Chemistry Essentials Final Exam Review Spring 2016 Answers

Part I Completion

Fill in the blanks with the word or phrase that best completes the sentence.

1. The _electron__ is the smallest atomic particle and is charged ___negatively______. It is found in the ______electron cloud____. 2. When _electrons___ are added or removed from an atom, a (n) ion_ is formed. 3. An ion that contains more protons than electrons is called a (n) _cation__. 4. The __neutron_ is the largest atomic particle and has _no_ charge. It is found in the _____nucleus_____ 5. An ion that contains more electrons than protons is called a (n) _anion___.The existence of isotopes__ explains why the atomic masses for elements are not express as whole numbers. 6. The proton is slightly smaller than the ___neutron__, and is charged ____positively____. It is found in the ___nucleus_____. 7. __Isotopes______of an element have the _same__ number of protons, but ___different____ numbers of ____neutrons_____. 8. A chemical equation is balanced by adding ____coefficients__ as needed. 9. At room temperature, bromine and mercury are ___liquids____. 10. At room temperature, there are 11 elements that exist as gases. They are H, He, N, O, F, Cl, Ne, Ar, Kr, XE, and Rn______.

11. The diatomic elements are H2, N2, O2, F2, Cl2 Br2, and I2______. 12. When a chemical equation is balanced, there are the _same_ types of atoms and the __same numbers of each type of atom on both sides of the equation. 13. Balancing an chemical equation satisfies the _Law of Conservation of Matter______. 14. When writing a chemical equation in standard form, _g_____ is used to indicate a gas. 15. When writing a chemical equation in standard form, l is used to indicate a liquid. 16. When writing a chemical equation in standard form, s__ is used to indicate a solid. 17. When writing a chemical equation in standard form, _aq__ is used to indicate a solution(dissolved in water.) 18. When an atom gains or loses _electrons____, an ion is formed. 19. An anion has more electrons__ than ___protons____. 20. A cation has more _protons___ than _electrons__. 21. __Ionic____ compounds are held together by electrostatic interactions. 22. Ionic compounds contain a metal and a(n) __nonmetal______. 23. The coefficients in a chemical equation are always the smallest ___whole__ numbers possible. 24. A sample of ruthenium has a mass of 101.1 amu. There are 6.02 x 1023_ ruthenium atoms are in that sample. 25. A sample of lead has a mass of 207.2 g. There are __6.02 x 1023_____ lead atoms are in that sample. 26. 6.02 x 1023 tungsten atoms has a mass(in grams) of _183.84______. 27. In order to calculate the mass of a product in a chemical reaction, you must have _a balanced chemical equation____, _the mass of another reactant or product______, and the molar mass of eh product______. 28. The mole ratio of two components in a chemical reaction is determined from the ___coefficients in the balanced equation______. 29. When using a balanced chemical equation to calculate the mass of a product produced from a known mass of a reactant, you must convert the mass of the reactant into___moles of reactant______.

30. Two correct interpretations of the balanced equation 2Al(s) + 3CuSO4(aq) → Al2(SO4)3(aq) + 3Cu(s)are 2 atoms of aluminum react with 3 molecules of copper(II) sulfate to form 2 molecules of aluminum sulfate and 3 atoms of copper and 2 moles of aluminum react with 3 moles of copper(II) sulfate to form 2 moles of aluminum sulfate and 3 moles of copper ______. 31. A chemical equation is balanced by adding ____coefficents___ as needed. 32. When a chemical equation is balanced, there are the same types of atoms and the same_ numbers of each type of atom on both sides of the equation. 33. Balancing an chemical equation satisfies the Law of Conservation of Matter____ 34. ______. 35. When an atom gains or loses electrons__, an ion is formed. 36. An anion has more _electrons_ than __protons. 37. A cation has more _protons_ than __elecytrons___. 38. The driving forces for a chemical reaction are transfer of electrons, formation of a gas, formatin of a solid, formatin of water 39. The loss of electrons is __oxidation_____. 40. The gain of electrons is __reduction_____. 41. The formula for an acid begins with _H___, and the formula for a base ends with_OH. Both of these types of compounds are _ionic and aqueous____. 42. A (n) indicator_ is a chemical that has two or more colors depending upon what type of ions are around it. 43. When a substance dissolves in water, it is considered _soluble___, and when it does not dissolve in water it is considered _insoluble__. 44. When a calcium atom becomes a calcium ion, it __loses__ 2 electrons. 45. When a bromine atom becomes a bromine ion, it __gains__ 1 electron. 46. The charge on any atom is __0_. 47. The charge on any compound is ___0_____. 48. . A student performed a series of experiments and recorded the information in the table below. Complete the table. Experiments Exp.# Description of Experiment Did a Chemical Reasons(s) Observations During & After Reaction Occur? 1 A clear colorless liquid was placed over a no Phase change flame. Bubbles formed in the liquid. 2 A clear orange liquid and a blue powder are yes Unexpected color mixed together. The result is a clear colorless change liquid. 3 A silvery solid is placed in a clear blue liquid. yes Production of a gas The liquid bubbles furiously and the solid gradually disappears. 4 Two clear colorless liquids are mixed. The yes Production of heat beaker becomes very warm to the touch. 5 A clear yellow liquid and a clear colorless yes Formation of a solid, liquid are mixed. The result is a cloudy blue unexpected color solution. Upon setting, the solution separates change into a clear colorless liquid and a bright blue solid.

