SEMESTER I FINAL EXAM REVIEW
Chemistry I Fall 2012
The Chemistry Final Exam will count for 20% of your final semester grade. The exam will consist
of 60-70 multiple choice questions. The exam will cover material from the entire semester.
Unit 1:
1. What are the SI base units for length, volume, mass, and density? What equipment is used to measure each?
2. Indicate the number of significant figures in each. Round each to have 2 significant figures.
a. 5.11 x 10-5
b. 8.6753
c. 0.0020045
d. 7.805
3. Calculate the density of an object that has a mass of 2.41g and a volume of 11.4mL. Round to the correct
number of significant figures.
4. Determine the volume of a sample of plastic if the mass is 3.0g and the density is 1.24g/cm3. Round to the
correct number of significant figures.
5. Why is density an intensive property?
6. What does the Law of Conservation of Mass state?
7. Draw particle diagrams of two containers with two different: a. masses
b. volumes
c. densities
Unit 2:
8. Draw and describe the particle models for gas, liquid, and solid.
9. a. How much heat is absorbed by 250g of water if its temperature changes from 25.2oC to 101oC?
The heat capacity of water is 4.18 J/goC.
b. Would more or less heat be absorbed by the same amount of alcohol over the same temperature
change? The heat capacity of alcohol is 2.5 J/goC. Justify your answer.
10. Ice heats from -10oC to 25oC.
a. Draw a heating curve. Label the state(s) of matter present in each section and type of energy
(Ei or Eth) being used.
b. Draw an energy bar diagram for the situation.
c. Is this an endothermic or exothermic process?
11. What is the difference between a physical property and a chemical property? Give three examples
of each
12.
Give three examples of physical change and three examples of chemical change.
13. Determine whether the following are exothermic or endothermic processes.
a. Mixing NaOH with water (gets hot)
b. Mixing the chemicals in a cold pack
c. Burning wood
14. List two examples of each: element, compound, homogenous mixture, heterogeneous mixture, physical
change, chemical change.
15. List three observations that indicate a chemical change has occurred.
Unit 3:
16.
17.
18.
19.
List the charge and location in an atom of the protons, neutrons, and electrons.
List the 4 statements of Dalton’s atomic theory. Which are no longer true?
What did Rutherford learn about the atom as a result of his gold foil experiment?
a. What is an isotope—what do they have in common, what is different?
b. List the number of protons, neutrons, and electrons in each isotope: 35Cl and 37Cl
20. Determine the number of protons, neutrons, and electrons in radon-222.
21. There are two naturally occurring isotopes of an element. One has a mass of 35 and is 75.77% abundant in
nature; the other has a mass of 37 and is 24.24% abundant. What is the average atomic mass of this element?
22. For each type of radiation (alpha, beta, gamma) tell what is it, the Greek symbol, the nuclear symbol, and
useful shielding materials.
23. Write and complete the nuclear equation: thorium-234 decays by beta emission.
24. How does fusion work on the sun and what elements are involved?
25. The half-life of a radioactive isotope is 20 minutes. How much of a 5.00 gram sample will remain after 80
minutes?
26. Use the graph below to determine the half-life of the substance.
700
Grams of Material
600
500
400
300
200
100
0
Start
5
10
15
20
25
30
35
40
days days days days days days days days
Time
Unit 4: Chemical Reactions
27.
Use the data below to determine the activity series for the unknown metals: A, B, C, and D. For
each test in the data table, determine the relative reactivity (i.e. Z > X) of the two metals
involved in the reaction.
Test # Solid Metal
Solution
Observations
Relative Reactivity
Bubbles
around
metal
1
A
BI2
No reaction
2
A
DI2
No reaction
3
A
CI2
No reaction
4
B
CI2
Metal turned dark black
5
D
BI2
No reaction
6
D
CI2
b. List the metals from most reactive to least reactive.
28.
Know how to name compounds and write correct chemical formulas from names
a. calcium carbonate
e. CCl4
i. sodium bicarbonate
b. tin (IV) phosphate
f. CaO
j. ammonia
a. boron trichloride
g. Be(NO3)2
k. sodium chloride
b. aluminum sulfide
h. Fe2O3
29. You combine solutions of SrCO3 and BaCl2. A chemical reaction occurs between these
solutions, producing a precipitate. The reaction product is filtered to separate the filtrate from
the precipitate. Flame tests are conducted on each compound and HCl is added to each
compound. The results:
Test Tube Contents
Flame Test
HCl Test
SrCO3
Bright red
Bubbles
BaCl2
Green
No bubbles
Filtrate
Bright red
No bubbles
Precipitate
Green
Bubbles
a. Write a balanced chemical equation (including state symbols) for the chemical reaction that
took place between strontium carbonate and barium chloride.
b. Identify the filtrate and the precipitate. Explain the experimental evidence that helped you
confirm which compound was the filtrate and which was the precipitate. You need to be
specific with what the flame test and HCl tests tell you about the filtrate.
30. Classify the following as chemical reaction or physical change:
a.
propane gas is burned to heat the grill to cook
b.
liquid ethyl alcohol turns into a solid when placed in a low-temperature freezer
c.
a mixture of sugar and water bubbles when yeast is added
d.
solid copper deposits on a piece of aluminum foil when the foil is placed in a blue e.
copper nitrate solution. The blue color of the solution fades.
31.Write the reaction products, balance the equation, then label it with the reaction type.
a. methane gas plus oxygen gas
d. electrolysis of water
b. aluminum metal plus iron (II) oxide e. nickel (II) nitrate plus potassium sulfide
c. sodium metal plus oxygen gas
Unit 5-6: The Mole & Stoichiometry
32.
Calculate the molar mass of the following:
a. Mg
b. N2
c. NaOH
d. Ca(NO3)2
33. How many moles of water are produced when 0.23 mole of H2 react with excess O2? (Hint:
Write the balanced chemical reaction first. Use this same equation for #33-36
34. How many grams of water are produced when 0.50 mole of O2 reacts with sufficient H2?
35. How many molecules of water are produced when 12.0 grams of hydrogen gas are combined
with excess oxygen gas?
36. How many grams of hydrogen gas are used when 16.2 grams of oxygen gas are used to produce
H2O?
37. What is the empirical formula of a compound that is 79.8% carbon and 20.2% hydrogen?
38. Determine the molecular formula of a compound with an empirical formula of CH2O and a
molar mass of 90 g/mol.
39. Determine the molecular formula of a compound if the compound is found to be 94.1% oxygen
and 5.9% hydrogen. The molecular formula mass is 34.02g/mol.
40. Consider the following reaction. CaO (s) + CO2 (g)  CaCO3 (s)
A chemist allows 14.4 g of CaO and 13.8 g of CO2 to react. When the reaction is finished, the
chemist collects 19.4 g of CaCO3. Determine the limiting reactant, theoretical, and percent
yield for the reaction.
Unit Organic
41. Draw a structure of each of the following alkanes.
a.
2-methylbutane
b. 3-ethyl-2methylhexane c. 3-isopropylheptane
d.
2,5-dimethyloctane
42.
Write the structural formulas for each of the two isomers of butane.