Chemistry SPRING SEMESTER REVIEW Name:________________________

RefreshersName the following compounds:N2O-Na2O-MgO-(NH4)2CO3-Ni(OH)2-HCl –H2SO4-

Write the formulas for the following compounds:

dinitrogen tetroxidecopper (I) sulfidecalcium nitratepotassium phosphatesulfuric acid

1. Write a balanced chemical equation for each word problem. Be sure to include all appropriate symbols, formulas, and balancing!a. Solid sodium metal reacts with liquid water to form hydrogen gas and solid sodium hydroxide.

Chemical Reactions and Conservation of Mass2. Solve for the missing mass in the following reactions

A) 3000g 2H2O2 1100g H2O + ____g O2

B) 2Na + I2 2NaI 37g + 52g _____g

C) According to the law of conservation of mass, how much zinc was present in the zinc carbonate?

3. Label each of the following reactions as Oxidation-Reduction (Red-Ox), Acid-Base, Precipitate or Nuclear reaction?

a. 2Ca(s) + O2(g) ⟶ 2CaO(s)b. Ba(NO3)2(aq) + Na2SO4(aq) ⟶ BaSO4(s) + 2NaNO3(aq)c. HBrO3(aq) + KOH(aq) ⟶ KBrO3(aq) + H20(l)d. NaBr(aq) + AgNO3(aq) ⟶ NaNO3(aq) + AgBr(s)e. CH4 + 2O2 2 H2O + CO2

f. U92235 Th90

231 + He24

4. Ethane (C2H6) is a colorless gas used as a fuel. Ethane burns in oxygen to produce carbon dioxide and water. The balanced chemical equation that describes this reaction is --

a. 2C2H6 + 7O2 ⟶ 4CO2 + 6H2Ob. 2C2H6 + 5O2 ⟶ 3CO2 + 6H2Oc. C2H6 + 7O2 ⟶ 4CO2 + 3H2O

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d. 3C2H6 + 7O2 ⟶ 6CO2 + 9H2O

5. Using the Activity series, determine if the reactions below will occur. If the reaction will not occur, write no reaction. If the reaction will occur, predict the products of the reaction and write a balanced chemical reaction.

a. Ag + Cu(NO3)2

b. MgCl2 + Ba

c. Zn + HCl

d. H2O + Sr

e. KI + Br2

6. Use the Solubility charts and Predict if the following compounds are soluble or insoluble. Write soluble (aq) or insoluble solid(s)

a. KCl ________________( ) e. NH4OH________________( )

b. K2S ________________( ) f. Ba(OH) 2 ________________( )

c. CaS ________________( ) g. PbS ________________( )

d. BaSO4 ________________( ) h. FeBr3________________( )

7. In the reaction shown below, what is the name of the precipitate? (Hint: figure out if your products form solids(s) or aqueous solutions (aq).. ..the solid is your precipitate)

BaCl 2(aq) + Na 2 SO 4(aq) → 2NaCl ( ___) + BaSO 4(___)

a. Barium Sulfateb. Sodium Chloride

8. Determine if the following pairs of solutions will form a precipitate. If they do form a precipitate, write the balanced reaction. If they do not form a precipitate, write NVR for no visible reaction.

a. AgNO3(aq) + CaCl2(aq)

b. Pb(NO3)2(aq) + KC2H3O2(aq) _____

c. CuSO4(aq) + KOH(aq)

9. Balance the following combustion reactions.

a. C4H8 + O2 CO2 + H2O

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b. C8H18 + O2 CO2 + H2O

10.Based on the oxidation numbers (charge) known for the other element(s), determine the oxidation number of the underlined element.

a. CaO _____ b. KCl _____ c. H2O _____d. Ag _____ e. AlCl3 _______

11. Define oxidation:12. Define reduction:13. In an oxidation-reduction reaction, reduction is defined as the —

A.

gain of protons.

B.

gain of electrons.

C.

loss of electrons.

D.

loss of protons.

14. Determine the oxidation numbers of each element in the following reaction and then complete the chart to determine which is being oxidized and which is being reduced. (label oxidation numbers for all elements) Fe + 2HCl H2 + FeCl2

A

.The manganese is oxidized, and its oxidation number changes from +2 to +4.

B. The manganese is reduced, and its oxidation number changes from +4 to +2.C. The manganese is oxidized, and its oxidation number changes from +4 to +2.D.

The manganese is reduced, and its oxidation number changes from +2 to +4.

