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CHEMISTRY I – FINAL EXAM STUDY GUIDE-SPRING 2011

Topic 1: Ionic Bonding

Objectives:

1. Predict the compound formed by combining two ions

2. Name ionically bonded compounds

Topic 2: Covalent Bonding

Objectives:

1. Name covalently bonded compounds

Topic 3: Balancing Equations

Objectives:

1. Balance Simple Chemical Equations

Topic 4: Molar conversions.

Objectives:

1. Convert (a) moles to mass (b) moles to atoms (c) atoms to mass and vice versa

2. Calculate molar mass

Topic 5: Chemical Reactions

Objectives:

1. Classify chemical reactions as double replacement, single replacement, combustion, synthesis ordecomposition

2. Apply LeChatlier’s Principle to a given scenario.

Topic 6: Experimental Design

Objectives:

1. Identify independent and dependent variable given an experimental scenario.

2. Select appropriate lab measurement tools and identify correct units

3. Write a number in scientific notation.

4. Setting up conversion factors

Topic 7: Empirical and Molecular Formula

Objectives:

1. Determine a compound’s empirical formula

2. Determine a compounds molecular formula

Topic 8: Stoichiometry.

Objectives:

1. Perform the following conversions (a) Moles to moles (b) Moles to mass (c) Mass to moles

2. Definition of limiting reactant and excess

Preliminary Final Exam Study Guide

Topic 9: Solutions

Objectives:

1. Explain how various factors affect solubility for solids (temperature, stirring, surface

area) and gases (temperature)

2. Calculate molarity

3. Explain definitions of saturated, unsaturated and supersaturated. Identify locations on solubility curve.

4. Define colligative properties and give examples.

5. Interpret solubility curves (How much solute will dissolve at a given temperature and is the solution saturated, unsaturated or supersaturated)

6. Describe the properties of strong and weak acids (extent of dissociation, conductivity, etc)

Topic 10: Acids and Bases

Objectives

1. Define and identify Bronsted-Lowery acid/base, Lewis acid/base or Arrhenius acid/base

2. Describe properties of acids and bases (response to litmus, taste, feel, pH, etc)

Topic 11: Thermochemistry and Thermodynamics

Objectives

1. Interpret an energy diagram (activation energy, exothermic or endothermic)

2. Calculate heat transfer, specific heat or temperature change using q = mc t

3. Explain whether phase changes are endothermic or exothermic (melting. Freezing, boiling, evaporating, condensing, etc) and identify each phase/phase change on a heating curve.

4. Describe why the temperature remains constant during phase changes

5. Describe the states of matter (solid, liquid, gas) in terms of volume, shape and entropy

Topic 12: Kinetics

Objectives:

1. Describe five factors that impact reaction rate

2. Define catalyst and explain its role in a reaction

3. Define activation energy and explain its role in a reaction

Topic 13: Gas Laws

Objectives:

1. Define absolute zero

2. Describe the Pressure/Volume, Volume/Temperature, and Pressure/Temperature relationships in a gas

3. Perform calculations using the combined gas law and ideal gas law (don’t forget to convert all temperatures

to Kelvin)

CHEMISTRY I – FINAL EXAM PRACTICE PROBLEMS – SPRING 2011

STRATEGY: Start by reading through your notes to refresh your memory on these topics. Then, use this review sheet as a starting point to identify the areas on which you need to spend more study time. For those areas, go back to homework assignments, quizzes, and reviews to practice more problems. Keep in mind that these questions are only samples and do not include specific examples of how vocabulary and other conceptual information might appear in a scantron format.

FORMAT:

♦ Questions will include multiple-choice and matching.

♦ A formula bank will be provided in addition to any values that you might need (solubility table, pressure conversions, etc.), but you will NOT begiven “formulas” for items listed in the VOCAB sections (molarity, % composition, etc).

