2nd Semester CHEMISTRY Final Exam Review Guide 2014-2015
What will the test look like?
 100 multiple choice problems
 All 2nd semester units, all material will be covered (see list below)
 2 hour time period, no extra time. Some will finish early, some will use whole 2 hours
 Will go in test category, will be worth roughly 10% of overall grade
The units we have covered are listed below with the approximate percentage of questions on the final exam
that will be from that unit:
Chapter
Concept(s)
Percentage of Q's
6&7
Ionic, Metallic, and Covalent Bonding
25
Formulas, Shapes, Polarity, Intermolecular Forces
End of 7
Empirical & Molecular Formulas
Percent Composition
5
8&9
Chemical Reactions
Stoichiometry
30-40
10 & 11
Gas Laws
Stoichiometry with Gases
25-30

You will get all the charts, tables, and gas law equations and constants we have used in class so far:
solubility table, activity series, electronegativity etc…

What you were required to memorize, you will still be required to have memorized: polyatomic ions,
common elements, etc....

As with other tests, you will be allowed to use a scientific non-graphing calculator. There are a
limited number of these you may borrow (not enough for everyone). Be prepared and bring one with
you to class. You will also need #2 pencils for the scantron.

You will not get extra time beyond the two hour time block reserved for this final. If you should
finish your final early, please bring something with you to keep you occupied (a magazine, etc.)

As usual, you may not have cell phones, ipods, headphones, etc… out during the test.

You may not leave the room while you are taking the test. Therefore, go to the bathroom BEFORE
the test starts. If you think you will get thirsty, bring a water bottle with you.

If you know you will be absent, schedule a make-up time with me. Please make these arrangements
as soon as possible.
Ch 6& 7: Bond Types, Properties, Structure of Molecules
Bond Types & Properties
 Be able to define and use the following terms: chemical bond, ion, molecule
 Know the 4 different bond types: ionic, metallic, network covalent, molecular covalent
 Given a chemical formula, identify what type of bonding is found in that substance.
 Compare and contrast the role of valence electrons in the 4 different types of bonds
 Identify what types of atoms are involved in the 4 bonding types (metals, non-metals, etc…)
 Describe the general properties of a certain type of bond (conducts electricity, dissolves, etc…)
 Given the properties of an unknown substance, identify type of bonding
Ionic Bonds
 What are ions? How do ions differ from neutral atoms?
 Predict ion charge for a given element. Relate ion charge to number of electrons gained or lost.
 Explain the relationship between ion charge, electron configurations, and the noble gases.
 Explain how cations and anions are formed.
 Predict the number of electrons an atom will gain or lose in an ionic bond
2+
 Write a correct symbol for any ion (ex: calcium ion = Ca )
 Given a metal and non-metal combination, write the correct chemical formula for the ionic compound. (ex:
when Ca and F combine, what compound will they make? Answer = CaF2)
 Given a chemical formula, correctly name any ionic compound
 Given a name, correctly write the chemical formula for any ionic compound
 Determine the charge of a metal cation based on its chemical formula (ex: what’s the charge of the iron
atom in the compound FeO?)
 Be able to explain what a polyatomic ion is.
 Know the names, formulas, and charge of the following polyatomic ions:
o
o
o
o
o
o
Sulfate
Carbonate
Nitrate
Ammonium
Hydroxide
Phosphate
Covalent Bonds:.
 Write a Lewis dot symbol for a single atom
 Predict the number of covalent bonds an element will form. Explain how this is related to the atom’s Lewis
dot structure.
 Write a Lewis dot structure for any molecule (using single, double, or triple bonds)
 Distinguish between bonded pairs and lone pairs in a Lewis Dot structure.
 Check Lewis Structures for correctness: make sure all electrons are paired, octet rule is satisfied, and each
atom forms the correct number of bonds.
 Explain the role of electrons in determining shape. Explain what electron domains are, and how electron
repulsion will create certain molecular shapes.
 Know the five molecular geometries we talked about in class:
o Linear
o Bent
o Trigonal planar
o Pyramidal
o Tetrahedral
Explain what it means for a molecule to be “polar” or “non-polar”
 Explain what electronegativity is.
 What is the difference between a polar-covalent, non-polar covalent, and ionic bond? What happens to the
electrons in each type of bond?
o Example: in a non-polar bond the electrons are shared equally between atoms, whereas in a
polar covalent bond the electrons are shared unequally- one atom “pulls harder” on the
electrons
 Be able to use the difference in electronegativity to predict which type of bond a pair of elements will form
(polar, non-polar, or ionic)
 Identify whether a molecule is polar or non-polar, and be able to explain why.
 Identify dipole charges (which sided of the molecule of bond are negative and positive)
 Given a chemical formula or Lewis dot structure, identify (name and/or draw) the 3-dimensional shape a
given molecule will have
 Know the 3 intermolecular forces and their relative strengths.
 Be able to predict which intermolecular forces are present in a given molecule.
 How do intermolecular forces (IMF) affect the properties of a molecule? Relate properties such as
solubility and boiling point to IMFs such as dipoles and hydrogen bonds
 Predict how molecules will orient themselves based on polarity and IMF.
 Be able to name covalent molecules and write formulas based on names