Stoichiometry Mole Math

16. Complete the mole road map. You will use this for the Mole Math and Stoichiometry sections.

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15. T

17. Find the mass of 0.89 mol of CaCl2.

18. Determine the number of moles of hydrogen that are in 362.8g of this diatomic element.

19. Determine the number of atoms that are in 0.58 mol of Se.

Stoichiometry (grams/liters/particles moles moles grams/liters/particles)

20. Define Stoichiometry:

21. State the law of conservation of mass:

22. Determine the molar mass of C6H12O6.

23. How many moles of sodium are needed to react with 6.58 mol of oxygen?___Na + ___O2 ___Na2O

24. How many atoms of Cr are needed to produce 27.0 grams of Cu?___Cr + ___CuSO4 ___Cr2(SO4)3 + ___Cu

25. How many grams of carbon dioxide would be produced when 89.2 L of oxygen reacts with C3H6?2C3H6 + 9O2 6CO2 + 6H2O

26. Stoichiometry calculations involving volumes of gases use the conversion 22.4 L = 1 mole. This is only valid under standard temperature and pressure conditions.

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Standard Temperature = °C = KStandard Pressure =____________atm

Limiting Reagents27. Define Limiting Reagent:

28. 4.7 grams of CO react with 0.54 grams of H2 to produce methanol, CH3OH.

CO + H2 CH3OH

a. Determine how many grams of methanol, CH3OH, can be produced by each of the reactants using stoichiometry.

b. Which reactant is the limiting reagent?

29. A 2.00 g sample of ammonia is mixed with 4.00 g of oxygen in the balanced reaction below.  How much nitrogen monoxide can be produced, in grams?

4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g)

Percent Yield30. Define Percent Yield:

31. How can you determine the theoretical yield?

32. How do you determine the actual yield?

33. The actual amount of product formed in a reaction during lab was 55 grams. You however calculated that 58 grams would form. What is the percent yield of this product?

34. Determine the percent yield if 86.0 g of ammonia is produced when 66.0 g hydrogen reacts with excess nitrogen. ___H2 + ___N2 ___NH3

Which reactant is the limiting reactant? _______________

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Gas Laws

35. In a sample of gas, what causes pressure? Answer in a complete sentence.

36. How does the behavior of gas particles change as temperature is decreased? How does this affect the pressure?

37. When performing gas law calculations, temperature must be in what units? _________________

38. A gas at constant temperature is in a container with a volume of 15 L. The container is contracted to 7.5 L. What will the pressure inside the container be if no gas has been removed?

Use the following graphs to answer the next 7 problems.

39.

Which graph represents Boyle’s Law?

40. Which graph represents Charles’s Law?

41. Which graph has the same shape as Avogadro’s Law?

42. In Graph #1, what is the pressure on 40mL of gas?

43. In Graph #2, what is the volume when the temperature reaches 30K?

44. Which graph represents an inverse relationship?

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5 10 15 20 25 30 35 40 45 50 550

50

100

Graph #2: The Relationship between Volume and Temperature of a Gas

Temperature (K)

Volu

me

(L)

0 20 40 60 80 1000

20

40

60

80

100

The Relationship between Volume and Pressure of a

Gas

Volume (mL)

Pres

sure

(kPa

)

45. Which graph represents a direct relationship? Show work (with units) for the following nine problems.46. What is the volume of a sample of gas that contains 1.2mol at 780 K and 1.3 atm?

47. A balloon inflated in a room at 24°C has a volume of 4.00L. The balloon is then heated to a temperature of 58°C. What is the new volume if the pressure remains constant?

48. Laughing gas is a mixture of dinitrogen monoxide and oxygen. If the total pressure is 1.2atm, what is the partial pressure of dinitrogen monoxide given that the partial pressure of oxygen is 0.45atm?

49. If the volume of a tire increases from 20.0L to 25.0L, what was the initial pressure if it is now at 1205mmHg?

50. What will the new volume of nitrogen gas be if your original sample is 10.0L, 32.0C, and 210kPa; and the temperature and pressure are increased to 45.0C and 355kPa, respectively?

51. The volume of oxygen in your lungs increased from 250.0mL to 375.0mL. If the initial temperature was 25C, what is the temperature now, in Celsius?

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52. What volume does 0.742mol of argon gas occupy at STP?

53. A 6.0 L sample at 25 °C and 2.00 atm of pressure contains 0.5 moles of a gas. If an additional 0.25 moles of gas at the same pressure and temperature are added, what is the final total volume of the gas?

