General Chemistry I Mock Exam
Course Introduction
General Chemistry I is an introductory course that lays the foundation for understanding the basic principles of chemistry. Topics covered include atomic and molecular structure, stoichiometry, chemical bonding, states of matter, thermochemistry, periodic trends, and properties of solutions. The course emphasizes problem-solving skills and laboratory techniques to help students develop a conceptual and practical understanding of chemical phenomena, preparing them for advanced studies in chemistry and related disciplines.
Recommended Textbook Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg
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Chapter 1: Keys to the Study of Chemistry
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Sample Questions
Q1) The mass of a sample is 550 milligrams. Which of the following expresses that mass in kilograms?
A) 5.5 × 10<sup>8 </sup>kg
B) 5.5 × 10<sup>5</sup> kg
C) 5.5 × 10<sup>-4 </sup>kg
D) 5.5 × 10<sup>-6 </sup>kg
E) 5.5 × 10<sup>-1 </sup>kg
Answer: C
Q2) A detailed explanation of natural phenomena that is generally accepted and has been extensively tested is called a A) theory.
B) hypothesis.
C) law.
D) fact.
E) postulate.
Answer: A
Q3) The ripening of fruit, once picked, is an example of physical change.
A)True
B)False
Answer: False
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Chapter 2: The Components of Matter
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Sample Questions
Q1) Which of the following ions occurs commonly?
A) P<sup>3+</sup>
B) Br<sup>7+</sup>
C) O<sup>6+</sup>
D) Ca<sup>2+</sup>
E) K<sup>-</sup>
Answer: D
Q2) Calcium hydroxide is used in mortar, plaster, and cement. What is its formula?
A) CaOH
B) CaOH<sub>2</sub>
C) Ca<sub>2</sub>OH
D) Ca(OH)<sub>2</sub>
E) CaHO<sub>2</sub>
Answer: D
Q3) The molecular formula of a compound provides more information than its structural formula.
A)True
B)False
Answer: False
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Chapter 3: Stoichiometry of Formulas and Equations
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Sample Questions
Q1) Gaseous methanol (CH<sub>4</sub>O) reacts with oxygen gas to produce carbon dioxide gas and liquid water. Write a balanced equation for this process.
Answer: 2CH<sub>4</sub>O(g) + 3O<sub>2</sub>(g) \(\to\) 2CO<sub>2</sub>(g) + 4H<sub>2</sub>O(l)
Q2) Magnesium fluoride is used in the ceramics and glass industry. What is the mass of 1.72 mol of magnesium fluoride?
A) 43.3 g
B) 62.3 g
C) 74.5 g
D) 92.9 g
E) 107 g
Answer: E
Q3) Balance the equation
B<sub>2</sub>O<sub>3</sub>(s) + NaOH(aq) \(\to\) Na<sub>3</sub>BO<sub>3</sub>(aq) + H<sub>2</sub>O(l)
Answer: B<sub>2</sub>O<sub>3</sub>(s) + 6NaOH(aq) \(\to\) 2Na<sub>3</sub>BO<sub>3</sub>(aq) + 3H<sub>2</sub>O(l)
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Chapter 4: The Major Classes of Chemical Reactions
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Sample Questions
Q1) In a redox reaction, the reducing agent undergoes loss of electrons.
A)True
B)False
Q2) What is the oxidation number of iodine in I<sub>2</sub>.
A) -1
B) 0
C) +1
D) +7
E) None of the above is the correct oxidation number.
Q3) The compound PS1U1B14S1U1B0OS1U1B110S1U1B0is used in refining sugar. Select the classification for the reaction in which it is synthesized. PS1U1B14S1U1B0(s) + 5OS1U1B12S1U1B0(g) \(\rarr\)PS1U1B14S1U1B0OS1U1B110S1U1B0(s)
A) combination
B) decomposition
C) displacement
D) acid-base
E) precipitation
Q4) Covalent compounds, dissolved in water, never produce conducting solutions.
A)True
B)False
Page 6
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Chapter 5: Gases and the Kinetic-Molecular Theory
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Sample Questions
Q1) Select the statement which does NOT apply to an ideal gas.
A) There are no attractive forces between the gas molecules.
B) There are strong repulsive forces between the gas molecules.
C) The volume occupied by the molecules is negligible compared to the container volume.
D) The gas behaves according to the ideal gas equation.
E) The average kinetic energy of the molecules is proportional to the absolute temperature.
Q2) The ozone layer is important because
A) ozone absorbs low energy radiation which warms the troposphere.
