CHEMISTRY YEAR END REVIEW
A. SAFETY and MEASUREMENT
1. Distinguish between exact and measured quantities.
2. Simplify (2.55 x 103) (6.5 x 10-10) =
(1.7 x 10-11)
3. Given that 2.2 lb. = 1.0 kg:
a) convert 35 lb. to kg
b) convert 22 kg to lb.
4. How many significant figures are in each of the following measured quantities:
a) 0.0052
b) 35.02
c) 35.00
d) 5 x 103
5. Add or subtract the following measured quantities (remember sig. fig. rules):
a) 25.02 - 3.511
b) 5.21 x 107 + 6.21 x 108
c) 25.22 +3.336
d) 6.19 x 103 - 1.13 x 102
6.
7.
8.
Multiply or divide the following measured quantities:
a) 25.35 x 32
b) 3.55 x 1022 x 4.1 x
A 25.5 g object has a volume of 35.5 mL. Calculate the density of the object.
Ethanol has a density of 0.789 g/mL.
a) What volume will 32.5 g occupy?
b) What mass will a volume of 301.0 mL have?
B. ATOMIC THEORY
1. Briefly describe the essential features the models of the atom developed by Dalton, Rutherford, Thomson and
Bohr.
2.
3.
4.
5.
6.
Describe the Law of Constant Composition and the Law of Conservation of Mass.
Which of the above laws explains why you can balance reactions? Explain.
Which of the above laws explains why water is always 2 parts H and 1 part O?
Pick one of the atomic theories in question #1 that first explained each of the observations below:
a) The spectral lines in a hydrogen discharge tube. ________________________________
b) Atoms combine in simple whole number ratio to form compounds. _________________
c) Matter contains positive and negative particles. _________________________________
Draw the Bohr (atomic) Diagram for F, Ca, K, Al, C and Br.
F
Ca
K
Al
C
Draw the Bohr (atomic) diagrams for F-, Al3+, P3- and Mg2+.
The number of p+, n0 and e- in Na+ are: __________, __________, __________.
An element has 6 p+, 8 n0 and 6 e-. What is its atomic mass?
Complete the following table:
p+
n0
eAtomic
Atomic
Number
Mass
C
14
C
3
H
10
10
7
15.00 amu
11. Define an isotope.
Br
7.
8.
9.
10.
Charge
0
0
12. Calculate the average atomic mass of magnesium given the following data:
24
25
26
Mg 78.7%
Mg 10.1%
Mg 11.2%
13. Distinguish between a chemical family and a chemical period.
14. Describe what happens to the atomic mass and number as you move to the right within a period on the periodic
table.
15. State three things that elements in the same chemical family have in common.
16. Circle each reactive atom or ion.
a) Na
b) Na+
c) S
d) S2e) He
17. Complete the following chart.
p+
Li+
Mg2+
S
S2Ne
n0
e-
No. of valence e-
C. CHANGES IN MATTER
1. Define physical and chemical change.
2. Circle each example of chemical change.
a) iron rusts
b) silver tarnishes
c) ham is sliced
d) water evaporates
e) ice sublimes
f) gasoline burns
g) salt dissolves
h) steam condenses
i) water is converted into H2 and O2
3. Define exothermic and endothermic change.
4. Describe each a exothermic or endothermic.
a) 2 H2 + O2
2 H2O + 500J
b) 200 J + 2 KClO3 (s)
2 KCl + 3 O2 (g)
c) H2O (s)
H2O (g)
d) H2O (g)
H2O (l)
5. Define the following changes in state: evaporation, freezing, condensation, sublimation and
melting.
6. At 100˚C determine the phase (solid, liquid or gas) for each of the following compounds:
a) m.p. 20˚C b.p. 120˚C
b) m.p. 105˚C
c) b.p. 25˚C 7. Define
element, compound and mixture.
8. Describe each as an element, compound or mixture.
a) A homogeneous brown substance has a constant composition. After heating in a vacuum it gives off a gas
and leaves a silver liquid.
b) A homogeneous purple solid has a constant composition. After heating the solid in a vacuum it evaporates
and re-condenses upon cooling. The molten solid does not decompose when an electrical current is
passed through it.
c) A liquid has a boiling point of 105˚C. After boiling the liquid for 10 minutes, the boiling point changes to
110˚C.
