General Chemistry I Exam Solutions - 3071 Verified Questions

General Chemistry I Exam Solutions

Course Introduction

General Chemistry I introduces students to the fundamental principles and concepts of chemistry, laying the groundwork for further study in the sciences. The course covers topics such as atomic structure, chemical bonding, stoichiometry, the periodic table, chemical reactions, thermochemistry, and the properties of gases, liquids, and solids. Through a combination of lectures, laboratory experiments, and problem-solving exercises, students develop critical thinking skills and gain hands-on experience in applying theoretical knowledge to real-world chemical problems. This foundational course is essential for students pursuing majors in science, engineering, health professions, or related fields.

Recommended Textbook Chemistry Structure and Properties 2nd Edition by Nivaldo J. Tro

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Chapter 1: Essentials: Units, Measurements, and Problem Solving

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Sample Questions

Q1) What decimal power does the abbreviation p represent?

A) 1 × 10⁶

B) 1 × 10⁹

C) 1 × 10^-1

D) 1 × 10^-12

E) 1 × 10^-15

Answer: D

Q2) How many µm³ are found in a shape that has a volume of 4.64 × 10^-9 cm³?

A) 4.64 × 10^-5 µm³

B) 4.64 × 10⁸ c = µm³

C) 4.64 × 10³ µm³

D) 4.64 × 10¹⁶ µm³

E) 4.64 × 10^-2 µm³

Answer: C

Q3) Define energy.

Answer: Energy is the capacity to do work.

Q4) Define the law of the conservation of energy.

Answer: Energy is neither created or destroyed.

Page 3

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Chapter 2: Atoms

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Sample Questions

Q1) How many protons (p)and neutrons (n)are in an atom of strontium-90?

A) 38 p, 52 n

B) 38 p, 90 n

C) 52 p, 38 n

D) 90 p, 38 n

Answer: A

Q2) Identify a solid.

A) copper

B) oxygen

C) water

D) nitrogen

E) air

Answer: A

Q3) an atom that has lost an electron is

A) a cation.

B) unlikely to be found in homogeneous mixtures.

C) electrically neutral.

D) likely to behave exactly like the parent atom.

E) an anion.

Answer: A

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Chapter 3: The Quantum Mechanical Model of the Atom

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Sample Questions

Q1) ________ are used to image bones and internal organs.

A) Ultraviolet light

B) Gamma rays

C) Microwaves

D) X-rays

E) Radio waves

Answer: D

Q2) Choose the one set of quantum numbers that contains an error.

A) n = 6, l = 3, m_l=+2

B) n = 3, l = 2, m_l =0

C) n = 4, l = 0, m_l = -3

D) n = 5, l = 4, m_l = -2

E) n = 3, l =1, m_l =-1

Answer: C

Q3) Give an example of a p orbital.

Answer: p_x,p_y,or p_z

Q4) Give an example of a d orbital.

Answer:

d_yz,d_xy,d_xz,d_x2_-y2,or d_z2

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Chapter 4: Periodic Properties of the Elements

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Sample Questions

Q1) Give the ground state electron configuration for Pb.

A) [Xe]6s²6p²

B) [Xe]6s²5d¹⁰6p²

C) [Xe]6s²5f¹⁴6d¹⁰6p²

D) [Xe]6s²4f¹⁴5d¹⁰6p²

E)

[Xe]6s²4f¹⁴5d¹⁰6s²6p²

Q2) Identify the number of valence electrons in Cl^-.

A) 6

B) 7

C) 8

D) 5

E) 4

Q3) What is the chemical symbol for iron?

A) In

B) Ir

C) Fe

D) I

Q4) Give the name of the element whose symbol is Na.

Q5) List the noble gas that has the highest ionization energy.

Q6) Define ionization energy.

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Chapter 5: Molecules and Compounds

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Sample Questions

Q1) What is the chemical formula for strontium hydride?

A) SrH₂

B) SrOH

C) SrOH₂

D) Sr(OH)₂

Q2) What is the chemical formula for barium oxide?

A) BaO

B) BaO₂

C) Ba₂O₂

D) Ba₂O

E) Ba₂O₃

Q3) Combustion analysis of 0.600 g of an unknown compound containing carbon,hydrogen,and oxygen produced 1.043 g of CO₂ and 0.5670 g of H₂O.What is the empirical formula of the compound?

A) C₂H₅O

B) C₂H₅O₂

C) C₂H₁₀O₃

D) C₃H₈O₂

Q4) Describe a structural formula.

Q5) Define empirical formula.

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Chapter 6: Chemical Bonding I

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Sample Questions

Q1) Place the following in order of increasing bond length. C-F C-S C-Cl

A) C-S < C-Cl < C-F

B) C-Cl < C-F < C-S

C) C-F < C-S < C-Cl

D) C-F < C-Cl < C-S

E) C-S < C-F < C-Cl

Q2) Describe the difference between a pure covalent bond and a polar covalent bond.