B) ozone purifies the atmosphere by reacting with excess fluorocarbons.
C) ozone absorbs ultraviolet radiation.
D) ozone reflects high energy radiation such as X-rays and gamma rays.
E) humans need to breathe air containing some ozone.
Q3) Nitrogen will behave most like an ideal gas
A) at high temperature and high pressure.
B) at high temperature and low pressure.
C) at low temperature and high pressure.
D) at low temperature and low pressure.
E) at intermediate (moderate) temperature and pressure.
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Chapter 6: Thermochemistry: Energy Flow and Chemical Change
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Sample Questions
Q1) The compound carbon suboxide, C<sub>3</sub>O<sub>2</sub>, is a gas at room temperature. Use the data supplied to calculate the heat of formation of carbon suboxide.
(Data: 2CO(g) + C(s) \(\to\) C<sub>3</sub>O<sub>2</sub>(g) \(\Delta\)H° = 127.3 kJ/mol
And: \(\Delta\)H<sub>f</sub><sup>°</sup> of CO(g) = -110.5 kJ/mol)
A) 116.8
B) -93.7
C) 227.8
D) -348.3
E) 93.7
Q2) Which of the following is not a state function?
A) internal energy
B) volume
C) work
D) pressure
E) enthalpy
Q3) For all processes, both q and \(\Delta\)E will have the same sign.
A)True
B)False
Page 8
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Chapter 7: Quantum Theory and Atomic Structure
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Sample Questions
Q1) What are the possible values for the following quantum numbers in an atom?
a. n
b. l
c. m<sub>l</sub>
Q2) A wave function for an electron is called an atomic orbital.
A)True
B)False
Q3) According to the Bohr theory of the hydrogen atom, the minimum energy (in J) needed to ionize a hydrogen atom from the n = 2 state is
A) 2.18 × 10<sup>-18 </sup>J.
B) 1.64 × 10<sup>-18 </sup>J.
C) 5.45 × 10<sup>-19 </sup>J.
D) 3.03 × 10<sup>-19 </sup>J.
E) none of the above.
Q4) Which word best describes the phenomenon which gives rise to a rainbow?
A) reflection
B) dispersion
C) diffraction
D) interference
E) deflection
Page 9
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Chapter 8: Electron Configuration and Chemical Periodicity
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Sample Questions
Q1) Select the diamagnetic ion.
A) Cu<sup>2+</sup>
B) Ni<sup>2+</sup>
C) Cr<sup>3+</sup>
D) Sc<sup>3+</sup>
E) Cr<sup>2+</sup>
Q2) Which of the following elements is paramagnetic?
A) Kr
B) Zn
C) Sr
D) V
E) Ar
Q3) Elements with _______________ first ionization energies and ___________ electron affinities generally form anions.
A) low, very negative
B) high, positive or slightly negative
C) low, positive or slightly negative
D) high, very negative
E) None of the above is generally correct.
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Chapter 9: Models of Chemical Bonding
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Sample Questions
Q1) Which of the following is an ionic compound?
A) H<sub>2</sub>S
B) NH<sub>3</sub>
C) I<sub>2</sub>
D) KI
E) CCl<sub>4</sub>
Q2) Which of the following contains ionic bonding?
A) CO
B) SrF<sub>2</sub>
C) Al
D) OCl<sub>2</sub>
E) HCl
Q3) When one mole of each of the following liquids is burned, which will produce the most heat energy?
A) C<sub>6</sub>H<sub>14</sub>
B) C<sub>5</sub>H<sub>12</sub>
C) C<sub>6</sub>H<sub>14</sub>O
D) C<sub>6</sub>H<sub>12</sub>O
E) C<sub>6</sub>H<sub>10</sub>O<sub>3</sub>
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Chapter 10: The Shapes of Molecules
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Sample Questions
Q1) Predict the ideal bond angles in GeCl<sub>4 </sub>using the molecular shape given by the VSEPR theory.
A) 90°
B) 109°
C) 120°
D) 180°
E) < 90°
Q2) Using SO<sub>2</sub> as an example, describe the sort of experimental data which might suggest that no single Lewis structure is an accurate representation of its bonding.
Q3) Predict the ideal bond angles in IF<sub>2</sub><sup>-</sup><sub> </sub>using the molecular shape given by the VSEPR theory.
A) 60°
B) 90°
C) 109°
D) 120°
E) 180°
Q4) List the three important ways in which molecules can violate the octet rule, and in each case draw one Lewis structure of your choice as an example.
