Honors General Chemistry Solved Exam Questions - 2309 Verified Questions

Honors General Chemistry

Solved Exam Questions

Course Introduction

Honors General Chemistry is an advanced, fast-paced course designed for students with a strong background and keen interest in chemical sciences. This course covers fundamental principles of chemistry, including atomic and molecular structure, stoichiometry, thermochemistry, electronic structure, periodic trends, bonding, states of matter, kinetics, equilibrium, and acid-base chemistry. Emphasizing conceptual understanding and problem-solving skills, the curriculum integrates laboratory experiments, theoretical analysis, and practical applications. Students are encouraged to think critically and apply chemical principles to real-world scenarios, preparing them for further studies in science, medicine, or engineering.

Recommended Textbook

Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg

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24 Chapters

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Chapter 1: Keys to the Study of Chemistry

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Sample Questions

Q1) Which of the following is not an SI base unit?

A) meter

B) ampere

C) second

D) gram

E) kelvin

Answer: D

Q2) When applying the scientific method, a model or theory should be based on experimental data.

A)True

B)False

Answer: True

Q3) Which of the following is an extensive property of oxygen?

A) boiling point

B) temperature

C) average kinetic energy of molecules

D) density

E) mass

Answer: E

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Chapter 2: The Components of Matter

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Sample Questions

Q1) All neutral atoms of tin have 50 protons and 50 electrons.

A)True

B)False

Answer: True

Q2) The molecular formula of a compound provides more information than its structural formula.

A)True

B)False

Answer: False

Q3) a. Give the names of the following ions:

(i) NH₄⁺ (ii) SO₃^2-

b. Write down the formulas of the following ions: (i) aluminum (ii) carbonate

Answer: a. (i) ammonium (ii) sulfite

b. (i) Al³⁺ (ii) CO₃^2-

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Page 4

Chapter 3: Stoichiometry of Formulas and Equations

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Sample Questions

Q1) Propane, C₃H₈, is commonly provided as a bottled gas for use as a fuel. In 0.200 mol of propane

a. what is the mass of propane?

b. what mass of carbon is present?

c. how many molecules of C₃H₈ are present?

d. how many hydrogen atoms are present?

Answer: a. 8.82 g

b. 7.21 g

c. 1.20 × 10²³ C₃H₈ molecules

d. 9.64 × 10²³ H atoms

Q2) In 0.20 mole of phosphoric acid, H₃PO₄

a. how many H atoms are there?

b. what is the total number of atoms?

c. how many moles of O atoms are there?

Answer: a. 3.61 × 10²³ H atoms

b. 9.64 × 10²³ atoms

c. 4.82 × 10²³ O atoms

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Chapter 4: The Major Classes of Chemical Reactions

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Sample Questions

Q1) A combination reaction may also be a displacement reaction.

A)True

B)False

Q2) Calculate the oxidation number of the chlorine in perchloric acid, HClO₄, a strong oxidizing agent.

A) -1

B) +4

C) +5

D) +7

E) None of the above is the correct oxidation number.

Q3) Which one of the following ionic compounds is insoluble in water?

A) Na₃PO₄

B) AgNO₃

C) NaCl

D) CaCO₃

E) MgCl₂

Q4) You are provided with a 250 mL volumetric flask, deionized water, and solid NaOH. How much NaOH should be weighed out in order to make 250. mL of 0.100 M solution?

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Page 6

Chapter 5: Gases and the Kinetic-Molecular Theory

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Sample Questions

Q1) Starting from the Ideal Gas Equation, derive an equation corresponding to Charles's Law, stating all important assumptions or conditions.

Q2) "The total pressure in a mixture of unreacting gases is equal to the sum of the partial pressures of the individual gases" is a statement of __________________ Law.

A) Charles's B) Graham's C) Boyle's D) Avogadro's E) Dalton's

Q3) A 255-mL gas sample weighing 0.292 g is at 52810 Pa and 127°C.

a. How many moles of gas are present?

b. What is the molar mass of the gas?

Q4) Briefly state the conditions corresponding to STP (standard temperature and pressure).

