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General Chemistry I > Final Exam Practice > Flashcards

Final Exam Practice Flashcards

Gain a better understanding of General Chemistry concepts

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1
Q

The basic unit for mass in the metric system is

A

the gram

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2
Q

Which units are less than a gram?

A

a milligram and a centigram

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3
Q

The meaning of the prefix milli is

A

1/1000

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4
Q

What is the length 485 cm in kilometers?

A

4.85*10^-3 km

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5
Q

How many significant digits are allowed in the answer to the following calculation?

(3736.4 mi)/(55.0 hr) = 67.934545 mi/hr

A

3 sig. digits

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6
Q

A person who weighs 100 pounds has a mass of about

1 kg = 2.20 lbs

A

45 kg

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7
Q

A length of 25 cm is about

A

10 inches

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8
Q

Which of the following relations is not correct?

a. 12 inches = 1 foot
b. 1 inch = 1/12 foot
c. 60 seconds = 1 minute
d. 1 hour = 60 minutes
e. all of these are correct

A

e. all of these are correct

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9
Q

A gallon is equal to 3.785 liters. How many gallons are in 245 liters?

A

64.7 gal

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10
Q

The symbol for the element silver is

A

Ag

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11
Q

Neutral atoms

a. are highly charged
b. have the same number of electrons and protons
c. have more protons than electrons
d. are more dense on the outside than on the inside
e. all of these

A

b. have the same number of electrons and protons

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12
Q

In terms of mass

a. electrons are much less massive than protons
b. electrons are much less massive than neutrons
c. protons and neutrons have similar masses
d. all of these
e. none of these

A

d. all of these

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13
Q

The atomic number of an element is equal to the

  1. number of protons in the nucleus of an atom
  2. number of electrons in a neutral atom of the element
  3. sum of the protons and neutrons
  4. sum of the protons and electrons
A

1 and 2

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14
Q

The nucleon number of an atom is

A

the number of electrons

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15
Q

The atomic weight of an element is

  1. the average number of protons
  2. the average number of protons plus neutrons
  3. the average number of protons, neutrons and electrons
  4. the average mass of an atom considering the contributions of all the natural isotopes
A

2

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16
Q

An atom with 50 protons, 50 electrons and 66 neutrons has a nucleon number equal to

A

116

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17
Q

Isotopes of an element have an equal number of

A

protons, different number of neutrons

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18
Q

A column in the periodic table is called a

A

group or family

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19
Q

The part of the Periodic Table with most of the elemental gases is

A

the upper right portion

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20
Q

A row in the periodic table is called a

A

period

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21
Q

Elements that are shiny, malleable and good conductors are

A

metals

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22
Q

Those elements that are brittle solids or gases and nonconductors are

A

nonmetal

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23
Q

The periodic table was first established by

A

Dmitri Mendeleev in 1869

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24
Q

An atom with atomic number 17, and nucleon number 37 has

  1. 37 protons, 17 neutrons
  2. 17 protons, 20 neutrons
  3. 37 electrons, 17 neutrons
  4. 17 electrons, 20 neutrons
A

2 and 4

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25
Q

The molecular formula of a particular solid is C8H4O8. Its molecular mass is

A

228 amu

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26
Q

What is the mass in grams of one propane, C3H8 molecule?

A

7.32*10^-23

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27
Q

Which of the following compounds contains the largest number of atoms?

a. 1.0 mol of CH3COCl
b. 2.0 mol of H2SO4
c. 3.0 mol of NH3
d. 4.0 mol of K2S
e. 5.0 mol of HBr

A

2.0 mol of H2SO4

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28
Q

What is the mass of oxygen atoms in 0.395 mol Fe(CO)5?

A

31.6 g

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29
Q

What is the mass in grams of 0.732 mol of glucose, C6H12O6?

A

132 g

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30
Q

How many moles of silver are contained in 7.00 kg of silver?

A

64.9 mol

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31
Q

How many molecules are there in 60.0 g of acetic acid, C2H4O2?

A

6.02*10^23

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32
Q

How many molecules are there in 2.47 mg of mannose, C6H12O6, which if a sweet-tasting sugar that has a bitter aftertaste?

A

8.26*10^18

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33
Q

The total number of oxygen atoms in 2.60 g of CaCo3 (MM = 100.0 g/mol) is

A

4.70*10^22

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34
Q

What is the percentage by mass of sulfur in copper (I) sulfide, Cu2S?

A

20.1%

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35
Q

Analysis of a compound showed that it contained 14.4% hydrogen atoms and 85.6% carbon atoms by mass. What is its empirical formula?

A

CH2

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36
Q

How do you find the empirical formula?