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Chapter 11: Theories of Covalent Bonding
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Q1) a. What simple experiment could you perform to show that a substance is paramagnetic?
b. What microscopic (atomic/molecular) feature must a substance possess in order to be paramagnetic?
c. Can it be predicted whether or not all homonuclear diatomic ions, X<sub>2</sub><sup>+</sup>, will be paramagnetic? Explain.
Q2) In one sentence state how molecular orbitals are usually obtained.
Q3) A molecule with the formula AX<sub>4</sub>E uses _________ to form its bonds.
A) sp<sup>2</sup> hybrid orbitals
B) sp<sup>3</sup> hybrid orbitals
C) sp<sup>3</sup>d hybrid orbitals
D) sp<sup>3</sup>d<sup>2</sup> hybrid orbitals
E) none of the above
Q4) Hybrid orbitals of the sp<sup>3</sup>d type occur in sets of four.
A)True
B)False
Q5) Explain what is meant by the term "bond order" and describe how it can be calculated using the information in a molecular orbital energy level diagram.
Q6) In one sentence state the basic principle of valence bond theory.
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Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes
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Sample Questions
Q1) Which of the following pairs is arranged with the particle of higher polarizability listed first?
A) CCl<sub>4</sub>, CI<sub>4</sub>
B) H<sub>2</sub>O, H<sub>2</sub>Se
C) C<sub>6</sub>H<sub>14</sub>, C<sub>4</sub>H<sub>10</sub>
D) NH<sub>3</sub>, NF<sub>3</sub>
E) none of the above
Q2) Iron crystallizes in the body-centered cubic lattice. What is the coordination number for Fe?
A) 4
B) 6
C) 8
D) 10
E) 12
Q3) Which of the following atoms should have the greatest polarizability?
A) F
B) Br
C) Po
D) Pb
E) He
Page 14
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Chapter 13: The Properties of Mixtures: Solutions and Colloids
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Sample Questions
Q1) If the density of a solution is less than 1.0 g/mL, its molarity will be greater than its molality.
A)True
B)False
Q2) Calculate the molarity of a solution prepared by diluting 1.85 L of 6.5 M KOH to 11.0 L.
A) 0.28 M
B) 0.91 M
C) 1.1 M
D) 3.1 M
E) 3.9 M
Q3) Which of the following pairs of ions is arranged so that the ion with the larger (i.e., more negative) heat of hydration is listed first?
A) Br<sup>-</sup>, K<sup>+</sup>
B) Mg<sup>2+</sup>, Sr <sup>2+</sup>
C) Ca<sup>2+</sup>, Sc<sup>3+</sup>
D) Na<sup>+</sup>, Li<sup>+</sup>
E) None of the above are arranged in this way.
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Chapter 14: Periodic Patterns in the Main Group Elements:
Bonding, Structure, and Reactivity
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Sample Questions
Q1) Predict the products for the following set of reactants. Cs(s) + Br<sub>2</sub>(l) \(\to\)
A) CsBr(s)
B) CsBr<sub>2</sub>(s)
C) CsBr(l)
D) CsBr<sub>2</sub>(l)
E) Cs<sub>2</sub>Br(s)
Q2) The chemical that ranks first in production among all industrial chemicals is
A) NH<sub>3</sub>, ammonia.
B) H<sub>3</sub>PO<sub>4</sub>, phosphoric acid.
C) NaOH, sodium hydroxide.
D) Na<sub>2</sub>CO<sub>3</sub>, sodium carbonate.
E) H<sub>2</sub>SO<sub>4</sub>, sulfuric acid.
Q3) Sulfur hexafluoride is a pollutant responsible for acid rain.
A)True
B)False
Q4) Oxygen gas (O<sub>2</sub>) is paramagnetic, but ozone (O<sub>3</sub>) is diamagnetic.
A)True
B)False
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Chapter 15: Organic Compounds and the Atomic Properties of Carbon
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Sample Questions
Q1) The protein amino acid sequence, the RNA base sequence, and the DNA base sequence are interrelated. Which of the following descriptions is correct?
A) The RNA base sequence determines the DNA base sequence which, in turn, determines the protein amino acid sequence.
B) The DNA base sequence determines the RNA base sequence which, in turn, determines the protein amino acid sequence.
C) The DNA base sequence determines the protein amino acid sequence which, in turn, determines the RNA base sequence.
D) The RNA base sequence determines the protein amino acid sequence which, in turn, determines the DNA base sequence.
E) None of the above statements is correct.
Q2) What is the difference between 1-butyne and 1-butene?
