Name_______________________________________________________ Period______
Chemistry Semester II Final Exam Review Guide
**Important note: sigfigs and scientific notation are also on the final but are not in this packet.
Make sure you know how to report answers using the correct number of sigfigs and how to use
scientific notation.**
Writing and Naming Chemical Compounds
Write the correct name or formula:
1. Ca(CO3)
2. KCl
3. FeSO4
4. LiBr
5. MgCl2
6. FeCl3
7. Zn3(PO4)2
8. NH4NO3
9. Al(OH)3
10. ammonium phosphate
11. iron (II) oxide
12. iron (III) oxide
13. calcium chloride
14. potassium nitrate
15. magnesium hydroxide
16. aluminum sulfate
17. copper (II) sulfate
18. lead (IV) chromate
19. Give the simplest formulas of the ionic compounds formed by:
a. calcium and iodine
b. potassium and sulfur
d. strontium and oxygen
e. rubidium and chloride
20. Give the simplest formula of a compound containing:
a. A1+ and X1- ions
b. B2+ and X1- ions
c. C3+ and X1- ions
c. aluminum and oxygen
f. aluminum and fluorine
d. A+ and Y2- ions
e. B2+ and Y2- ions
f. C3+ and Y2- ions
21. Name the following compounds:
a. FeCl2
b. FeCl3
c. CuNO3
22. Give the formulas of:
a. sodium carbonate
b. magnesium sulfate
c.. cobalt (III) nitrate
23. Give the ionic charge for each metal.
a. K
b. Ca
c. Sr
24. Give the ionic charge for each non-metal?
a. S
b. F
d. Li
c. N
e. Al
d. I
e. O
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Writing & Balancing Equations
Balance equations 1-13. Label the type of reaction wherever you can (synthesis, decomposition, double or single
exchange, combustion)
1.
___CoCl2 + ___Fe → ___FeCl3 + ___Co
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2.
___KClO3 → ___KCl + ___O2
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3.
___H2O + ___Mg3N2 → ___Mg(OH)2 + ___NH3
___________________________________________
4.
___FeCl3 + ___NaOH → ___Fe(OH)3 + ___NaCl
___________________________________________
5.
___H2O + ___P2O5 → ___H3PO4
___________________________________________
6.
___ NaNO3 → ___ NaNO2 + ___O2
___________________________________________
7.
___BCl3 + ___P4 + ___H2 → ___BP + ___HCl
___________________________________________
8.
___H2S + ___O2 → ___H2O + ___SO2
___________________________________________
9.
___Fe + ___H2O → ___Fe3O4 + ___H2
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10. ___NH3 + ___H2SO4 → ___ (NH4)2SO4
___________________________________________
11. ___ Fe2O3(s) + ___ H2 (g) → ___ Fe(s) + ___ H2O(l)
___________________________________________
12. ___ K(s) + ___ Br2(l) → ___ KBr(s)
___________________________________________
13. ___ C2H2(g) + ___ O2(g) → ___ CO2(g) + ___ H2O(l)
___________________________________________
Interpreting Chemical Statements
Write the balanced chemical equation from the written descriptions below.
14. Solid aluminum + aqueous hydrogen chloride yield aqueous aluminum chloride + hydrogen gas
15. Calcium hydroxide + phosphoric acid (hydrogen phosphate) yield calcium phosphate + water
16. Solid copper + sulfuric acid (dihydrogen sulfate) yield aqueous copper (II) sulfate + water + gaseous sulfur
dioxide (SO2)
17. Gaseous hydrogen + nitrogen monoxide (NO) yield water + nitrogen gas
18. Assuming the coefficients in the following equations are correct, determine the subscripts, x and y in the
incomplete formulas.
a. N2(g) + 3H2(g) → 2NxHy(g)
x = ________
y = ________
b. N2(g) + 2O2(g) → 2NxOy(g)
x = ________
y = ________
c. CxHy(l) + 8O2(g) → 5CO2(g) + 6H2O(l)
x = ________
y = ________
Chemical Reactions & Conservation of Matter
1. Define the Law of Conservation of Matter.
2. Define physical and chemical changes. Give two examples of each.
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3. Consider the following reaction between sodium cyanide and a solution of hydrochloric acid. Describe in
detail what you would observe if you carried out this reaction.
