Level One Chemistry 2012-2013
Final Exam (Tuesday, 6/18 at 8:30am) and Quarter Four Cumulative Exam (Thursday, 6/13)
Review Sheet
The final exam is worth 12% of your year grade and it contains questions from each unit.
In addition to this list of topics, sample review questions are posted on the class website
Q1 (approximately 10% of quarter 4 cumulative exam)
Unit One- Matter and Measurement
Chapter 2: Properties of matter, Classification, Chemical vs. physical changes
Chapter 3: Reading scales, Metric units, Significant Digits, Density
Chapter 4: Problem solving techniques, Dimensional analysis
Chapter 18: Types of solutions, reading solubility curves, factors affecting solubility and rate of dissolving, colligative
properties
Unit Two- Atomic Structure and the Periodic Table
Chapter 5: History of atomic models, Atomic structure, Periodic table arrangement, isotopes, ions
Chapter 13: Atomic structure, Electron configuration (up to atomic number 20), Lewis dot symbols and valence
electrons
Chapter 14: Period & Group trends (atomic & ionic radius, ionization energy, electronegativity, reactivity)
Section 15.1: Electron configuration of ions
Q2 (approximately 10% of quarter 4 cumulative exam)
Unit Three- Chemical Bonding
Ch 6: Types of compounds, properties of compounds, Names & formulas of ionic compounds, binary molecular
compounds and acids
Chapters 15&16: Ionic & covalent bonding, Lewis structures, bond polarity, shape & polarity of molecules,
intermolecular forces
Chapter 18: like dissolves like
Unit Four- Chemical Reactions
Chapter 8: Writing word and formula equations, Balancing equations, Classifying types of reactions, Solubility,
Predicting products
Section 19.1: Rates of reactions and factors affecting rates
Q3 (approximately 10% of quarter 4 cumulative exam)
Unit Five - Chemical Quantities
Chapter 7: Mole, Conversions between mole, particle, gram, liter, Empirical & molecular formulas, Molar mass, Percent
composition
Chapter 18.2: Molarity, Dilution
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Chapter 9: Stoichiometry, Limiting & excess reagents, Percent yield
Unit Six – States of Matter and Thermochemistry
Chapter 10: Kinetic theory, Pressure, Temperature, Phase diagrams, Changes of state, Heating/Cooling curves
Chapter 11: Endothermic vs. exothermic, Specific heat, Thermochemical equations, Solving for heat transfer
Chapter 19.3: Driving Forces (entropy and enthalpy), Spontanteity of reaction
Q4 (approximately 70% of the Q4 cumulative exam)
Unit Seven- Gases
Chapter 12: Kinetic theory, P,V,T, n relationships and calculations, Dalton’s Law, Graham’s Law
Section 18.1: Factors affecting gas solubility- temperature and pressure
Unit Eight- Equilibrium and Acids/Bases
Section 19.3: Reversible reactions, Equilibrium; writing, calculating, and interpreting Keq
Le Chatelier’s Principle- predict the shift of equilibrium when a stress (pressure, temperature,
concentration) is applied
Chapters 20 & 21: Definitions (Bronsted-Lowry and Arrhenius) & properties of acids & bases (ions present, pH values,
color when phenolphthalein present), writing dissociation and ionization equations, Kw, pH scale
& calculations, strengths of acids and bases related to Ka and Kb, writing neutralization reactions,
calculating the concentration of an acidic or basic solution based on a titration
Unit Nine- Nuclear Chemistry
Chapter 28: Sources of radiation and exposure effects, writing and interpreting nuclear equations, half-life,
Types of nuclear reactions (alpha, beta, gamma, fission, fusion)
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Unit One- Matter and Measurement