Q5) Aluminum metal shavings (10.0 g) are placed in 100. mL of 6.00 M hydrochloric acid. What is the maximum volume of hydrogen, measured at STP, which can be produced? 2Al(s) + 6HCl(aq) \(\to\) 2AlCl₃(aq) + 3H₂(g)

Q6) State Avogadro's Law.

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Chapter 6: Thermochemistry: Energy Flow and Chemical Change

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Sample Questions

Q1) Starting from equations relating pressure to force and force to work, derive the relationship w = - P\(\Delta\)V, explaining the steps in your argument.

Q2) For which one of the following reactions will \(\Delta\)H be approximately (or exactly) equal to \(\Delta\)E?

A) H₂(g) + Br₂(g) \(\to\) 2HBr(g)

B) H₂O(l) \(\to\) H₂O(g)

C) CaCO₃(s) \(\to\) CaO(s) + CO₂(g)

D) 2H(g) + O(g) \(\to\) H₂O(l)

E) CH₄(g) + 2O₂(g) \(\to\) CO₂(g) + 2H₂O(l)

Q3) Which one of the following relationships is always correct?

A) potential energy + kinetic energy = constant

B) E = q + w

C) .\(\Delta\)E = \(\Delta\)H - P\(\Delta\)V

D) H = E + PV

E) .\(\Delta\)H = q_v

Q4) 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) The following combinations of quantum numbers are not allowed.

Q2) Who was the first scientist to propose that the atom had a dense nucleus which occupied only a small fraction of the volume of the atom?

A) Planck

B) Bohr

C) Rydberg

D) Rutherford

E) Thomson

Q3) In not more than three lines for each answer, briefly outline one important scientific contribution of each of the following:

a. Planck

b. de Broglie

c. Heisenberg

Q4) a. Use Bohr's equation to calculate how much energy (in J) is needed to promote an electron from the H-atom ground state to the n = 4 level.

b. If a photon provides the energy in (a), what is its wavelength in nm?

Q5) What is the minimum uncertainty in the position of a neutron if the uncertainty in its speed is 0.0250 m/s?

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Chapter 8: Electron Configuration and Chemical Periodicity

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Sample Questions

Q1) Define what is meant by ionization energy, and write a balanced chemical equation to represent the relevant process for element X.

Q2) Which of the following elements is paramagnetic?

A) Kr

B) Zn

C) Sr

D) V

E) Ar

Q3) Select the correct electron configuration for Te (Z = 52).

A) [Kr]5s²5p⁶4d⁸

B) [Kr]5s²5d¹⁰5p⁴

C) [Kr]5s²4d¹⁰5p⁶

D) [Kr]5s²4f¹⁴

E) [Kr]5s²4d¹⁰5p⁴

Q4) Moseley's measurements of nuclear charges of the elements provided the basis for arranging the elements of the periodic table in order of increasing atomic number.

A)True

B)False

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Chapter 9: Models of Chemical Bonding

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Sample Questions

Q1) The melting points of metals are only moderately high because A) metallic bonding is weak.

B) metals have fewer bonding electrons than non-metals.

C) metals also have relatively low boiling points.

D) the melting process does not break the metallic bonds.

E) metals prefer to be bonded to non-metals.

Q2) Using the bond energies provided below, calculate \(\Delta\)H° for the reaction CH₄(g) + 4Cl₂(g) \(\to\) CCl₄(g) + 4HCl(g)

Bond energies: C-H = 413 kJ/mol, Cl-Cl = 243 kJ/mol, C-Cl = 339 kJ/mol, H-Cl = 427 kJ/mol

A) 1422 kJ

B) 440 kJ

C) 110 kJ

D) -110 kJ

E) - 440 kJ

Q3) The stronger the bonds in a fuel, the more energy it will yield.

A)True

B)False

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Chapter 10: The Shapes of Molecules

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Sample Questions

Q1) Which of the following molecules has a net dipole moment?