A

Percent to mass
Mass to mole
Divide by small
Multiply til’ whole

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37
Q

Of the following the only empirical formula is

a. C2H8
b. C5H14
c. C6H12
d. H2O2
e. O3

A

C5H14

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38
Q

A 0.0111-mol sample of urea, NH2CONH2 contains

A

5.35*10^22 atoms

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39
Q

What is Avogadro’s number?

A

the number of atoms or molecules in one mole of a substance, equal to 6.023 × 10^23

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40
Q

How many molecules are there in 62.0 g of ethylene glycol, HO(CH2)2OH?

A

6.02*10^23

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41
Q

How many molecules are there in 60.0 g of acetic acid, C2H4O2?

A

6.02*10^23

42
Q

The correct name for FeO is

A

iron (II) oxide

43
Q

The formula for hypochlorous acid is

A

HClO

44
Q

How many moles of hexachlorobenezene, C6Cl6, are there in 5.86 g of C6Cl6?

A

0.0206 mol

45
Q

Which of the following ions is most likely to form an insoluble sulfate?

a. K+
b. Li+
c. Ca2+
d. S2-
e. Cl-

A

Ca2+

46
Q

Which of the following compounds is soluble in water?

a. Sc2S3
b. BaS
c. Cs2S
d. Hg2S2
e. Ag2S

*See solubility rules

A

BaS

47
Q

Which of the following compounds is insoluble in water?

a. (NH4)2CO3
b. FeCO3
c. K2CO3
d. Li2CO3
e. Na2CO3

*See insolubility rules

A

FeCO3

48
Q

Identify the spectator ions in the following reaction.

Cu(OH)2(s) + 2H+(aq) + 2Cl-(aq) produces Cu2+(aq) + 2Cl-(aq) + 2H2O(l)

A

Cl-

49
Q

What are spectator ions?

A

An ion that exists in the same form on both the reactant and product sides of a chemical reaction

50
Q

When a solution of lithium chloride and a solution of ammonium sulfate are mixed,

A

no reaction occurs

51
Q

When solutions of barium chloride and sodium sulfate are mixed, the spectator ions in the resulting reaction are

A

both Na+ and Cl-
BaCl2(aq) + Na2SO4(aq) produces BaSO4(s) + 2NaCl(aq)

Ba2+ + 2Cl- +Na+ + 2SO42- produces BaSO4(s) + 2Na+ + Cl-

52
Q

Which of the following is not an Arrhenius base?

a. LiOh
b. NH3
c. Ca(OH)2
d. Sr(OH)2
e. KOH

A

NH3

53
Q

Which of the following reactions best describes the dissolution of solid RbOH(s) in water?

a) RbOH(s) → RbO–(aq) + H+(aq)
b) RbOH(s) → RbO+(aq) + H–(aq)
c) RbOH(s) → Rb+(aq) + OH–(aq)
d) RbOH(s) + H2O(l) → RbO–(aq) + H3O+(aq)
e) RbOH(s) → RbOH(aq)

A

c) RbOH(s) → Rb+(aq) + OH–(aq)

54
Q

Which of the following test describes the acid-base properties of ammonia (NH3) in aqueous solution?

a) NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)
b) NH3(aq) + H2O(l) → NH4+(aq) + OH–(aq)
c) NH3(aq) + 3H2O(l) N3–(aq) + 3H3O+(aq)
d) NH3(aq) + 3H2O(l) → N3–(aq) + 3H3O+(aq)
e) NH3(aq) → N3–(aq) + 3H+(aq)

A

a) NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)

55
Q

Which of the following reactions best describes the dissolution of gaseous hydrogen chloride (HCl) in water?

a. HCl(aq) + H2O (l) –- > H3O+(aq) + Cl- (aq)
b. HCl(aq) + H2O (l) –- > H3O+(aq) + ClO- (aq)
c. HCl(aq) + 2H2O (l) –- > H3O+(aq) + ClO- (aq) + H2(g)
d. 2HCl(aq) + H2 (g) –- > Cl2 (g)
e. HCl(aq) + 5H2O (l) –- > H3O+(aq) + ClO4- (aq) + 4H2(g)

A

a. HCl(aq) + H2O (l) –- > H3O+(aq) + Cl- (aq)

56
Q

Which of the following is a strong base aqueous solution?

a. NH3
b. Ca(OH)2
c. CH3OH
d. CH3CH2OH
e. HOCH2CH2OH

A

b. Ca(OH)2

57
Q

Which of the following is a strong base in aqueous solution?

a. HClO4
b. CH3OH
c. NH3
d. KCl
e. KOH

A

e. KOH

58
Q

Which of the following is a weak acid in aqueous solution?