End of Ch 7: Percent Composition, Empirical & Molecular Formulas
 Be able to find the percent composition of a compound
 Be able to use percent composition to determine the empirical formula of a compound
 Be able to use formula (molar) mass and percent composition to determine the molecular formula of
a compound
Ch. 8 & 9: Reactions & Stoichiometry
Reactions:
 Translate a chemical equation into words, and vice versa
 Use appropriate symbols when writing a reaction (, (s), etc…) and know what they represent
 Explain what the following symbols represent: (s), (g), (l), and (aq).
 Describe what a reaction will look like based on a chemical equation
o Describe what you will see- will you bubbles? Will you see a cloudy mixture form?
 Identify reaction type, and/or predict the products of any of these reaction types:
o Single displacement
o Combustion
o Double displacement
o Decomposition
o Synthesis
 Use coefficients to balance chemical equations
 Know the 7 diatomic molecules (HOFBrINCl), and apply this knowledge appropriately when writing
chemical equations
 Use the activity series to determine whether a single displacement reaction will occur
o Solitary metal must be higher in activity than the one it is trying to displace
Stoichiometry:
 Determine the molar mass of an element or compound using the periodic table
 Use the coefficients in a balanced chemical equation to determine the mole ratios in a reaction.
 Be able to do any stoichiometry calculation!
o grams ↔ grams, moles ↔ moles, moles ↔ grams, etc…
 Given a chemical reaction and actual lab data, be able to calculate percent yield.
o Know terms: theoretical yield, actual yield, percent yield
Percent yield =
actual yield___x100
Theoretical yield
Ch 10 & 11: Kinetic Molecular Theory and Gas Laws:
 Explain the Kinetic molecular theory of gases (know different properties of gases)
 What is the effect of temperature on gas particles?
 What is gas pressure? What is it caused by?
 What happens to pressure when you change volume? Using what you know about the motion of
o molecules, explain why this happens.
 How does changing temperature affect volume? How does changing temperature affect pressure?
What causes these relationships? Explain by discussing the motion of gas molecules.
 Be able to convert between Celsius and Kelvin
 Be able to draw examples of barometers at different altitudes.
 What is absolute zero? Be able to theoretically define it in terms of gas particles.
 Be able to use these formulas to find the unknown pressure, volume, or temperature in different
scenarios.
P1 V1 = P2 V2




