Page 1

First Semester Final Exam Study Guide Name: ______________________________

Chemistry 1 2 3 4 5 6 7 60 multiple choice questions Chapter 1

1. Define matter and list five examples. Chapter 2

2. Define pure substance.

3. Define element.

4. Define compound.

5. Define chemical property and give 3 examples.

6. Define physical property and give 3 examples.

7. Define chemical change and give 3 examples.

8. Define physical change and give 3 examples.

Page 2

Chapter 3

9. Write the following in scientific notation:

a) 153,000 kg b) 602200000000000000000000 atoms

c) 0.256 m

d) 0.000000000044 mL

10. Write the following in expanded form:

a) 2.34 x 108 cm b) 1.56 x 1015 m

c) 9.67 x 10-13 g

d) 4.87 x 10-6 cm

11. Perform the following metric conversions.

a) 300 g to mg b) 4000 g to kg

c) 60 kg to g

d) 45 km to mm

Chapter 4

12. Explain the law of conservation of mass.

13. What mass of product is produced in the following reaction? 2 grams H2 + 3 grams Cl2

Page 3

14. Complete the chart.

Subatomic Particle Charge Size Location in Atom

15. Which subatomic particle makes up most of an atom’s mass?

16. Most of an atoms volume is really what?

17. Define isotope and give an example.

18. Why is the atomic mass of a given element usually NOT listed on the periodic table as a whole number?

19. What particle does the atomic number listed on the periodic table represent, and where does this particle reside?

20. Write the equation that relates the mass # to the number of protons and neutrons in the nucleus of an atom.

Page 4

21. Complete the chart.

Name Symbol Protons Neutrons Electrons Atomic Mass Atomic Number

Potassium

Pb 82

Arsenic 33

47 63

260 101

22. Which elements from problem 36 were isotopes?

Chapter 5

23. Fill in the orbital diagram below for Tungsten and then write the full electron configuration below.

W:

Page 5

24. Which of the following is an incorrect orbital diagram? Why?

25. Write the unabbreviated electron configuration for the following atoms:

a) Helium

b) Oxygen

c) Aluminum

d) Neon 26. Write the abbreviated electron configuration for the following atoms:

a) Silver

b) Copper

c) Tungsten

d) Titanium 27. Define ground state. 28. Define excited state.

Page 6

29. How do atoms jump from ground state to excited state and back?

30. Which atomic particle changes energy states? Chapter 6

31. Define period.

32. Define group. What is another name for group?

33. Define electron affinity.

34. Define electronegativity. How is it used to determine bonds?

35. Define atomic radius.

36. Define ionization energy.

37. Fill in the following chart with the properties of metals, nonmetals and metalloids.

Metal Nonmetal Metalloid

Conductivity

React w/ acid

State of matter

Malleable

Brittle

Page 7

38. Label the following on the attached periodic table (at the end of SG):

1) Atomic radius trend

2) Electronegativity trend

3) Ionization energy trend

4) Alkali metal

5) Alkaline earth metals

6) Halogens

7) Noble gases

8) Actinide

9) Lanthanide

10) Metals

11) Nonmetals

12) Metalloids

Chapter 7 & 8

39. How is an ion different from an atom? 40. Define cation and give an example.

41. Define anion and give an example.

42. Compare and contrast ionic and covalent bonds.

43. How many electrons are shared in a covalent bond?

44. Define octet rule. Which family on the period table has a complete octet and thus does not react to form compounds?

45. Define valence electron.

46. How many valence electrons do the following atoms have:

a) Magnesium b) Bromine

c) Calcium

d) Phosphorous

Page 8

47. The electron configuration of oxygen is 1s2 2s2 2p4. How many more electrons does oxygen need to satisfy the octet rule?

Chapter 9 48. Name the following ionic compounds.

a) Al(SO4)2 b) PbF4

c) KBr2

d) Ca(OH)2

e) Mg(NO2)2

49. Write the formula for the following ionic compounds.

a) Magnesium chlorate

b) Potassium cyanide

c) Tin (IV) chloride

d) Zinc nitrate

e) Calcium chlorite

50. Name the following covalent (in other words molecular) compounds.

a) As3P5

b) IF7

c) NO2

d) SrH2

e) SiCl4

Page 9

51. Write the formula for the following covalent (molecular) compounds.

a) dinitrogen pentoxide b) carbon monoxide

c) trisulfur difluoride

d) triarsenic pentanitride

e) dihydrogen monoxide

Chapter 10

52. Circle the subscripts in the following chemical formula.

K2SO4 53. What information does a subscript reveal and how would changing a subscript in a chemical formula

change the chemical it represents?