a. HNO3
b. Hal
c. H2SO4
d. H2SO3
e. HClO4

A

d. H2SO3

59
Q

Which of the following is necessary in order for a metal to be oxidized?

a. addition of hydrogen
b. removal of oxygen
c. removal of electrons
d. addition of electrons
e. addition of oxygen

A

c. removal of electrons

* oxidized - oxidation number increases meaning electrons are lost

60
Q

Which of the following is an oxidation-reduction reaction?

a. CO2(g) + H2O(l) –> H2CO3(q)
b. SO3(g) + H2O(l) –> H2SO4(aq)
c. Na2O(s) + H2O(l) –> 2NaOH(aq)
d. PbCO3(s) + 2HNO3(aq) –> Pn(NO3)2(aq) + CO2(g) + H2O(l)
e. C2H4(g) + H2(g) –> C2H4(g)

A

e. C2H4(g) + H2(g) –> C2H4(g)

61
Q

All of the following are oxidation-reduction reactions except

a. H2(g) + F2(g) –> 2HF(g)
b. 6Li(s) + N2(g) –> 2Li3N(s)
c. Ca(s) + H2(g) –> CaH2(s)
d. Mg3N2(s) + 6H2O(l) –> 3Mg(OH)2(s) +2NH3(g)
e. 2K(s) + 2H2O(l) –> 2KOH(aq) + H2(g)

A

d. Mg3N2(s) + 6H2O(l) –> 3Mg(OH)2(s) +2NH3(g)

62
Q

In the following reaction, which species is oxidized?

a.

A

a. Na2S

63
Q

In the following reaction, which species is oxidized?
3Na2S(s) + 8H+(aq) + 2NO3(aq) –> 6Na(aq) + 3S(s) +2NO(g) +4H2O
a. Na2S
b. H+
c. No3-
d. NO
e. Na+

A

a. Na2S

64
Q 
For the reaction that occurs in a lead storage battery,
Pb(s) + PbO2(s) + 2H+ (aq) + (2HSO4)- (aq) --> 2PbSO4 (s) +2H2O (l),
The oxidizing agent is?
a. Pb
b. PbO2
c. H+
d. (HSO4)-
e. PbSO4
A

b. PbO2

65
Q

What does oxidized mean?

A

When an atom is oxidized, it loses electrons, and its oxidation number increases.

66
Q

What does reduced mean?

A

When an atom is reduced, it gains electrons, and its oxidation number decreases.

67
Q

The reactant containing the element being oxidized is called the

A

reducing agent

68
Q

The reactant containing the atom being reduced is called the

A

oxidizing agent

69
Q

What are the six strong acids?

A

HCl, HBr, HI, HNO3, HClO4, and H2SO4

70
Q

Name a few weak acids

A

HF, HNO2, H3PO4

71
Q

What are the common strong bases?

A 
LiOH - lithium hydroxide
NaOH - sodium hydroxide
KOH - potassium hydroxide
RbOH - rubidium hydroxide
CsOH - cesium hydroxide
*Ca(OH)2 - calcium hydroxide
*Sr(OH)2 - strontium hydroxide
*Ba(OH)2 - barium hydroxide
72
Q

If q = -31 kJ for a certain process, that process

a. occurs rapidly
b. is exothermic
c. is endothermic
d. cannot occur
e. requires a catalyst

A

b. is exothermic

73
Q

At constant pressure, the sign q for the process H2O(l) –> H2O(s) is expected to be

a. negative, and the process is endothermic
b. positive, and the process is exothermic
c. positive, and the process is endothermic
d. negative, and the process exothermic
e. impossible to predict

A

d. negative, and the process exothermic

74
Q

H2 and F2 react according to the following equation, forming HF.
H2(g) + F2(g) → 2HF(g); ∆H° = –271 kJ

If H2(g) and F2(g) were mixed in a thermally insulated vessel, the reaction that occurred would be

a. exothermic, and the temperature of the reaction system would rise.
b. endothermic, and the temperature of the reaction system would rise.
c. endothermic, and the temperature of the reaction system would fall.
d. exothermic, and the temperature of the reaction system would fall.
e. We could not tell unless the original and final temperatures were given.

A

a. exothermic, and the temperature of the reaction system would rise.

75
Q

Which of the following statements is incorrect?

a. The value of q is positive when heat flows into a system from the surroundings.
b. Heat flows from a system into the surroundings in an endothermic process.
c. Enthalpy is a state function.
d. Internal energy is a state function.
e. The value of q is positive in an endothermic process.

A

b. Heat flows from a system into the surroundings in an endothermic process.