A) 1 carbon atom
B) 1 carbon atom and 2 hydrogen atoms
C) 2 hydrogen atoms
D) 4 hydrogen atoms
E) 1 carbon atom and 4 hydrogen atoms
Q3) Explain what is meant by "complementary" in the context of DNA strands.
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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions
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Sample Questions
Q1) The radioactive isotope tritium decays with a first-order rate constant k of 0.056 year<sup>-1</sup>. What fraction of the tritium initially in a sample is still present 30 years later?
A) 0.19
B) 0.60
C) 0.15
D) 2.8 × 10<sup>-38</sup>
E) none of the above
Q2) In a reversible reaction, a catalyst will speed up the forward reaction but not affect the reverse reaction.
A)True
B)False
Q3) If the activation energy of a reaction decreases by 10.0 kJ/mol, from 100.0 to 90.0 kJ/mol, what effect will this have on the rate of reaction at 298K?
A) The rate will increase, by a factor of more than 50.
B) The rate will decrease, by a factor of more than 50.
C) The rate will increase, by a factor of less than 50.
D) The rate will decrease, by a factor of less than 50.
E) The rate will not change unless temperature changes.
Page 18
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Chapter 17: Equilibrium: the Extent of Chemical Reactions
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Q1) When a chemical system is at equilibrium,
A) the concentrations of the reactants are equal to the concentrations of the products.
B) the concentrations of the reactants and products have reached constant values.
C) the forward and reverse reactions have stopped.
D) the reaction quotient, Q, has reached a maximum.
E) the reaction quotient, Q, has reached a minimum.
Q2) A chemical reaction will reach equilibrium when the limiting reactant is used up.
A)True
B)False
Q3) A good catalyst for a reaction will speed up the forward reaction and slow down the reverse reaction.
A)True
B)False
Q4) When a reaction system reaches equilibrium, the forward and reverse reactions stop.
A)True
B)False
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Chapter 18: Acid-Base Equilibria
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Q1) Iodine trichloride, ICl<sub>3</sub>, will react with a chloride ion to form ICl<sub>4</sub><sup>-</sup>. Which species, if any, acts as a Lewis acid this reaction?
A) ICl<sub>4</sub><sup>-</sup>
B) ICl<sub> 3</sub>
C) Cl<sup>-</sup>
D) the solvent
E) None of the species acts as a Lewis acid in this reaction.
Q2) A solution is prepared by adding 0.10 mol of lithium nitrate, LiNO<sub>3</sub>, to 1.00 L of water. Which statement about the solution is correct?
A) The solution is basic.
B) The solution is neutral.
C) The solution is weakly acidic.
D) The solution is strongly acidic.
E) The values for K<sub>a</sub> and K<sub>b</sub> for the species in solution must be known before a prediction can be made.
Q3) Arrhenius bases raise the hydroxide ion concentration when dissolved in water.
A)True
B)False
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Chapter 19: Ionic Equilibria in Aqueous Systems
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Q1) A 50.0-mL sample of 0.50 M HCl is titrated with 0.50 M NaOH. What is the pH of the solution after 28.0 mL of NaOH have been added to the acid?
A) 0.85
B) 0.75
C) 0.66
D) 0.49
E) 3.8
Q2) The solubility of magnesium phosphate is 2.27 × 10<sup>-3 </sup>g/1.0 L of solution. What is the K<sub>sp</sub> for Mg<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub>?
A) 6.5 × 10<sup>-12</sup>
B) 6.0 × 10<sup>-14</sup>
C) 5.2 × 10<sup>-24</sup>
D) 4.8 × 10<sup>-26</sup>
E) 1.0 × 10<sup>-26</sup>
Q3) What is the pH of 375 mL of solution containing 0.150 mol of propenoic acid (HA) and 0.250 mol of sodium propenoate (NaA)? (K<sub>a</sub> for propenoic acid is 5.52 × 10<sup>-5</sup>.)
Q4) Make a clear distinction between buffer range and buffer capacity.
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Chapter 20: Thermodynamics: Entropy, Free Energy, and the Direction of Chemical Reactions
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Q1) For a chemical reaction to be spontaneous only at high temperatures, which of the following conditions must be met?
A) .\(\Delta\)S° > 0, \(\Delta\)H° > 0
B) .\(\Delta\)S° > 0, \(\Delta\)H° < 0
C) .\(\Delta\)S° < 0, \(\Delta\)H° < 0
D) .\(\Delta\)S° < 0, \(\Delta\)H° > 0
E) .\(\Delta\)G° > 0
Q2) Which relationship best describes \(\Delta\)S° for the following reaction? CO(g) + H<sub>2</sub>O(g) \(\to\) CO<sub>2</sub>(g) + H<sub>2</sub>(g)
A) .\(\Delta\)S° = \(\Delta\)H°
B) .\(\Delta\)S° = \(\Delta\)H°/T
C) .\(\Delta\)S° > 0
D) .\(\Delta\)S° < 0
E) .\(\Delta\)S°<sup> </sup> \(\approx\) 0
Q3) In a spontaneous process, the entropy of the system always increases.