A) BeCl₂

B) SF₂

C) KrF₂

D) CO₂

E) CCl₄

Q2) For the chlorate ion, ClO₃^-, draw two different valid Lewis structures, as follows:

a. a structure in which the octet rule is obeyed

b. a structure in which formal charges are minimized

Q3) What is the molecular shape of BeH₂as predicted by the VSEPR theory?

A) linear

B) bent

C) angular

D) trigonal

E) none of the above

Q4) Name and outline the concept which is introduced when more than one valid Lewis structure can be drawn for a given molecule or ion. Use appropriate diagrams of the formate ion (HCO₂^-, carbon is the central atom) to illustrate.

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Chapter 11: Theories of Covalent Bonding

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Sample Questions

Q1) The angles between sp² hybrid orbitals are 109.5°.

A)True

B)False

Q2) According to valence bond theory, overlap of bonding orbitals of atoms will weaken a bond, due to electron-electron repulsion.

A)True

B)False

Q3) Which one, if any, of the following statements about the MO treatment of the bonding in benzene is correct?

A) MO theory uses resonance to describe the delocalized nature of the bonding in benzene.

B) The lowest energy pi-bonding orbital in benzene can hold up to 6 electrons.

C) MO theory predicts that benzene should be paramagnetic.

D) The lowest energy pi-bonding MO in benzene has two hexagonal lobes, above and below the plane of the carbon atoms.

E) None of the above statements is correct.

Q4) Overlap of two sp² hybrid orbitals produces a \(\pi\) bond.

A)True

B)False

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Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes

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Sample Questions

Q1) A temperature increase causes __________________ in the conductivity of a conductor.

A) a decrease

B) an increase

C) an increase or decrease (depending on the conductor)

D) a modulation

E) no change

Q2) Mercury melts at -39°C and boils at 357°C. Draw a diagram of the heating curve of mercury. Label all lines and axes, and clearly indicate the melting and boiling points on your diagram.

Q3) A metal with a body-centered cubic lattice will have ______ atom(s) per unit cell.

A) 1

B) 2

C) 3

D) 4

E) 9

Q4) How do the electrical properties of semiconductors differ from those of metals?

Q5) Use molecular orbital band diagrams to explain why metals are good conductors but semiconductors are not.

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Chapter 13: The Properties of Mixtures: Solutions and Colloids

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Sample Questions

Q1) Based only on the relative lattice energies of the compounds below, which one would be expected to have the lowest solubility in water?

A) NaBr

B) CaS

C) NaOH

D) KI

E) CsCl

Q2) What is the mole fraction of Ar in a mixture containing 10.1 g of Ne, 79.9 g of Ar, and 83.8 g of Kr?

A) 0.40

B) 0.25

C) 0.20

D) 0.14

E) 0.058

Q3) A 7.112 M solution of sulfuric acid (H₂SO₄) in water has a density of 1.395 g/mL. What is its molality?

Q4) List five important intermolecular forces that operate in globular proteins, and explain what parts of the molecule are involved in these forces.

Page 15

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Chapter 14: Periodic Patterns in the Main Group Elements:

Bonding, Structure, and Reactivity

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Sample Questions

Q1) Phosphoric acid (H₃PO₄) is a strong acid.

A)True

B)False

Q2) Select the element with the lowest first ionization energy.

A) Se

B) S

C) Sn

D) Sr

E) H

Q3) Sulfuric acid is produced when sulfur dioxide dissolves in water.

A)True

B)False

Q4) Which element forms compounds which are involved in smog and acid rain?

A) carbon

B) fluorine

C) chlorine

D) boron

E) nitrogen

Page 16

Q5) Name the three different classes (types) of hydride, and list some of their important characteristics.

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Chapter 15: Organic Compounds and the Atomic Properties of

Carbon

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Sample Questions

Q1) Which of the following features do cyclohexene and 3-methyl-1-pentyne have in common?

A) same physical properties

B) same chemical properties

C) same boiling point

D) same molecular weight

E) same number of double bonds

Q2) Both alkenes and alkynes exhibit geometric (cis/trans) isomerism.