54. If 4.5g of oxygen gas is placed in a 450mL container at 450K, calculate the pressure of this container.

55. At high altitudes, pilots have to supplement their supply of oxygen. In this mixture, there are oxygen and nitrogen gases. If nitrogen’s partial pressure is 550mmHg, calculate the partial pressure of oxygen if the total pressure is 945mmHg.

Solutions Review

56. Which characteristic of water explains its ability to dissolve a great variety of things?

a. its transparency in lightb. its electrical conductivityc. its physical state of matterd. its molecular arrangement and shape

57. Does Ce2(SO4)3 resemble a solid or gas__________________

58. Explain your answer to #56.

59. What is the solubility of Pb(NO3)2 at 20°C? __________

60. If you have 30g of KClO3 at 40°C, the solution is:a. unsaturated

b. saturatedc. supersaturatedd. slightly saturated

61. How many grams of K2Cr2O7 are needed to saturate 200g of H2O at 50°C? _______62. Increasing the pressure of 35g solid NaCl at 20°C would _____________the solubility.

increase decrease not affect

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63. You need to make a solution containing 150g of potassium chloride in 300g of water. What temperature is required? (Use solubility graph on the previous page!)

64. Circle the condition that causes salt to dissolve faster in water.rock salt or in granulated form?when allowed to stand or when stirred?at a higher temperature or a lower temperature?

65. Where would a data point be on a line graph if a solution is? unsaturated:___________________, saturated:_______________, or supersaturated:______________________

66. If two liquids dissolve each other, they are said to be ___________________.

67. How does the solubility of solid substances change as the temperature of the solvent increases?a. The solubility increases for most solids.b. The solubility decreases for most solids.c. The solubility remains constant.

68. How does the solubility of a gas change with the increase in temperature?

69. What types of compounds when dissolved in water will cause the solution to be electrolytic?

70. Label the partial charges and the element symbols on the water molecule below.

71. Which of the following solutions are considered electrolytes and would cause a light bulb to give a bright light :

a. sugar solutionb. NO2 solutionc. CaCl2 solutiond. Ag

72. Match the following terms with their correct definitiona) Hydrogen bond b) Polar molecule c) Nonpolar moleculed) Electrolytee) Nonelectrolyte

_____ Molecule in which there is a partial negative charge on one end and a partial positive charge on the

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other end _____ A substance that does not dissolve to form ions in an aqueous solution_____ Bond occur between a hydrogen and the negatively charged end of another molecule_____ A substance which forms ions in an aqueous solution_____ Molecule which has no separation of charge, so no positive or negative poles are formed.

73. Describe the settling of particles for each of the following:

Solutions-

Suspensions-

Colloids-

74. How would you make a solution more concentrated?

How would you make a solution more dilute?

75. What does water, a polar covalent molecue, dissolve? (Circle one or more.)a. Ionic compoundsb. Polar compoundsc. Nonpolar compounds

76. How many liters are there in a 5.2 M solution that contains 0.5 moles of solute?

77. What is the Molarity of 0.107 moles of NaOH that are dissolved in 2.0 L of water?

78. How many moles of solute are there in 1.5 L of a 3.25 M NaOH solution? How many grams of NaOH would be in this solution?

79. If enough water is added to a 138mL solution of 5.0M HCl to increase the volume to 238mL, what is the concentration of the new, dilute solution?

80. What volume of 16.0M sulfuric acid must be used in order to prepare l.5L of 0.10M H2S04 solution?

81. What are the three factors that influence the rate of dissolving? For each one, list how to increase the dissolution using the factor.

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ACIDS/BASES82. Define Arrhenius acid:83. Define Bronsted-Lowry Acid:

84. What is an amphoteric substance?

85. Calculate the pH of the following concentrations and determine if it is an acid, base or neutral solution.

a. [H+] = 1.0 x 10-8 ___________________ ______b. [H+] = 1.0 x 10-3 ___________________ ______c. [H+] = 2.6 x 10 -4 ___________________ ______ d. [H+] = 9.1 x 10 -13 ___________________ ______

86. As hydrogen ion concentration of an aqueous solution increases, the hydroxide ion concentration of this solution

a. increasesb. decreasesc. remains constant

87. Which is not a property of an acid?a. conducts electricityb. reacts with metalsc. corrosived. taste bitter

88. Which represents the most acidic substance?a. Apple Juice (pH = 3.0)b. Black Coffee (pH = 5.0)c. Hand Soap (pH = 10.0)d. Household Ammonia (pH = 12.0)

89. Which substance will produce the highest pH?a. HCl(aq)b. H2O(l)c. NaCl(aq)d. NaOH(aq)

90. Which pH indicates a solution that has the highest concentration of OH- ions?a. 2b. 4c. 6d. 8

91. Which pH indicates a solution that has the highest concentration of H+ ions?a. 2b. 4c. 6d. 8

92. Using the Kw constant and equation, solve for the following:a. Calculate the [OH-] of a solution with an [H+] of 1.6X10-6.