1. Perform the following conversions:
a. 50 cm = ________m = _____________mm
b. 95g = _____________kg = _____________mg
c. 117 cm3 = ____________ mL = ____________ L
2. How many significant figures are there in each of these numbers?
a. 0.0120 m
b. 180. Mg
c. 3050 mL
3. Write the following numbers in scientific notation.
a. 0.00161 m
b. 310000 g
4. What is the density of a solid with the mass of 1.02 g and volume of the 0.8111 m3?
5. Complete the following temperature conversions
a. 20C = _____K
b. 422 K = _____C
6. What is the difference between a chemical and physical change? Give an example of each.
7. Classify the following as an element, compound, heterogeneous mixture, homogeneous mixture (solution):
a. Co(s)
b. CO(g)
c. NaCl(aq)
d. oil + water
8. Define:
a. saturated solution
b. unsaturated solution
9. Use the solubility curve below to answer the following questions.
a. How many grams of KCl and dissolve in 100 g
of water at 30oC?
b. 120 g of sodium nitrate is added to 100 g of
water at 40oC, is this solution saturated,
unsaturated or supersaturated?
c. If a saturated solution of KI is cooled from
20oC to 10oC , how many grams of KI will
precipitate?
d. What mass of sodium chloride can dissolve in
200 g of water at 80oC?
10. You have two aqueous solutions of the same concentration. One is a molecular solute, and the other solute is
magnesium chloride. Which solution will have the lower freezing point? Explain.
Unit Two- Atomic Structure and the Periodic Table
11. Define an isotope and an ion. Write the nuclear symbol for two isotopes of potassium and an ion of potassium.
12. Give the charge and location in the atom of each of the following particles:
Particle
Charge
Location in Atom
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Proton
Electron
Neutron
13. Fill in the following table:
Atom or ion
Atomic No.
Hg
K
K+
17
17
Fe+2
S–2
Uranium-235
Mass No.
201
#p
#e
#n
Avg. atomic mass
21
39
17
18
18
18
56
30
14. Give the electron dot (Lewis) structures for:
Ba
S–2
K+
C
15. a. Write out the complete election configuration (1s22s2p6 etc.) of Ar.
b. Diagram its orbital notation (draw orbitals as lines and electrons as arrows. Label all orbitals).
c. Write the shorthand electron configuration of phosphorus.
d. Write the complete electron configuration of chloride ion.
16. On the periodic table, identify:
a. noble gases
b. alkali metals
c. halogens
d. transition metals
e. metalloids f. transuranium elements
17. Define ionization energy. What elements have the lowest ionization energy? highest? Why?
18. Define electronegativity. Which elements have the largest electronegativity? the smallest? Why?
19. a. Which element in Group 1A has the largest atomic radius and why?
b. What element in Period 3 has the largest atomic radius and why?
c. For the same element, rank in order of size, the neutral element, the cation and the anion.
Unit Three- Chemical Bonding
20. Write the correct formulas for the following. Include states of matter when possible (assume room temp).
a. iron (II) hydroxide
b. dihydrogen monoxide
c. silver sulfate
d. nitrous acid
e. carbon tetrachloride
f. ammonium chromate
g. potassium sulfide
21. Name the following compounds:
a. CrCl3
b. Mg(NO3)2
c. HF(aq)
d. Na2SO3
e. PH3
f. Ba(HCO3)2
g. AlI3
22. a. How is an ionic bond formed? What type(s) of elements is(are) needed?
b. How is a covalent bond formed? What type(s) of elements is(are) needed?
23. a. What is polar molecule? Give an example.
b. What is a nonpolar molecule? Give an example.
24. Polar molecules dissolve best in a. polar or
b. non-polar solvents?
25. Draw electron dot (Lewis) structures, name the shape, and determine if the molecule is polar.
a. H2O
Polar
c. NHF2
Polar
molecule?
molecule?
Shape
Shape
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b. CCl4
Polar
molecule?
d. SF2
Shape
Polar
molecule?
Shape
26. a. Which intermolecular force exists in each of the molecules above?
b. Which would have a higher boiling point, H2O or CCl4? Explain your answer.