Part II Names & Formulas 49. Identify the type and write the name of each of the following compounds. See end of answer key

a. FeBr2 o. H2SO4(aq)

b. CoS p. HBr(aq)

c. Co2S3 q. B2H6

d. SnO2 r. GeH4

e. N2O4 s. FeI3

f. XeF6 t. CoCl2

g. OF2 u. HC2H3O2(aq)

h. AsI3 v. PbO

i. Fe(C2H3O2)3 w. Cr(CN)3

j. BeO x. MnF2

k. MgI2 y. SnBr4

l. Na2S z. HNO2(aq)

m. CaH2

n. HCl(aq) 50. Identify the type and write the formula of each of the following compounds. See end of answer key a. calcium chloride n. silver perchlorate b. aluminum sulfide o. copper(I) bicarbonate c. beryllium bromide p. potassium acetate d. hydrosulfuric acid q. cesium sulfite e. potassium hydride r. iron (II) phosphate f. magnesium iodide s. nickel(II) oxalate g. cesium fluoride t. ammonium peroxide h. sulfur dioxide u. gold(III) iodate i. dinitrogen monoxide v. lead(IV) phosphate j. tetraphosphorus decoxide w. manganese(II) hydroxide k. sulfur hexafluoride x. sulfurous acid l. nitrogen dioxide y. acetic acid m. copper(II) nitrate z. hydroiodic acid

Part III Equations For each of the following, a) underline the reactants and circle the products, b) write a word equation, c) write and balance a symbolic equation. 51. Hydrogen and bromine react to form gaseous hydrogen bromide.

H2(g) + Br2(l)  2HBr(g)

52. Solutions of iron (III) chloride and sodium hydroxide react to form solid iron (III) hydroxide and a solution of sodium chloride.

FeCl3(aq) + 3NaOH(aq)  Fe(OH)3(s) + 3NaCl(aq) 53. Liquid carbon disulfide reacts with oxygen to produce carbon dioxide gas and sulfur dioxide gas.

CS2(l) + 2O2(g)  CO2(g) + SO2(g)

53. A piece of copper is placed in a solution of silver nitrate. Silver metal and a solution of copper (II) nitrate are formed.

Cu(s) + 2AgNO3(aq)  2Ag(s)+ Cu(NO3)2(aq)

54. When heated, ammonium nitrate decomposes to form dinitrogen monoxide gas and gaseous water.

4NH4NO3(s)  4N2O(g) +8t H2O(g)

Balance the following equations.

55. 2Na(s) + _2_H2O(l) → 2 NaOH(aq) + _____H2(g)

56. _____CuCl2(s) + _____Na2SO4(aq) → 2__NaCl(aq) + _____CuSO4(s)

57. __3__Li(s) + _____AuCl3(aq) → __3__LiCl(aq) + _____Au(s)

58. _____C10H8(s) +_12_ O2(g) → _10__CO2(g) + _4__H2O(l)

59. _2__NaN3(s) → __2_ Na(s) + _3_N2(g)

60. _____BaCl2(aq) + _____K2CO3(aq) → _____BaCO3(s) + __2_KCl(aq)

61. _____H2SO4(aq) + __2__KOH(aq) → _____K2SO4(aq) + __2__HOH(l)

62. _____(NH4)3PO4(aq) + _____CrBr3(aq) → _3__NH4Br(aq) + _____CrPO4(s)

63. _____Al(s) + _3_CuCl(aq) → _3___Cu(s) + _____AlCl3(aq)

For each of the following reactions give the type of reaction. If the reaction would occur 1) predict the product(s) and balance the equation (show your work) 2) if it is a redox reaction determine which species is oxidized and which species is reduced (show your work.) If it is not a redox reaction, write NOT A REDOX REACTION. If no reaction would occur, write NO REACTION and give a reason.

64. HNO3(aq) + Sr(OH)2(aq) →

65. Mg(s) + N2(g) →

66. Na(s) + Cl2(g) →

67. Sn(s) + AgNO3(aq) →

68. CH4(g) + O2(g) →

69. AgNO3(aq) + NaCl(aq) →

70. Al(s) + H2(g) →

71. H2SO4(aq) + KOH(aq) →

72. Na2CO3(aq) + NH4Cl(aq) →

73. C6H12O6(aq) + O2(g) →

74. HgO(s) →

75. Cr(s) + Pb(NO3)2(aq) →

76. KOH(aq) + MgI2(aq) →

77. HCl(aq) + Ba(OH)2(aq) →

78. HBr(aq) + NaOH(aq) →

79. Ag(s) + Zn(NO3)2(s) →

80. K(s) + Cl2(g) →

81. Mg(s) + O2(g) →

82. Na3PO4(aq) + Ca(OH)2(aq) →

For each of the following reactions give the type of reaction. If it is a redox reaction determine which species is oxidized and which species is reduced (show your work.) If it is not a redox reaction, write NOT A REDOX REACTION.