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b. Calculate the pH of a solution with an [OH-] concentration of 1.85 X 10-10. (Hint: 2 steps)

93. What is the Molarity of 3.50 L of solution that contains 60.0g of sodium chloride?

94. Write a complete and balanced equation for the following acid-base reaction:

a. H3PO4 + Ca(OH)2

b. HCl + Mg(OH)2

95. Calculate the concentration of sulfuric acid if 65.0mL of this acid is neutralized by 23.0mL of .15M potassium hydroxide? ___ H2SO4 + ____ KOH ____ K2SO4 + ____H2O

Use:MaVa = MbVb Na nb

COVALENT BONDING

96. Define ionic and covalent bonding.

97. Define the octet rule. What are the exceptions?

98. How many electrons are in a single bond _______ double bond ______ triple bond _______?

99. Draw the Lewis structure (dot diagram) for the element nitrogen. How many valence electrons?

100. Determine if the following bonds are ionic, polar, or nonpolar using the electronegativity values in the table below.

A. N - N

B. N – O

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Element

Electronegativity

H 2.1Br 2.8C 2.5N 3.0O 3.5F 4.0Cl 3.0I 2.5

Na 0.9P 2.1

C. C – Cl

D. O – I

E. Na - Br

F. P - F

For the polar bonds, circle which bond was the most polar? Also, assign the positive pole(ς+) and negative pole(ς -) for the bond.

101. Draw the Lewis structure for the following molecules. a. Label the type of bonds (polar or nonpolar covalent). b. What is the shape? c. Is the molecule polar or nonpolar?

a. CO2 b. H2O c. NH3 d. CCl4 e. AlCl3

a. b. c. d. e.

b. b. b. b. b.

c. c. c. c. c.

102. According to VSEPR theory, why do molecules adjust their shapes?

103. Which two molecular shapes are polar no matter what elements are involved? Explain why.

104. In the Lewis structure for CCl4, how many unpaired electrons would appear on the structure?

THERMODYNAMICS

105. Complete the following chart:

Type of Reaction Sign of ∆Hrxn Which has more chemical (potential) energy: reactants

or products?

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Exothermic  Endothermic  

106. Make enthalpy diagrams for the two chemical reactions from above, showing clearly the amount of energy released or gained. Label the reactants (R), products (P), activation energy (Ea), and ∆Hrxn for each diagram.

Exothermic Change       Endothermic Change      

 

107. Classify each of the following as an exothermic or endothermic process. Melting ice cubes _____________________ Baking Bread _____________________ Burning a candle ____________________ Splitting a gas molecule apart _____________Evaporation of water____________________ Formation of snow in clouds ______________

108. Using the enthalpy of formation chart, (a) find the ΔH for the following reactions (b) classify as endothermic or exothermic, and (c) rewrite the equation with heat included on the correct side of the equation.

Compound ΔHf° (kJ/mol)NO (g) +90.4NO2 (g) +33.9NaOH (s) -426.7HCl (g) -92.3NaCl (s) -411.0H2O (l) -285.8H2O (g) -241.8

I) 2 NO(g) + O2(g) ---> 2 NO2(g) a.ΔH =

b. endothermic or exothermic

c. Rewrite the Thermochemical equation:

II) NaOH(s) + HCl(g) ----> NaCl(s) + H2O(g) a. ΔH =

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P.E.

Time Time

P.E.

b. endothermic or exothermic

c. Rewrite the Thermochemical equation:

109. Calculate the number of joules needed to warm 100 grams of water from 25.00C to 80.00C.

110. A sample of zinc, specific heat of 0.386 J/g0C, released 1,964 joules of heat when it cooled from 92.50C to 65.00C. What was the mass of the zinc sample?

111. Students conduct an experiment where a reaction occurs in a calorimeter. Calculate the heat released/absorbed in Joules to the nearest whole number. The specific heat capacity of water is 4.184 J/g⁰C. Is the reaction endothermic or exothermic?

Mass(g)

Initial Temperature of Water (ºC)

Final Temperature of Water (ºC)

100.0 25.0 10.8

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