27. What properties distinguish an ionic solid from a molecular solid?
Unit Four: Chemical Reactions
28. Complete and balance the following equations and write states of matter for reactants and products. Label the type
of reaction.
a. ___ Na + ___ H2O 
b. ___ KClO3 
c. ___ KI + ___ Cl2 
d. ___ AgNO3 + ___ NaBr 
e. ___ NaCl +___ K2O 
29. Write a chemical equation to show the complete combustion of propane (C3H8)
30. Write the chemical equation with states of matter for the reaction between silver nitrate and sodium sulfide.
31. Using your solubility rules, identify whether each of the following compounds are soluble or insoluble in water:
a. Na2CO3
b. Pb(NO3)2
c. Mg3(PO4)2
d. AgCl
32. How does each of the following affect the rate of reaction?
a. concentration?
b. temperature?
c. catalyst?
Unit Five: Chemical Quantities
33. Convert to Moles.
a. 25 g of Al2(CO3)3
b. 189 mL CO2 at STP
c. 6.02 x 1022 atoms of Li
34. Convert to Grams.
a. 10. moles of NaOH
b. 1.2 L of CO2
c. 3.00 x 1024 molecules of O2
35. How many moles of hydroxide ions are in 3.5 mol of Ca(OH)2?
36. Determine the % of Na in NaOH. Show work
37. Determine the empirical formula for a compound composed of 70.0% Fe and 30.0% O.
38. The empirical formula of a compound is CH, its molecular mass is 78. What is its molecular formula?
39. How many grams of water should be produced when 5.00 moles of oxygen react with excess hydrogen? Calculate
the percent yield if 150 grams of water is actually produced. Be sure to start with the balanced chemical reactions.
40. Calcium hydroxide, used to neutralize acid spills, reacts with hydrochloric acid according to the following equation:
Ca(OH)2 (aq) + 2HCl (aq)  CaCl2 (aq) + 2H2O (l)
If you have spilled 6.3 mol of HCl and put 2.8 mol of Ca(OH)2 on it, which substance is the limiting reactant?
41. 20. 0 g of NaOH is dissolved in enough water to make 2.00 liters of solution. Calculate the molarity of the solution.
42. How many grams must be weighed out to make 100.ml of 0.40M NaOH?
43. What volume of 12.0 M HCl is needed to make 100 mL of a 3M HCl solution?
Unit Six: Energy and States of Matter
44. Explain what phase or phase change is occurring at each letter in the graph to the
left.
A _______________________________________
B________________________________________
C________________________________________
D _______________________________________
E ________________________________________
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Is this a heating or cooling graph?
45. Define:
a) Sublimation
Why is there a plateau at B? (What is occurring on the molecular level?)
b. Vaporization
c. Condensation
46. How many Joules of energy are needed to heat 5.0 grams of H2O from 0C to
15C?
(specific heat of H2O = 4.184 J/gC)
47. Answer the following questions based on the phase diagram to the right.
a. What is the state of matter of the substance at 20 mmHg and 25oC?
A
b. Explain what is occurring at point A
48. Identify the measurements needed to calculate heat transfer using calorimetry and
list some possible sources of error.
49. What are the two driving forces of a chemical reaction?
50. Use the following diagram to answer the questions below:
a. Is this potential energy curve endothermic or
25
exothermic?
b. Label each of the following with the numerical
value from the diagram:
___ pot. energy of the products
15
___ H, heat of reaction (forward)
___activation energy (forward)
___H of reverse reaction
___pot. energy of activated complex (forward)
c.
Do
the products have more or less potential energy
0
than the reactants?
d. Which of the following is the correct
thermochemical equation for a reaction with the
kJ
potential energy curve on the left?
A2(s) + B2(s) +heat  2 AB (l)
A2(s) + B2(s)  2 AB (l) + heat
e. Is this reaction increasing or decreasing in entropy?
f. Is the reaction spontaneous always, sometimes, or
never? How do you know?