First Semester Topics

1. Give the longhand electron configuration for arsenic.

2. The largest atoms are in the ___ corner of the table.

3. Classify the following as chemical or physical changes (3-5).

A. rusting of iron

B. digestion of meat

C. boiling water

4. State the law of conservation of matter.

5. Write formulas for the compounds in 7-10.

A. magnesium fluoride

B. dinitrogen pentoxide

C. sodium sulfate

D. phosphoric acid

6. Name the compounds in 11-14.

A. KNO3

B. HBr

C. SO3

D. FeCl3

7. Draw the Lewis diagram & specify the molecular polarity (15-16).

A. AsH3

B. BF3

Chemical Reactions – Ch. 9

8. Write a word equation for the following reaction (incl. how many? of what? what state?).

Ba(ClO3)2(s) → BaCl2(s) + 3O2(g)

9. Rewrite and balance the following word equation using chemical formulas, physical states, and energy. – When solid sodium chlorate absorbs energy, it

10. produces solid sodium chloride and oxygen gas.

11. Predict the products and balance (19-22). Write N.R. if no reaction will occur.

Cu(s) + MgSO4(aq) →

C5H12(l) + O2(g) →

NH4Cl(aq) + Pb(NO3)2(aq) →

Fe2O3(s) →

12. For each of the reactions in #11, specify whether it is combustion, synthesis, decomposition, single replacement, ordouble replacement.

13. Identify as endothermic or exothermic.

A. PE of products is lower than PE of reactants.

B. PE of products is higher than PE of reactants.

C. When substances are mixed, the test tube feels cold.

D. In your car’s engine, fuel is burned to produce energy.

E. List three conditions required for a successful collision according to Kinetic Molecular Theory.

Name four ways to increase the rate of a reaction.

VOCABULARY

• endothermic

• exothermic

• catalyst

The Mole – Ch. 10

30. How many magnesium sulfate molecules are in 25.0 g?

31. Find the molarity of a 750 mL solution containing 346 g of

potassium nitrate.

32. Calculate the number of grams required to make a 50.0 mL

solution of 6.0M NaOH.

33. Find the % composition of copper (II) chloride.

34. The percent composition of a compound is 40.0% C, 6.7% H,

and 53.7% O. The molecular mass of the compound is 180.0

g/mol. Find its empirical and molecular formulas.

35. Combustion analysis gives the following: 26.7% C, 2.2% hydrogen,

71.1% oxygen. If the molecular mass of the compound is 90 g/mol,

determine its molecular formula.

36. A substance is decomposed and found to consist of 53.2% C, 11.2%

H, and 35.6% O by mass. Calculate the molecular formula of the

unknown if its molar mass is 90 g/mol.

VOCABULARY

• Avogadro’s number

• empirical formula

• percent composition

• molecular formula

Stoichiometry – Ch. 11

37. How many grams of copper would be produced from 49.48 g of chromium?

Cr +CuSO4 → Cu + Cr2(SO4)3

38. How many grams of chromium are required to react with 125 mL of 0.75M CuSO4.

Cr +CuSO4 → Cu + Cr2(SO4)3

39. How many grams of ZnS are required to react with 12.6 L of oxygen gas at STP?

ZnS + O2 → ZnO + SO2

40. 6.45 g of lithium reacts with 9.20 g of oxygen gas to produce lithium oxide. How

many grams of Li2O are formed?

41. What are the limiting and excess reactants in #40?

42. The actual yield of the reaction in #40 is 12.5 g. What is the percent yield of this

reaction?

VOCABULARY

• Theoretical yield

• limiting reactant

• percent yield

• excess reactant

Gases – Ch. 13

Identify the gas laws that explain these situations (43-45). Specify

the variables involved and direct/inverse relationship.

43. A balloon pops after floating high into the atmosphere.

44. A balloon pops in a hot car on a summer day.

45. Do not store aerosol cans at temperatures above 120°F.

Danger of explosion.

Identify the gas law and solve the problem (46-47).