P1
T1
=
P2
T2
Be able to explain a given phenomenon by discussing the movement of gas molecules and the
relevant variables (P, T, V) that are involved.
Be able to calculate volume, moles, and mass of gases using molar volume
Be able to calculate temperature, pressure, volume, moles, and mass using the Ideal Gas Law: PV =
nRT
Be able to solve stoichiometry problems involving gases
Name:______________________________ Date:___________ Period:__________
Semester 2 Final Exam Review Practice Problems
Ch 6 & 7:
1. How many electrons are lost or gained in forming each ion?
a. Mg2+
b. Brc. Ag+
d. Fe3+
2. Classify each of the following as a covalent compound or an ionic compound.
a. CO2
b. NaCl
c. MgCl2
d. N2
e. H2O
3. What types of elements tend to combine to form covalent compounds? What exactly is a covalent bond?
4. What types of elements tend to combine to form ionic compounds? What exactly is an ionic bond?
5. Will solid copper sulfate conduct electricity? Will aqueous copper sulfate conduct electricity? Explain.
6. Which type of bond results in a “sea of electrons”? Why do we call it that?
7. What type of bonding do you expect there to be
in the following molecules?
a. NH3 (g)
b. NaCl (s)
c. Mg (s)
d. H2O (l)
e. a hard solid which does not dissolve in water and does not conduct electricity
f. a substance which conducts electricity as a solid.
8. What charge of ion will each of the following elements form? Write the symbol for the ion they will
form.
a. Potassium
b. Magnesium
c. Bromine
9. Write the formula (including charge) for each ion.
a. carbonate ion _______
c. sulfate ion _________
e. phosphate ion ________
b. nitrate ion ________
d. hydroxide ion _________ f. ammonium ion ________
10. Explain why transition metals need roman numerals in their compound names.
11. List 3 elements that will combine with Oxygen in a 1:1 ratio. Explain why this is.
12. Write the formulas for these ionic compounds.
a. magnesium oxide_____________
d. sodium sulfide _______________
b. potassium iodide_____________
e. aluminum chloride ____________
c. tin(II) fluoride ______________
f. Iron (II) nitride ______________
g. Ba2+, Cl- _________
h. Ca2+, S2- __________
i. Al3+, O2- __________
j. Ag+, I- __________
k. K+, Br- __________
l. Fe2+, O2-___________
13. Name the following ionic compounds.
a. MnO2 _______________________
e. SrBr2 ___________________________
b. CaCl2 _______________________
f. K2S ____________________________
c. NiCl2 ________________________
g. CuCl2___________________________
14. Write formulas for the following ionic compounds.
a. sodium phosphate ______________
d. potassium nitrate ________________
b. sodium hydroxide ______________
e. ammonium chloride _______________
c. magnesium sulfate ______________
f. potassium carbonate_______________
15. Write formulas for compounds formed from these pairs of ions.
a. NH4+, SO42- _____________
c. barium ion and hydroxide ion _______________
b. K+, NO3-
d. lithium ion and carbonate ion _______________
_____________
16. Name the following compounds.
a. Na2CO3 ________________________
d. FeCl3_____________________________
b. Li2SO4 _________________________
e. K2CO3____________________________
c. Cu(OH)2 ________________________
f. LiNO3_____________________________
17. Draw the Lewis dot structure for nitrogen trichloride.
18. Draw the electron dot configuration for acetylene, C2H2.
19. How many bonding pairs of electrons are in a Lewis dot diagram of PH3?
(A) 1
(B) 2
(C) 3
(D) 4
20. In a single bond, the atoms share
(A) 1 electron
(B) 2 electrons
(C) 3 electrons
(D) 4 electrons
21. If the electronegativity difference between two atoms is extremely large, what type of bond will they
form? What if the electronegativity difference is very small?
22. If the bond in the molecule HI is polar, on which end of the hydrogen iodide molecule would you find a
partial negative charge?
23. Describe how you determine the shape of a molecule. What do “electron domains” have to do with it?
How and why do the “electron domains” affect shape?
24. Draw a structural formula for C4H10O and C2H4. Describe the shape around each central atom. Are the
molecules polar or nonpolar?
25. Draw & name the shape for carbon dioxide, CO2. Is this molecule polar or nonpolar?
26. Draw & name the shape of fluorine monoxide, F2O. Is this molecule polar or nonpolar?
27. What type of bond; nonpolar covalent, polar covalent, or ionic will form between each pair of atoms?
a. Na and O
b. O and O
c. P and O
28. What is the strongest Intermolecular force acting between the following molecules?
a. CO2 _______________________________
b. CH4O _____________________________
c. CH2O _____________________________
d. NF3 _______________________________
29. Which of the molecules from question 28 would evaporate the fastest? The slowest?
30. Which of the following molecules would have the highest boiling point? Circle and explain why.
31. For the molecules below, answer each of the following questions:
A) Draw the Lewis dot structure
B) Name the shape of the molecule
C) Identify the molecule as polar or non-polar
D) Identify the intermolecular force present
E) Explain how you determined which IMF was present
F) Draw at least 3 molecules of each type and show how they would attract to each other
a. CH2O
b. CH4O
c. CF4
d. PBr3
32. Rank the following molecules in order from weakest to strongest intermolecular attractions. Explain your