9.
Define metals and nonmetals.
10. Describe how you could separate a mixture of sand and water.
11. Describe how you would separate a mixture of ethanol and water by distillation. Draw a labeled diagram of the
apparatus used. In your explanation include the boiling points of water and ethanol and the terms: distillation
flask, condensation tube and collecting flask.
D. CHEMICAL COMPOUNDS
1. Circle each ionic compound.
a) NaBr
b) KOH
c) HCl
d) NH4C1
2.
Write correct ionic formula for each combination.
a) Mg ; C1
b) NH4 ; O
c) Fe ; S
d) A1 ; SO4
3.
4.
Name each compound in the above question.
Define and give an example of a cation and anion.
e) CH4
f) NH3
5.
Circle each cation.
a) K+
b) C1c) NaBr
d) NH4OH
6. Circle the formulas that have formula units.
a) CCl4
b) H20
c) C12
d) NaF
e) NH4Br
7. Circle the compounds that conduct electricity.
a) Na (s)
b) NaCl (s)
c) NaCl (aq)
d) NaCl (l)
e) CCl4 (l)
f) H2O (l)
g) CH3OH (aq)
h) NH4OH (aq)
8. Write dissociation equations for:
a) Na2CO3
b) Fe2(SO4)3
9. Draw Lewis structures for:
a) CCl4
b) H2
c) CO2
d) C2H2
10. Arrange the compounds into the four categories: ionic, acids, or molecular
Na2SO4
H2SO4
CCl4
C2H6
KOH
NH4OH
CH3COOH
NH4C1
11. Explain why NaCl and CCl4 are stable compounds.
12. Circle the molecular compounds.
a) H2O
b) HCl
13. Name each binary molecular compound.
a) CCl4
b) CO2
c) NaBr
d) NH4OH
c) P2O5
d) SiCl
E. REACTIONS
1. Complete, balance, and identify each reaction type. Write NR for those reactions that do not occur (use
Activity Series to determine this - only for SR reactions). All phases must be included!
a) C6H14
+ O2
CO2
+
H2O
b) Al2(SO4)3 + AgNO3
c) NaCl
+ KOH
d) Au
+ Cu(NO3)2
e) Cu
+ Au(NO3)3
f) H2SO4
+ NaOH
g) HgO
h) Ca
+ H2O
i) Al
+ I2
2. Write a balanced chemical equation to show how solid iron reacts with oxygen from the air to produce solid
iron (III) oxide. Include all phase symbols.
3. Write a normal double replacement for Cu(NO3)2 (aq) + NaOH (aq). Write the total ionic and net ionic equations
for this reaction.
F. MOLE CONCEPT
1. State Avogadro’s number and hypothesis.
2. Define molar mass.
3. If 3.50 moles of an element has a mass of 14.0 g, calculate its atomic molar mass. What is the element?
4.
5.
6.
7.
8.
9.
10.
11.
12.
13.
14.
15.
Calculate the number of moles in 50.0 L of H2 (g).
Calculate the number of grams in 35.2 moles of cobalt.
Calculate the number of atoms in 6.25 moles of copper atoms.
Calculate the number of moles of water molecules that is equal to 7.55 x 10 22 molecules of water.
Calculate the number of formula units of NaCl in 5.60 g of NaCl.
Calculate the mass of 2.65 x 1027 atoms of iron.
Calculate the percent composition of each element in C2H6O.
A compound is 32.38% Na, 22.58% S and 45.07% O. Calculate the empirical formula of the substance.
Check your formula with the table of radicals.
What is the empirical formula of C6H6?
A molecular compound has an empirical formula of CH 2 and a molecular mass of 84 g/mol. Determine the
molecular formula.
How many moles of H2 gas at STP are in 100.0 L?
How many litres of NH3 gas at STP are in 2.5 moles?
G. BONDING
1. Explain how two atoms bond ionically. Explain how two atoms bond covalently. Illustrate your answers with
examples.