A)True

B)False

Q3) Excluding cyclic compounds, how many possible isomers exist for C₄H₈?

A) 2

B) 4

C) 5

D) 6

E) 7

Q4) All ketone molecules are capable of hydrogen bonding to other ketone molecules.

A)True

B)False

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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions

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Sample Questions

Q1) Briefly outline the key arguments in the collision theory of reaction rates for the elementary reaction

C + D \(\to\) products

Show that this theory predicts a second-order rate law, and how it predicts the form of the rate constant k.

Q2) Chlorine atoms act as heterogeneous catalysts in the destruction of ozone in the stratosphere.

A)True

B)False

Q3) The rate law for the reaction 3A \(\to\) 2B is rate = k[A] with a rate constant of 0.0447 hr^-1. What is the half-life of the reaction?

A) 0.0224 hr

B) 0.0645 hr

C) 15.5 hr

D) 22.4 hr

E) 44.7 hr

Q4) Briefly list the features/properties common to all catalysts and how they work. Draw a labeled reaction energy diagram as part of your answer.

Page 18

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Chapter 17: Equilibrium: the Extent of Chemical Reactions

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Sample Questions

Q1) If all the reactants and products in an equilibrium reaction are in the gas phase, then K_p = K_c.

A)True

B)False

Q2) A chemical reaction has an equilibrium constant of 2 × 10⁶. If this reaction is at equilibrium, select the one correct conclusion that can be made about the reaction.

A) The forward and back reactions have stopped.

B) The limiting reactant has been used up.

C) The forward and reverse rate constants are equal.

D) The forward and reverse reaction rates are equal.

E) None of the above conclusions is correct.

Q3) In a chemical reaction, if the starting concentrations of reactants are increased, then the equilibrium constant K_c will also increase.

A)True

B)False

Q4) For some gas-phase reactions, K_p = K_c.

A)True

B)False

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Chapter 18: Acid-Base Equilibria

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Sample Questions

Q1) A solution of sodium acetate (CH₃COONa) in water is weakly basic.

A)True

B)False

Q2) Ammonium chloride is used as an electrolyte in dry cells. Which of the following statements about a 0.10 M solution of NH₄Cl, is correct?

A) The solution is weakly basic.

B) The solution is strongly basic.

C) The solution is neutral.

D) The solution is acidic.

E) The values for K_a and K_b for the species in solution must be known before a prediction can be made.

Q3) The ammonium ion, NH₄⁺, is a weak acid.

A)True B)False

Q4) All strong acids have weak conjugate bases.

A)True B)False

Q5) Describe what is meant by the "leveling effect". Use a real acid as an example, and write an appropriate equation.

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Chapter 19: Ionic Equilibria in Aqueous Systems

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Sample Questions

Q1) Which one of the following aqueous solutions, when mixed with an equal volume of 0.10 mol L^-1 aqueous NH₃, will produce a buffer solution?

A) 0.10 mol L^-1 HCl

B) 0.20 mol L^-1 HCl

C) 0.10 mol L^-1 CH₃COOH

D) 0.050 mol L^-1 NaOH

E) 0.20 mol L^-1 NH₄Cl

Q2) Assuming that the total volume does not change after 0.200 g of KCl is added to 1.0 L of a saturated aqueous solution of AgCl, calculate the number of moles of Ag⁺ ion in the solution after equilibrium has been reestablished. For AgCl, K_sp = 1.8 × 10^{- 10}.

A) 1.8 × 10^-10 mol Ag⁺

B) 9.0 × 10^-10mol Ag⁺

C) 9.0 × 10^-9mol Ag⁺

D) 6.7 × 10^-8mol Ag⁺

E) 1.3 × 10^-5mol Ag⁺

Q3) 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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Sample Questions

Q1) Consider the reaction 11ec6d53_97e1_49fd_9cd1_030bf8888a28_TB5833_

If the concentrations of the Cu⁺ and I^- ions in equilibrium at 298 K are both equal to 1.03 × 10^-6 M, what is the value of \(\Delta\)G° for the reaction?