83. KClO3(s) → KCl(s) + O2(g)

84. C(s) + O2(g) → CO2(g)

85. CaCO3(s) → CaO(s) + CO2(g)

86. K2SO3(s) → K2O(s)+ SO2(g)

Part IV Calculations:

87. How many atoms of lead are in 6.25 x 10-7 moles of lead?

88. How many moles of barium are in 8.08 g of barium?

89. What is the molar mass of Sr3(PO4)2?

6 90. What is the mass (in grams) of 3.03 x 10 moles of MgSO4? The molar mass of MgSO4 is 119.28 g MgSO4/1mole MgSO4.

91. For the balanced chemical equation, 4Fe(s) + 3O2(g) → 2Fe2O3(s), how many -3 moles of Fe2O3 are formed from 6.35 x 10 moles of iron? The molar mass of Fe2O3 is 159.7 g Fe2O3/1 mole Fe2O3.

92. For the balanced chemical equation, 2LiOH(s) + CO2(g) → Li2CO3(s) + H2O(g), what 6 mass(in grams) of LiOH required to produce 8.78 x 10 g of Li2CO3? The molar mass of LiOH is 23.95 g LiOH/1 mole LiOH. The molar mass of Li2CO3 is 73.89 g Li2CO3/ 1 mole Li2CO3. 49.

a. FeBr2 Binary Ionic Type II iron(II) bromide b. CoS Binary Ionic Type II cobalt(II) sulfide

c. Co2S3 Binary Ionic Type II cobalt (III) sulfide

d. SnO2 Binary Ionic Type II tin(IV) oxide

e. N2O4 Binary Covalent dinitrogen tetroxide

f. XeF6 Binary Covalent xenon hexafluoride

g. OF2 Binary Covalent oxygen difluoride

h. AsI3 Binary Covalent arsenic triiodide

i. Fe(C2H3O2)3 Nonbinary Ionic Type II iron (III) acetate j. BeO Binary Ionic Type I beryllium oxide

k. MgI2 Binary Ionic Type I magnesium iodide

l. Na2S Binary Ionic Type I sodium sulfide

m. CaH2 Binary Ionic Type I calcium hydride

n. HCl(aq) Binary acid hydrochloric acid

o. H2SO4(aq) oxyacid sulfuric acid

p. HBr(aq) binary acid hydrobromic acid

q. B2H6 Binary covalent diboron hexahydride

r. GeH4 Binary covalent germanium tetrahydride

s. FeI3 Binary Ionic Type II iron (III) iodide

t. CoCl2 Binary Ionic Type II cobalt (II) chloride

u. HC2H3O2(aq) oxyacid acetic acid v. PbO Binary Ionic Type II lead (II) oxide

w. Cr(CN)3 Nonbinary Ionic Type (II) chromium(III) cyanide

x. MnF2 Binary Ionic Type (II) manganese(II) fluoride

y. SnBr4 Binary Ionic Type (II) tin(IV) bromide

z. HNO2(aq) oxyacid nitrous acid

50. a. calcium chloride Binary Ionic Type I CaCl 2

b. aluminum sulfide Binary Ionic Type (I) Al2S3

c. beryllium bromide Binary Ionic Type I BeBr2

d. hydrosulfuric acid Binary Acid H2S(aq) e. potassium hydride Binary Ionic Type I KH

f. magnesium iodide Binary Ionic Type I MgI2 g. cesium fluoride Binary Ionic Type I CsF

h. sulfur dioxide Binary covalent SO2

i. dinitrogen monoxide Binary Covalent N2O

j. tetraphosphorus decoxide Binary covalent P4O10

k. sulfur hexafluoride Binary Covalent SF6

l. nitrogen dioxide Binary Covalent NO2

m. copper(II) nitrate Nonbinary Ionic Type (II) Cu(NO3)2

n. silver perchlorate Nonbinary Ionic Type I AgClO4

o. copper(I) bicarbonate Nonbinary Ionic Type II CuHCO3

p. potassium acetate Nonbinary Ionic Type I KC2H3O2

q. cesium sulfite Nonbinary Ionic Type I Cs2SO3

r. iron (II) phosphate Nonbinary Ionic Type II Fe3(PO4)2

s. nickel(II) oxalate Nonbinary Ionic Type II NiC2O4

t. ammonium peroxide Nonbinary Ionic Type I (NH4)2O2 this is an exception

u. gold(III) iodate Nonbinary Ionic Type II Au(IO3)3

v. lead(IV) phosphate Nonbinary Ionic Type II Pb3(PO4)4

w. manganese(II) hydroxide Nonbinary Ionic Type II Mn(OH)2

x. sulfurous acid oxyacid H2SO3(aq)

y. acetic acid oxyacid HC2H3O2(aq)

z. hydroiodic acid Binary acid HI(aq)