A)True
B)False
Q4) In the expression, S = k ln W, W is called the number of microstates. Explain clearly the meaning of the word "microstate", and why a system under a given set of conditions normally has many microstates.
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Chapter 21: Electrochemistry: Chemical Change and Electrical Work
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Q1) Which of the following statements about voltaic and electrolytic cells is correct?
A) The electrons in the external wire flow from cathode to anode in both types of cell.
B) Oxidation occurs at the cathode only in a voltaic cell.
C) The free energy change, \(\Delta\)G, is negative for an electrolytic cell.
D) The cathode is labeled as positive (+) in a voltaic cell but negative (-) in an electrolytic cell.
E) Reduction occurs at the anode in an electrolytic cell.
Q2) Calculate the potential of a voltaic cell (E°<sub>cell</sub>) if it is required to do 5.43 × 10<sup>-3</sup> kJ of work when a charge of 2.50 C is transferred.
A) 2.17 × 10<sup>3</sup> V
B) 2.17 × 10<sup>-3</sup> V
C) 2.17 V
D) 13.6 V
E) 1.36 × 10<sup>-2</sup> V
Q3) For the reaction occurring in a voltaic (galvanic) cell, \(\Delta\)G > 0.
A)True
B)False
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Chapter 22: The Elements in Nature and Industry
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Q1) Nitrogen fixation occurs through atmospheric, industrial, and biological processes. Which of these fixes the most nitrogen?
A) biological
B) atmospheric
C) industrial
D) industrial \(\approx\) atmospheric
E) industrial \(\approx\) biological
Q2) The alkali metals are isolated from non-aqueous systems. Why is this necessary?
A) The electrolysis of aqueous solutions of the alkali metals requires more energy than electrolysis of the molten salts.
B) The dissolved alkali earth halides are too reactive to be electrolyzed.
C) The aqueous metal ions are more difficult to reduce than water.
D) The reduction potentials of the alkali metals are more positive than the reduction potential of water.
E) The aqueous metal ions react violently with water.
Q3) The bond energy in the hydrogen molecule (H<sub>2</sub>) is greater than that of the tritium molecule (T<sub>2</sub>).
A)True
B)False
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Chapter 23: The Transition Elements and Their Coordination Compounds
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Q1) Which of the following ions is least likely to form colored compounds?
A) Mn<sup>2+</sup>
B) Cr<sup>5+</sup>
C) Sc<sup>3+</sup>
D) Fe<sup>3+</sup>
E) Co<sup>2+</sup>
Q2) Apply the valence bond theory to predict the electronic structure and hybridization pattern of chromium in the complex ion Cr(NH<sub>3</sub>)<sub>6</sub><sup>3+</sup>.
Q3) What geometry is particularly common for complexes of d<sup>10</sup> metal ions?
Q4) What is the difference between a coordination compound and a complex ion?
Q5) The compound Rh(CO)(H)(PH<sub>3</sub>)<sub>2</sub> forms cis and trans isomers. Use this information to predict the geometry of this complex, and draw the geometric isomers.
Q6) Why is the +2 oxidation state so common among transition elements?
Page 25
Q7) Tetrahedral complexes can exhibit both optical and linkage isomerism. A)True B)False
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Chapter 24: Nuclear Reactions and Their Applications
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Q1) Gamma rays are not deflected by an electric field.
A)True
B)False
Q2) So-called "magic numbers" of particles are thought to convey extra stability to certain nuclei. These magic numbers refer to which of the following particles?
A) protons only
B) electrons only
C) positrons only
D) neutrons only
E) protons and neutrons
Q3) A 9.52 × 10<sup>-5</sup> mol sample of rubidium-86 emits 8.87 × 10<sup>16</sup>
\(\alpha\) particles in one hour. What is the half-life of rubidium-86?
A) 2.23 × 10<sup>-3</sup> h
B) 1.55 × 10<sup>-3</sup> h
C) 448 h
D) 645 h
E) none of the above
Q4) Write a complete, balanced equation to represent the beta decay of thallium-207.
Q5) Write a complete, balanced equation to represent the alpha decay of radon-210.
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