A) -68 kJ

B) 68 kJ

C) -30. kJ

D) 30 kJ

E) 34 kJ

Q2) For a process with \(\Delta\)S < 0, which one of the following statements is correct?

A) The process will definitely be spontaneous if \(\Delta\)H < 0.

B) The process will be definitely be spontaneous if \(\Delta\)H < T\(\Delta\)S.

C) The process can never be spontaneous.

D) The process will definitely be spontaneous, regardless of \(\Delta\)H.

E) The process will definitely be spontaneous if \(\Delta\)S_surr > 0.

Q3) For the reaction of xenon and fluorine gases to form solid XeF₄, \(\Delta\)H° = -251 kJ and \(\Delta\)G° = -121 kJ at 25°C. Calculate \(\Delta\)S° for the reaction.

Page 22

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Chapter 21: Electrochemistry: Chemical Change and Electrical Work

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Q1) Electrolytic cells utilize electrical energy to drive non-spontaneous redox reactions.

A)True

B)False

Q2) A voltaic cell consists of an Au/Au³⁺ electrode (E° = 1.50 V) and a Cu/Cu²⁺ electrode (E° = 0.34 V). Calculate [Au³⁺] if [Cu²⁺] = 1.20 M and E_cell = 1.13 V at 25°C.

A) 0.001 M

B) 0.002 M

C) 0.01 M

D) 0.02 M

E) 0.04 M

Q3) In the shorthand notation for cells, a double vertical line is used to separate the reduced and oxidized forms of a redox couple.

A)True

B)False

Q4) Electrons are produced at the cathode of a voltaic cell.

A)True

B)False

Page 23

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Chapter 22: The Elements in Nature and Industry

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Sample Questions

Q1) Describe, and give reasons for, the fundamental bonding differences between transition metal sulfides and transition metal oxides.

Q2) What gas is produced during the Hall-Héroult process for production of aluminum?

A) chlorine, Cl₂

B) oxygen, O₂

C) hydrogen, H₂

D) ammonia, NH₃

E) carbon dioxide, CO₂

Q3) Plants extract phosphate from the soil

A) by converting it to the dihydrogen phosphate ion by addition of acid to the soil near roots.

B) by converting it to phosphoric acid.

C) by osmosis.

D) by leaching.

E) by fixation.

Q4) What gas is produced at the anode in the Downs cell, in which molten NaCl is electrolyzed?

Q5) Give two important reasons why there is so much more Na⁺ in the oceans than K⁺, despite their similar abundances in the earth's crust.

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Chapter 23: The Transition Elements and Their Coordination Compounds

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Q1) What geometry is particularly common for complexes of d¹⁰ metal ions?

Q2) A certain transition element has the stable oxidation states of +2, +3, +4, +5, and +6. In which state will the element be most likely to form a covalent bond with chlorine?

A) +2

B) +3

C) +4

D) +5

E) +6

Q3) Which of the following ions could exist in only the high-spin state in an octahedral complex?

A) Cr²⁺

B) Mn⁴⁺

C) Fe³⁺

D) Co³⁺

E) Ni²⁺

Q4) Valence Bond theory rationalizes octahedral geometry by assuming a d²sp³ hybridization pattern.

A)True

B)False

Page 25

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Chapter 24: Nuclear Reactions and Their Applications

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Q1) Bombardment of uranium-238 nuclei by carbon-12 nuclei produces californium-246 and neutrons. Write a complete, balanced equation for this nuclear process.

Q2) After 4 half-lives, the fraction of a radioactive isotope which still remains is approximately one-eighth.

A)True

B)False

Q3) Gamma rays are not deflected by an electric field.

A)True

B)False

Q4) A certain isotope has a specific activity of 7.29 × 10^-4Ci/g. How many \(\alpha\) particles will a 75.0 mg sample emit in one hour?

A) 9.99 × 10⁴

B) 2.02 × 10⁶

C) 7.28 × 10⁹

D) 1.29 × 10¹²

E) none of the above

Q5) Briefly, explain the relationship between the rad and the rem as units of radiation dosage.

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Page 26