2. What does electronegativity refer to?
3. Find the electronegativity of the following atoms from your periodic table: Na, F, Br, S, O, Ca, Fr, Cu.
4. Describe the Hydrogen bonding provide drawings to illustrate.
H. PERIODIC TABLE
1. Draw an outline of the periodic table and indicate where each chemical family is located. State where the
metals and nonmetals are located. Describe the trends of ionization energy and atomic radius.
2. State what is meant by ionization energy and provide an equation to show the ionization of Ca atoms
3. Which is larger metal atoms or metal ions. Explain.
4. Which is larger nonmetal atoms or nonmetal ions. Explain.
5. In which chemical family is each element located? Cl, K,Ca, Xe, Fe,
6.
State the periodic law.
I.
1.
2.
3.
4.
5.
6.
7.
8.
9.
J.
1.
STOICHIOMETRY
How many chlorine atoms are in 5 formula units of MgCl2?
How many moles of Br are in 10.0 moles of AlBr 3?
Given the reaction: 2 H2 + O2
2 H2O How many moles of H2O are produced by a) 5 moles O2 and b) 5.0 g
O2 ?
If the actual yield is 50.0 g and the theoretical yield is 70.0 g, what is the percentage yield?
Given: 2 Fe2O3 + 3 C
4 Fe + 3 CO2. If 50.0 g of Fe2O3 completely reacts with excess carbon, how many
grams of carbon dioxide are produced?
C3H8 (g) + 5 O2 (g)
3 CO2 (g) + 4 H2O (g). If 205g of C3H8 (g) reacts with excess O2 (g), calculate the
theoretical yield of CO2 (g). If the actual yield of CO2 was 500.0 g, calculate the percentage yield.
2 H2 (g) + O2 (g)
2 H2O (l)
How many grams of H2O are produced when 50.0 L of H2 (g) at STP reacts completely to form water.
2 NaCl 2 Na + Cl2. How many litres of Cl2 (g) are produced at STP when 100.0 g of NaCl completely react.
2 H2 (g) + O2 (g)
2 H2O (l)
If 8 moles of H2 (g) reacts with 3 moles O2 (g), determine the amount (moles) of H2O (l) produced, the limiting
reactant and the amount of excess reactant.
SOLUTION CHEMISTRY
Pick the compounds that would be soluble in CCl4.
a) NaCl
b) CBr4
c) C2H2
d) HF
e) H2O
2. Calculate the number of grams NaCl required to prepare 100.0 mL of a 0.200 M solution.
3. Calculate the number of litres of NaCl required to obtain 0.50 g if the solution is 0.20 M.
4. How many moles NaCl are in a 0.20 L of a 0.50 M solution.
5. 25.0 moles NaOH is dissolved in 50.0 L of H2O. What is the molarity?
6. If 25.0 g NaOH is dissolved in 5.0 L of H2O, what is the concentration?
7. If 25.0 mL of a 6.0 M NaOH solution is diluted to 100.0 mL, what is the new molarity?
8. You are given the equation H2SO4 + 2 NaOH  Na2SO4 + 2 H2O and that 50.0 mL of
0.20M NaOH is required to neutralize 25.00 mL H2SO4. What is the acid concentration?
GAS LAWS
1. What is the final pressure if the temperature of a gas at 101 kPa is increased from 28oC to 48oC and
the volume is increased from 2.5 liters to 5.0 liters? What gas law is used h
2.The volume of 2.5 moles of gas is 15.0 L at a pressure of 98.5 kPa. What is the temperature of the
gas? What gas law is used here?
3. A gas is collected over water in a graduated cylinder. Find the pressure of the gas if water
vapor pressure is 6 kPa and atmospheric pressure is 101 kPa. What gas law is used here?
Thermochemistry
46. Heating and cooling curves. Indicate where melting, vaporization, freezing and condensation are on
each one: What happens to kinetic energy as it relates to temperature?
During a phase change, what remains constant? _
47. Sketch energy diagrams of exothermic and endothermic reactions. Indicate where the reactants,
products, activation energy, H are on each. What is the effect of adding a catalyst on the activation
energy? What is the sign of H for exothermic vs endothermic reactions?
48. H = mass . Cp . T
If a 35.0 g sample changes from 10.0 C to 22.0 C and the heat capacity
is 4.180 J/mol C. Find the change in heat?