46. Hydrogen gas is collected over water at 35°C to give a total

pressure of 0.80 atm. Find the pressure of the dry hydrogen

gas in kPa. (see p.899 for necessary data)

47. A jar is tightly sealed at 22°C and 772 torr. What is the pressure

inside the jar after it has been heated to 178°C?

Gases – Ch. 13 (continued)

48. 300.0 mL of gas has a pressure 75.0 kPa. When the volume is

decreased to 125.0 mL, what is its pressure?

49. Hydrogen diffuses 3.72 times faster than an unknown gas. Find

the molar mass of the unknown gas.

50. 50.0 L of gas has a temperature of 75°C. What is the temp in

Celsius when the volume changes to 110 L?

51. What is the volume of a container that holds 48.0 g of helium at

a pressure of 4.0 atm and temperature of 52°C?

52. Neon diffuses at a rate of 688 m/s. What is the speed of

ammonia (NH3) at the same temperature and pressure?

53. A gas occupies 325 L at 25°C and 98.0 kPa. What is its volume at

70.0 kPa and 15°C?

54. What volume of SO2 is produced from 32.5 g of ZnS at 23°C and 103.3

kPa? ZnS + O2 → ZnO + SO2

55. Define real gases. When do they act like ideal gases?

56. Explain Graham’s law. How does molar mass affect the rate of

diffusion?

VOCABULARY

• Kelvin diffusion

• STP effusion

•

Solutions – Ch. 14

57. Explain the effect of adding more solute to unsaturated,

saturated, and supersaturated solutions.

58. Explain how temperature and pressure affect solubility.

State whether each pair is soluble or insoluble (75-78).

59. KCl in water

60. ammonia in oil

61. wax in C6H6

62. CH4 in water

63. Read solubility curves (See Nature of Solutions w/s and

Solutions Quiz).

64. How many grams of AlCl3 are required to make a 2.25 m

solution in 30.0 g of water?

65. What volume of 12M HCl is needed to prepare 250 mL of

0.20M HCl?

66. Explain the difference in preparing solutions based on molarity

versus molality.

67. Which will have the greatest effect on ΔTf at the same molality:

C12H22O11, MgBr2, AlCl3, or NH4NO3?

68. When 26.4 g of NaBr dissolves in 0.20 kg of water, what is the

freezing point of the solution? (see p.438)

VOCABULARY

• Solvation

• solubility

• dissociation

• ionization

• molality

• strong/weak/nonelectrolyte

Acids and Bases – Ch. 18

State whether the following are acids or bases (69-72).

69. Have a sour taste.

70. React with metals.

71. Feel slippery

72. Turn blue litmus paper red.

73. Define acids and bases according to Arrhenius, Brønsted-Lowry,

and Lewis.

74. Identify each substance as acid, base, conjugate acid, or

conjugate base. H2S + H2O → HS – + H3O+

75. Give the conjugate acids of: NH3 and Br –.

76. Give the conjugate bases of: H3O+ and HSO4–.

77. Find the pH of 0.75M HCl.

78. Find the molarity of a KOH solution with a pH of 9.5.

79. Is the solution in #94 acidic or basic?

80. When a neutralization reaction between a strong acid and a weak

base reaches the equivalence point, will the solution be acidic,

basic, or neutral?

81. If 43.5 mL of 0.15 M HBr is required to neutralize 25.0 mL of

Ca(OH)2, what is the molarity of Ca(OH)2?

VOCABULARY

• hydronium ion

• neutralization reaction

• amphoteric substance

• strong/weak acid/base

Reaction Rates – Ch. 16

82. What does the symbol in the following chemical equation

mean?

83. How does the nature of reactants affect the reaction rate?

84. How does concentration affect reaction rate?

85. How does surface area affect reaction rate?

86. How does temperature affect reaction rate?

87. How do catalysts and inhibitors affect reaction rate?

88. What is activation energy?

89. What is the difference between a catalyst and an inhibitor?

90. What is an example of a heterogeneous catalyst?

91. Tell me as much as you can about the graph below

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