hydrocarbon benzene (C6H6) is burned in excess oxygen.
27. Calcium metal is heated strongly in nitrogen gas.
28. Chlorine gas is bubbled into a solution of potassium iodide.
Write down a set of rules for
assigning oxidation numbers.
Assign oxidation numbers to
each element in the
substances shown.
Define oxidation and
reduction.
Define oxidizing agent and
reducing agent.
29. H2SO4
H _____
S _____
30. NI3
N _____
I _____
31. MoO42-
Mo _____
O _____
32. MgCr2O7
Mg _____
Cr _____
O _____
O _____
In each of the following
reactions, determine the
element oxidized, the
element reduced, the
oxidizing agent and the
reducing agent.
33. Mg + 2AgNO3 → Mg(NO3)2 + 2Ag
34. Mg(ClO3)2 → MgCl2 + 3O2
Write down the steps used to
balance redox reactions in
acidic solution.
Balance the following redox
reactions that occur in acidic
solution.
35. S2O32- + OCl- → Cl- + S4O62-
36. NO3- + Cu(s) → NO2(g) + Cu2+
Write down the steps used to
balance redox reactions in
basic solution.
Balance the following redox
reaction that occur in basic
solution.
37. SO32- + MnO4- → SO42- + MnO2(s)
Write down the molarity
equation. Define each
variable and list the common
unit for each.
Write down the dilution
equation.
Complete the following
calculations.
38. Calculate the molarity of a solution containing 10.4 g of calcium chloride in 225
mL of solution.
39. In a 0.270M sodium chloride solution, what volume (in mL) contains 2.14 grams
of dissolved solute?
40. Calculate the mass of KI in grams required to prepare 500. mL of a 2.80 M
solution.
41. You have a 4.00M stock solution of HNO3. What volume of stock solution must
be used to prepare 500.0 mL of 0.125M HNO3 solution?
Explain the process of
preparing a solution from a
solid.
Explain the process of
preparing a solution from a
stock solution.
Answer the questions.
42. How would you prepare 100. mL of a 0.50M solution of NaOH from a 6.0M stock
solution?
43. Explain how to prepare 500. mL of a 0.10M solution of K2SO4 from the solid salt.
Complete the following
solution stoichiometry
problems.
44. What are the concentrations of all ions in a 0.50M solution of (NH4)2S?
45. What mass of precipitate forms when excess K3PO4 reacts with 25mL of 1.50M
CaCl2?
46. When 50.0mL of 6.0M AgNO3 and 40.0mL of 5.0M Na2SO4:
(a) What mass of precipitate can be formed?
(b) What is the molar concentration of all ions remaining in solution?
Sketch and label a typical
titration apparatus.
Define equivalence point and
end point. Compare and
contrast the two.
Complete the following
calculation.
47. What volume of 0.50M HCl solution is needed to neutralize completely each of the
following:
(a) 25.0 mL of 0.40M NaOH solution
(b) 10.0 mL of 1.50M Ba(OH)2 solution
Answers:
1. single replacement
2. decomposition
3. combustion
4. double replacement
5. K2O + H2O  2KOH
6. 4Al + 3O2  2Al2O3
7. K2CO3  K2O + CO2
8. 2Al(OH)3  Al2O3 + 3H2O
9. yes
10. no reaction
11. no reaction
12. yes
13. soluble
14. insoluble
15. insoluble
16. soluble
17. soluble
18. insoluble
19. CaSO4 + H2O
20. Ba(NO2)2 + H2O
21. H2O - base, CH3COOH - acid, CH3COO- - conjugate base, H3O+ - conjugate acid
22. HF - acid, OH- - base, H2O - conjugate acid, F- - conjugate base
23. 2Ag+ + SO42- → Ag2SO4
26. 2C6H6 + 15 O2 → 12 CO2 + 6H2O
24. 2Al + 3Fe2+ → 2Al3+ + 3Fe
27. 3Ca + N2 → Ca3N2
25. Ca3(PO4)2 + 6H+ → 2H3PO4 + 3Ca2+
28. Cl2 + 2I- → 2Cl- + I2
29. H: +1
30. N: +3
31. Mo: +6
32. Mg: +2
33. ox: Mg
S: +6
I: -1
O: -2
red: Ag
OA: AgNO3
RA: Mg
34. ox: O
O: -2
Cr: +6
red: Cl
O: -2
OA and RA: Mg(ClO3)2
35. 2S2O32- + OCl- + 2H+ → S4O62- + Cl- + H2O
36. 4H+ + 2NO3- + Cu → 2NO2 + Cu2+ + 2H2O
37. 3SO32- + 2MnO4- + H2O → 3SO42- + 2MnO2 + 2OH38. 0.427M
39. 136 mL of solution
40. 232 g KI
41. 15.6 mL of stock solution
42. Use a volumetric pipette to measure out 8.3 mL of the stock solution. Place in a 100mL volumetric flask. Add water
to the mark, cap and fully mix.
43. Measure out 8.7g of solid K2SO4. Place in a 500mL volumetric flask. Add sufficient water to fully dissolve solid. Add
water to the mark, cap and fully mix.
44. 1.0M NH4+, 0.50M S246. (a) 47g
47. 20. mL HCl
(b) [Ag+] = 0M,
45. 3.9 g of precipitate
[Na+] = 4.4 M,
[NO3-] = 3.3 M,
[SO42-] = 0.56M
48. 60. mL HCl