reasoning for your rankings:
NH3 NF3 N2
End of Ch 7:
33. Find the percent composition of all the elements in NaOH:
34. Find the percent composition of all the elements in (NH4)2S
35. An oxide of chromium is found to have the following % composition: 68.4% Cr and 31.6% O.
Determine this compound’s empirical formula.
36. A 170.00 g sample of an unidentified compound contains 29.84 g sodium, 67.49 g chromium and 72.67
g oxygen. What is the compound’s empirical formula?
37. The empirical formula for trichloroisocyanuric acid, the active ingredient in many household bleaches is
OCHCl. The molar mass of this compound is 232.41 g/mol. What is the molecular formula of
trichloroisocyanuric acid?
38. Phenyl magnesium bromide is used as a Grignard reagent in organic synthesis. Determine its empirical
and molecular formula if its molar mass is 181.313 g/mol and it contains 39.7458% C, 2.77956% H,
13.4050% Mg, and 44.0697% Br.
Ch 8 & 9:
39. Balance and identify the types of reactions below:
a. _____C3H8 + _____O2  _____CO2 + _____H2O ____________________________
b.
_____Al + _____Fe3N2  _____AlN + _____Fe _____________________________
c. _____H2O2  _____H2O + _____O2
_____________________________
d. ____Mg + O2  _____MgO
______________________________
40. Predict the products of the following reactions and balance the equations.
a. _____LiCl + _____F2 
b. _____Cu + _____Fe(NO3)3 
c. _____Ca3(PO4)2 + _____LiCl 
d. _____Na3N 
41. Write a balanced equation showing states of matter for the formation of aqueous sulfuric acid (H2SO4)
from water and sulfur trioxide gas (SO3).
42. Write the balanced equation for the reaction between hydrochloric acid and calcium metal.
43. Write the balanced equation for the reaction between iron(III) chloride and sodium hydroxide.
44. Use the activity series of metals and your knowledge of the relative reactivity of the halogens to predict
whether the following reactions will occur. Write balanced equations for those reactions that do occur.
a. Br2(l) + NaCl(aq) 
b. Ca(s) + Mg(NO3)2(aq) 
c. K(s) + H2SO4(aq) 
d. Zn(s) + NaOH(aq) 
45. Calculate the number of moles of chlorine needed to form 14 moles of iron(III)chloride.
2Fe(s) + 3Cl2(g)  2FeCl3(s)
46. Calculate the mass of oxygen produced from the decomposition of 75.0 g of potassium chlorate.
2KClO3(s)  2KCl(s) + 3O2(g)
47. Calculate the mass of silver needed to react with chlorine to produce 84 g of silver chloride.
Hint: Write a balanced equation first.
48. Calculate the percent yield of Cl2(g) in the decomposition of hydrogen chloride if 25.8 g of HCl produces
13.6 g of chlorine gas. Hint: write the balanced chemical equation, do some stoichiometry!
49. What is the actual amount of magnesium oxide produced when excess carbon dioxide reacts with 42.8 g
of magnesium metal? The percent yield of MgO(s)for this reaction is 81.7%.
2Mg(s) + CO2(g)  2MgO(s) + C(s)
50. What is a diatomic element? List all seven.
Ch 10 & 11:
51. List at least three properties of gases as discussed in class.
52. Draw a picture of 2 barometers:
*One at sea level (760 mm Hg)
* One placed at a higher elevation than sea level.
53. When you spray an aerosol can (such as hairspray) you may notice that the temperature of the container
will drop. Explain why this is(1) – what variables are involved(1)? Which gas law is relevant (1)?
54. Laura is practicing volleyball. The ball is not bouncing right so she pumps some more air into it.
What happens to the mass (weight) of the ball with this change?
(a) The mass increases.
(b) The mass decreases.
(c) The mass stays the same.
55. Convert -38°C to Kelvin.
56. How many degrees Celcius is 585 K? _________________________
57. What type of relationship exists between temperature and pressure?
58. Solve the combined gas law equation for T2:
59. Why can’t we use temperature in °C when we’re doing gas law calculations?
60. What happens to pressure when volume is doubled? What happens to temperature?
61. Draw a picture and explain why the coke can was crushed when it was heated and then placed upside
down in cold water.
Solve the following problems. For each… 1) identify the formula used, 2) show your work, including units,
3) draw a box around your answer
62. A 5.0 liter container of nitrogen has a pressure of 4.8 atm. What volume would be necessary to decrease
the pressure to 2.0 atm?
63. Chlorine gas occupies a volume of 75 mL at 350 K. What volume will it occupy at 700 K?
64. A 3.0 L sample of oxygen gas has a pressure of 0.65 atm at 32 C. What is the new pressure if the
temperature is increased to 200 C?
65. A child is holding a balloon with a volume of 1.75 liters. The temperature of the balloon when it was
filled was 220 C and the pressure was 0.96 atm. If the child were to let go of the balloon and it rose into the
sky where the pressure is 0.68 atm and the temperature is -70 C, what would the new volume of the balloon
be?
66. Find the mass in grams of 2.80 L CO2 at STP.
67. Find the volume in Liters of 3.50 g CO at STP.
68. Calculate the pressure in atmospheres of 2.15 moles of CO2 occupying a volume of 750 mL at 57°C.
69. Find the mass of 3.50 L of NH3 at 0.921 atm and 27°C.
70. Propane burns according to the following equation.: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
What volume of water vapor measured at 250.°C and 1.00 atm is produced when 3.0 L of propane at STP is
burned?