NaCN (s) + HCl (aq) → NaCl (aq) + HCN (g)
4. What is the correct description of this reaction? Cu (s) + HCl (aq) → CuCl2 (aq) + H2 (g)
a. hydrogen gas reacts with a solution of copper chloride to produce solid copper metal in a solution
of hydrochloric acid
b. a solution of copper metal reacts with solid hydrochloric acid to produce a solution of copper
chloride and hydrogen gas
c. hydrogen gas reacts with solid copper chloride to produce solid copper metal in a solution of
hydrochloric acid
d. solid copper reacts with a solution of hydrochloric acid to produce a solution of copper chloride and
hydrogen gas
5. The following reaction takes place in a beaker. If the mass of the beaker and its contents are measured, how
will the final mass compare to the initial mass? Explain your reasoning.
CuCO3(s) + H2SO4(aq) → CO2(g) + CuSO4(aq) + H2O(l)
Mole Conversions
1. Determine the mass in grams of one mole of
a. C2H2
b. CO2
c. K2S
2. Determine the number of grams in 2.50 mol of
a. C2H2
b. H2S
c. KNO3
3. How many moles are there in 1.00 g of
a. N2
b. HF
c. COCl2
d. KMnO4
4. Glucose has the molecular formula of C6H12O6.
a. What is the molar mass of glucose?
b. What is the mass in grams of one mole of glucose?
c. How many molecules are there in a mole of glucose?
d. How many atoms are there in a mole of glucose?
5. In a molecule of the explosive “TNT” there are 7 carbon atoms, 5 hydrogen atoms, 3 nitrogen atoms and 6
oxygen atoms.
a. What is the molecular formula of TNT?
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b. What is the molar mass of TNT?
c. Convert 1.20 mol of TNT to grams.
Molarity & Dilutions
Calculate the molarities of the following solutions
1. 2.3 moles of sodium chloride in 0.45 liters of solution.
2. 0.75 moles of lithium fluoride in 65 mL of solution.
3. 120 grams of calcium nitrite in 240 mL of solution.
4. 98 grams of sodium hydroxide in 2.2 liters of solution.
5. A chemist starts with 50.0 mL of a 0.40 M NaCl solution and dilutes it to 1000. mL. What is the
concentration of NaCl in the new solution?
6. A chemist wants to make 500. mL of 0.050 M HCl by diluting a 6.0 M HCl solution. How much of the HCl
solution should be used?
7. How much 2.0 M NaCl solution would you need to make 250 mL of 0.15 M NaCl solution?
8. What would be the concentration of a solution made by diluting 45.0 mL of 4.2 M KOH to 250 mL?
9. What would be the concentration of a solution made by adding 250 mL of water to 45.0 mL of 4.2 M KOH?
10. What would be the concentration of a solution made by adding 250 mL of water to 45.0 mL of 4.2 M KOH?
Acids & Bases
1. What are the properties of a base? Name at least 3 common bases.
2. What are the properties of an acid? Name at least 3 common acids.
3. Sodium hydroxide, a base, reacts with phosphoric acid (H3PO4). Write a balanced reaction for this
neutralization.
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Types of Reactions
Write the balanced chemical equation for the reactions described below, remembering the correct states of matter.
Then, indicate the type of reaction in the space provided.
Reaction
Balanced Chemical Equation
Type
Magnesium + oxygen →
Potassium chloride reacts with
lithium to produce lithium chloride
and potassium
Copper II sulfate reacts with
chromium III carbonate
Sodium chloride decomposes into
its elements
CH4 burns in oxygen to produce
carbon dioxide and water
Barium sulfate becomes barium
sulfide and oxygen
Octane (C8H18) burns in oxygen to
produce carbon dioxide and water
Nickel (II) carbonate + strontium
react to form strontium carbonate
and nickel
NaOH + CoCl2 →
Hydrogen peroxide decomposes
into its elements
1. What is the complete balanced equation for the reaction between zinc hydroxide and hydrogen acetate?
Determine whether or not the products will be aqueous or solid (assume the reactants are all aqueous).
2. What are the formulas for the products of the reaction between sodium hydroxide and sulfuric acid
(dihydrogen sulfate)? Determine whether or not the products will be aqueous or solid (assume the
reactants are all aqueous).