54. Define molar mass.

55. Calculate the molar mass of the following:

a) Au =

b) H2O =

18

c) Mg(MnO4)2 =

262 d) CH3CH2COOH =

75

56. Define molecular formula.

57. Define empirical formula.

Page 10

58. What is the empirical formula for C6H12 ?

59. What is the empirical formula for Hg2Cl2 ?

60. What is the percent composition of the following compounds:

a) potassium cyanide

KCN 60% K, 18% C, 22% N b) butane, C4H10

83% C, 17% H c) sulfuric acid, H2SO4

33% S, 2% H, 65% O

61. Find the empirical formula of a compound that is 63.52 % iron and 36.48 % sulfur.

FeS 62. Calculate the empirical formula of a compound containing 1.0 g K, 0.70 g Cr, and 0.82 g of O.

K2CrO2 63. A white powder is analyzed and found to have an empirical formula of P2O5. The compound has a

molar mass of 283.88 g/mol. What is the compound’s molecular formula?

P4O10 64. A compound with the empirical formula CH4O was found to have a molar mass of approximately 192

g/mol. What is the molecular formula of the compound?

C6H24O6

Page 11

65. How many atoms are there in a mole of gold?

6.022 x 1023 66. Convert 3.5 moles of nickel into atoms.

2.11 x 1024 67. How many moles are in 1.20 x 1024 atoms of Zinc?

1.99 mol 68. What is the mass of 5 moles of C6H12O6, glucose?

900 g 69. How many grams are in 3 moles of HCl, hydrochloric acid?

108 g

70. How many moles are in 1 gram sample of gold?

5 x 10-3 moles 71. A scientist has a sample of magnesium that has a mass of 5 grams. How many atoms of magnesium

does the scientist have?

1.25 x 1023 moles Chapter 11 & 12

72. Define product.

73. Define reactant.

Page 12

74. Label the products and reactants in the following chemical reactions:

H2 (g) + O2 (g) H2O (l)

Copper sulfate (aq) + iron (s) iron sulfate (aq) + copper (s)

H2O + CO2 + light C6H12O6 + O2

75. What is changed when balancing a chemical equation?

76. Identify the type of reaction and balance:

a) ___ Al + ___ NaOH → ___ Na3AlO3 + ___ H2 _____________________________

b) ___ C12H22O11 + ___ O2 → ___ CO2 + ___ H2O _____________________________

c) ___ Ca + ___ H2SO4 → ___ CaSO4 + ___ H2 _____________________________

d) ___ Cu + ___ HNO3 → ___ Cu(NO3)2 + ___ NO + ___ H2O _____________________________

77. Predict the products and balance these chemical reactions:

a) ___ Cu + ___ HNO3 →

b) ___ SiF4 + ___ NaOH →

c) ___ MnO2 + ___ K →

d) ___ K + ___ Cl2 →

e) ___ C2H8 + ___ O2 →

Page 13

78. Define limiting reactant.

79. Define excess reactant.

80. Mg(OH)2 + 2 HCl MgCl2 + 2 H2O

If the reaction begins with 25.44g of Mg(OH)2, what mass of H2O is produced from the reaction?

81. 6 Li + Ca3(PO4)2 2 Li3PO4 + 3 Ca

If the reaction above produced 84.3 grams of calcium, what mass of lithium was used to start the

reaction?

82. Ba3(PO4)2 + 6 KCl 3 BaCl2 + 2 K3PO4

The above reaction begins with 3.71 g of Ba3(PO4)2 and 3.71 g of KCl. What is the limiting reactant?

What is the excess reactant?

83. 2 N2H4 + N2O4 3 N2 + 4 H2O

The above reaction begins with 153.89 g of N2H4 and 542.92 g of N2O4, what is the limiting reactant? If

the reaction actually produces 190.4g of N2, what is the percent yield?

Page 14