76
Q

Which of the following statements about enthalpy is false?

a. Enthalpy is a state function
b. At constant pressure, the enthalpy change is equal to the heat absorbed or released.
c. Enthalpy is an extensive property.
d. The change in enthalpy of a process cannot be negative.
e. The SI unit of enthalpy is J.

A

d. The change in enthalpy of a process cannot be negative.

77
Q

The phrase “the heat absorbed or released by a system undergoing a physical or chemical change at constant pressure” is

A

the change in enthalpy of the system

78
Q

The quantity of heat required to raise the temperature of a sample of a substance by 1°C is the sample’s

A

heat capacity

79
Q

The units for heat capacity are

A

J/°C

80
Q

The units for specific hear are

A

J/(g • °C)

81
Q

Consider the following metals.

Metal

A

e. lead

82
Q

How much heat is gained by nickel when 29.2 g of nickel is warmed from 18.3°C to 69.6°C? The specific heat of nickel is 0.443 J/(g • °C).

A

6.64*10^2 J

83
Q

Exactly 273.5 J will raise the temperature of 10.0 g of a metal from 25.0°C to 60.0°C. What is the specific heat capacity of the metal?

A

0.781 J/(g • °C)

84
Q

Which of the following has a standard enthalpy of formation value of zero at 25°C?

a. H2O(l)
b. C6H12O(s)
c. FeSO4(s)
d. FeSO4(aq)
e. S8(s)

A

e. S8(s)

* pure element

85
Q

From the following information, determine the enthalpy of formation C2H4(g).
1/2C2H4(g) → C(s) + H2(g); ΔH = –26.2 kJ
a. –26.2 kJ/mol
b. 26.2 kJ/mol
c. 104.8 kJ/mol
d. –52.4 kJ/mol

A

d. -52.4

86
Q 
Consider the following changes: 
H20(s)----> H2O(l) ΔH1
H20(l)----> h20(g)  ΔH2
H2O(g) ---> H2(s)  ΔH3
Using Hess Law, the sum ΔH1 +ΔH2 +ΔH3 is

a. >0
b. = 0
c. < 0
d. sometimes > 0 and sometimes < 0
e. cannot be determined from this information

A

b. = 0

87
Q

The energy associated with a speeding bullet is called

A

kinetic energy

*energy which a body possesses by virtue of being in motion.

88
Q

The energy associated with a motionless rock on the top of Mount Everest is

A

potential energy
*the energy possessed by a body by virtue of its position relative to others, stresses within itself, electric charge, and other factors

89
Q

All of the following statements are true for an exothermic reaction except

a. The products have a higher heat content than the reactants.
b. Heat passes from the reaction system to the surroundings.
c. The temperature of the reaction system decreases.
d. The temperature of the surroundings increases.
e. The enthalpy change for the reaction is negative

A

a. The products have a higher heat content than the reactants.

90
Q

Making sure that the number of symbols of each element is the same on both sides of an equation is called

A

balancing the equation

91
Q

Which of the following statements is false? A balanced chemical equation

a. has the same number of combined atoms of each element on each side
b. corresponds to the law of conservation of matter
c. uses coefficients to indicate the number of formula units of each reactant and product
d. requires changes of subscripts and formulas to balance
e. none of these

A

d. requires changes of subscripts and formulas to balance

92
Q

A chemical reaction

A

can be described by a chemical equation

93
Q

How many grams of potassium carbonate are needed to make 200 mL of a 2.5 M solution?

A

69.1 grams

94
Q

What is the concentration of an aqueous solution with a volume of 450 mL that contains 200 grams of iron (II) chloride?

A

3.51 M

95
Q

How many liters of 4 M solution can be made using 100 grams of lithium bromide?

A

3.47 L

96
Q

For the reaction
2S(s) + 302(g) –> 2SO3(g)

If 6.3 g of S is reacted with 10.0 g of O2 show by calculation which one will be the limiting reactant.

A

Sulfur

97
Q

For the reaction
CaC03(s) + 2HCl(aq) –> CaC12(aq) + CO2(g) + H20(l)

68.1 g solid CaC03 is mixed with 51.6 g HCl. What number of grams of CO2 will be produced?

A

29.9 g CO2

98
Q

Consider the reaction
Mg2Si(s) + 4H20(l) –> 2Mg(OHh(aq) + SiH4(g)

Which of the reactants is in excess if we start with 50.0 g of each reactant?

A

H2O

99
Q

A handheld “red dot” laser pointer that is commonly used to certain its ( and humans), has a wavelength of 650 nm.

Calculate the frequency of this light.

A

4.612*10^14 Hz

100
Q

Continued:

Calculate the energy of a single photon of this light.

A
  1. 06*10^-14

* E=hv

General Chemistry I (1 decks)

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