3. Write the balanced equation for the following combinations of reactants. Label the products as aqueous or
solid based on the solubility rules. Assume the reactants are aqueous.
a. zinc nitrate and copper I sulfide
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b. ammonium sulfide and barium hydroxide
c. potassium chloride and zinc phosphate
Stoichiometry/Limiting Reactants
1. How many moles are in 6.70 gram of potassium chloride?
2. If you start with 3.23 x 1027 molecules of hydrochloric acid and excess barium hydroxide, how many grams
of barium chloride can you produce?
2 HCl + Ba(OH)2 → BaCl2 + 2 H2O
3. If you have 56 grams of ammonia and 3.2 x 10-3 molecules of oxygen, which is the limiting reagent?
4 NH3 + 3 O2 → 2 N2 + 6 H2O
4. If you have 3 grams of Al and 10.0 moles of oxygen, how many grams of aluminum oxide can you form?
4 Al + 3 O2 → 2 Al2O3
5. Chlorine is used by textile manufacturers to bleach cloth. Excess chlorine is destroyed by its reaction with
sodium thiosulfate, Na2S2O3:
Na2S2O3 + Cl2 + H2O → NaHSO4 + HCl
a. How many moles of Na2S2O3 are needed to react with 0.12mol of Cl2?
b. How many moles of HCl can form from 0.12mol of Cl2?
c. How many moles of H2O are required for the reaction of 0.12mol of Cl2?
d. How many moles of H2O react if 0.24mol HCl is formed?
6. In nitric acid, HNO3, copper metal dissolves according to the following equation:
3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O
How many grams of HNO3 are needed to dissolve 11.45g of Cu?
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7. For the reaction: K + O2 →K2O, How many grams of K2O will be produced from 0.50 g of K and 0.10 g of O2?
8. For the reaction: Na2O + H2O →NaOH
a. What mass of NaOH could be made from 12.4 g of Na2O and 42.1 g of H2O?
b. What would be the limiting reactant if 100 g each of Na2O and H2O were allowed to react?
9. For the reaction: C + 2H2 → CH4
a. How many moles of CH4 can be made from 7.0 moles of H2 and 5.0 moles of C?
b. What mass of CH4 will be made when 10.0 g of H2 reacts with 5.0 g of C?
Heat
1. Definitions:
a. Endothermic
b. Exothermic
c. System
d. Surroundings
e. Specific heat capacity
f.
Heat of fusion
g. Heat of vaporization
h. Thermal equilibrium
2. On the graph to the right label the melting point and boiling point. What phase (or phases) is the substance
at for each of the 5 sections?
A.
B.
C.
D.
E.
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3. In a lab experiment, a student dissolves ammonium acetate, NH4C2H4O2, in water and observes that the
temperature of the solution changes from 25 °C to 10 °C.
a. What is the value of ΔT for this reaction?
b. What is the direction of heat transfer between the system and the surroundings?
c. Is this process endothermic or exothermic?
4. What energy change occurs between water and the surrounding environment as water vapor condenses to
form liquid water? Think about energy transfer between the system and the surroundings.
5. Calculate the final temperature for each of these situations. Show your work.
a. 110. mL of water at 20°C are mixed with 250 mL of water at 85°C.
b. 500 cal of heat are added to 100 mL of water, starting at 5 °C.
c. 500 cal of heat are added to 100 g of aluminum, starting at 5 °C. (Cp = 0.21 cal/g • °C)
c. 2400 cal of heat are added to 1000 g of ice at 0 °C. (ΔHfus = 79.9 cal/g)
6. A 30.4 g copper coin is heated to 400 °C with a Bunsen burner, and then dropped into a beaker containing
100 mL of 25 °C water. The water absorbs 1000 cal of heat when this happens.
a. What is the final temperature of water in the beaker?
b. Based on your answer for part a, what will be the final temperature of the coin?
c. What is the specific heat capacity of the copper coin?
7. When a 2.00 g peanut is burned, it warms 400. mL of water 25.0 °C. What was the energy content of the
peanut?
8. Which of these substances will combust? For each one that does, write out a balanced chemical reaction.
a. Fe
b. NaCl
c. C3H8
d. CH3OH
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