Unit Seven: Gases (Ch. 12)
51. What is STP? Equal moles of H2(g) and CO2(g) are in a balloon at STP. Which gas molecules are moving faster
and why?
52. A popcorn kernel contains a hard outer shell, with a water and starch mixture on the inside.
Explain, using kinetic molecular theory, why heating the popcorn kernels causes them to pop. In your answer be sure to
explain changes to pressure, volume, temperature as the kernel is popped.
53. A gas occupies 11.2 L at a pressure of 655 kPa and a temperature of 227°C. What will its volume be at STP?
54. If 30 ml of a gas exerts a pressure of 2 atm, what will be the pressure if the volume of the gas is reduced to 20ml?
55. An engineer pumps 5.00 mol of carbon monoxide gas into a cylinder that has a capacity of 20.0 L. What is the
pressure in atmospheres of CO inside the cylinder at 25°C?
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Unit Eight: Equilibrium & Acids, Bases (Ch 19.1, 19.2 , Ch 20 (Sections 1-4), Section 21.1)
56. Consider the equilibrium:
N2O4(g)  2NO2(g)
H = +13.9 Kcal.
a. What will happen if you increase the temperature?
b. The pressure?
57. Write the equilibrium constant expression for:
2NO(g) + O2(g) 2NO2(g)
58. Define:
a. Arrhenius acid
c. Bronsted-Lowry acid
b. Arrhenius base
d. Bronsted-Lowry base
59. A pH of 7 indicates a(n) a. neutral solution b. acid c. base
60. What is the pH, pOH, [H+} and [OH-] of the following solutions?
a. [H+] = 1.3 x 10–2 M
d. pH = 6.5
b. [OH-] = 2.4 x 10–8 M
e. pOH = 4.0
c. [H+] = 0.001M
61. Circle all of the following which are acids:
a. H2SO4 b. NaOH c. NaCl d. HCl e. HNO3 f. RbOH g. HCOOH
62. In a titration, 20.0ml of 0.60 M KOH are required to neutralize 5.00 ml of HCl. What is the concentration of the
acid?
63. What makes an acid strong or weak?
Unit 9: Nuclear Chemistry (Ch 28)
64. Complete the following nuclear questions:
a.
b.
234
90
234
92
Th 
U 
0
e
+
He
+
1
4
2
Type of decay ____________________
Type of decay ____________________
c. Carbon-14 decays by beta emission. Write the nuclear equation.
65. Carbon-14 has a half life of 5730 years. How long has an animal been dead if it started with 200g of Carbon-14
and when it was measured, it had 25g?
66. Define fission and fusion. Give one example of where each reaction occurs.
Lab Equipment and Techniques
67. Match the following pieces of equipment to their use:
___ Used in titration to add exact amounts of
a) crucible
liquid
flask
b) volumetric flask
g) Erlenmeyer
h) balance
___ Used to make a solution of exact Molar
concentration
c) burette
i) evaporating dish
d) graduated cylinder
j) barometer
___ Used to measure atmospheric pressure
e) pipette
68. How could you test for the presence of oxygen, hydrogen and carbon dioxide? What would the results of the
test(s) be?
69. Record the measurements, in cm, indicated by arrows A, B,C, & D.
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B
C
A
0
0.01
0.02
D
0.03
A ________
B ________
C ________
D ______
Level One Chemistry Final Exam Formula Sheet
Constants
Avogadro’s number (NA) = 6.02 x 1023
Gas Constant (R) = 0.0821 (Latm)/(molK) OR
8.314 (LkPa)/(molK)
Conversions
1 atm = 760 mmHg = 101.3 kPa
K = oC + 273
Equations
PV=nRT
P1V1T2 = P2V2T1
P1V1=P2V2
T1
Kw = [H3O+1][OH-1]
M = moles/liters
pH= - log[H3O+1]
T2
[H3O+1] = 1 x 10-pH
q= mCT
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