Q3) Place the following elements in order of increasing electronegativity. Ba S Li

A) Ba < Li < S

B) Li < S < Ba

C) Ba < S < Li

D) S < Ba < Li

E) S < Li < Ba

Q4) How many lone pairs of electrons are on the Br atom in BrF₄^+?

A) 0

B) 1

C) 2

D) 3

Q5) Define formal charge.

Page 8

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Chapter 7: Chemical Bonding Ii

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Sample Questions

Q1) Draw the Lewis structure for the molecule CH₂CHCH₃.How many sigma and pi bonds does it contain?

A) 8 sigma, 1 pi

B) 9 sigma, 0 pi

C) 9 sigma, 1 pi

D) 7 sigma, 2 pi

E) 8 sigma, 2 pi

Q2) Give the electron geometry (eg),molecular geometry (mg),and hybridization for XeF₄.

A) eg = tetrahedral, mg = tetrahedral, sp³

B) eg = trigonal pyramidal, mg = trigonal pyramidal, sp³

C) eg = octahedral, mg = square planar, sp³d²

D) eg = octahedral, mg = octahedral, sp³d²

E) eg = trigonal bipyramidal, mg = seesaw, sp³d

Q3) Describe a sigma bond.

A) side by side overlap of d orbitals

B) end to end overlap of p orbitals

C) s orbital overlapping with the side of a p orbital

D) overlap of two d orbitals

E) p orbital overlapping with an f orbital

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Chapter 8: Chemical Reactions and Chemical Quantities

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Sample Questions

Q1) Carbonic acid can form water and carbon dioxide upon heating.How much carbon dioxide is formed from 12.4 g of carbonic acid?

H₂CO₃ H₂O + CO₂

A) 8.80 g

B) 17.5 g

C) 12.4 g

D) 3.60 g

E) 42.7 g

Q2) How many molecules of HCl are formed when 90.0 g of water reacts according to the following balanced reaction? Assume excess ICl₃. 2 ICl₃ + 3 H₂O ICl + HIO₃ + 5 HCl

A) 5.00 × 10²⁴molecules HCl

B) 3.00 × 10²⁴molecules HCl

C) 9.00 × 10²⁴molecules HCl

D) 6.00 × 10²⁴molecules HCl

E) 5.00 × 10²⁵molecules HCl

Q3) Balance the following equation. ________ C₁₀H₁₂ + ________ O₂ ________ H₂O + ________ CO₂

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Chapter 9: Introduction to Solutions and Aqueous Reactions

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Sample Questions

Q1) Which of the following pairs of aqueous solutions will form a precipitate when mixed?

A) KOH + Li₂S

B) (NH₄)₂SO₄ + LiCl

C) Sr(C₂H₃O₂)₂ + Li₂SO₄

D) KNO₃+ LiOH

E) None of the above solution pairs will produce a precipitate.

Q2) How many grams of NaOH (MW = 40.0)are there in 500.0 mL of a 0.175 M NaOH solution?

Q3) Determine the oxidation state of P in PO₃^3-.

A) +3

B) +6

C) +2

D) 0

E) -3

Q4) Determine the oxidation state of Mn in KMnO₄.

Q5) Describe the difference between complete ionic and net ionic equations.

Q6) How can you tell if a reaction is an oxidation-reduction reaction?

Q7) Define an electrolyte.

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Chapter 10: Thermochemistry

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Sample Questions

Q1) Which of the following processes are endothermic?

A) the reaction associated with the lattice energy of LiCl

B) the reaction associated with the ionization energy of potassium

C) the reaction associated with the heat of formation of CaS

D) the formation of F₂ from its elements in their standard states

E) None of the above are endothermic.

Q2) Give the temperature and pressure for the standard state for a liquid.

Q3) Which of the following is NOT a standard state?

A) For a liquid, it is 25°F.

B) For a solid, it is 25°C.

C) For a solid, it is 1 atm.

D) For a solution, it is 1 M.

E) For a liquid, it is 1 atm.

Q4) Calculate the amount of heat (in kJ)necessary to raise the temperature of 53.8 g benzene by 50.6 K.The specific heat capacity of benzene is 1.05 J/g°C.

A) 1.61 kJ

B) 16.6 kJ

C) 2.59 kJ

D) 2.86 kJ

E) 3.85 kJ

Page 12

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Chapter 11: Gases

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Sample Questions

Q1) How many grams of water are required to produce 5.50 L of hydrogen gas at 25.0°C and 755 mm Hg pressure according to the chemical equation shown below?

BaH₂(s)+ 2 H₂O(l) Ba(OH)₂(aq)+ 2 H₂(g)

A) 2.01 g

B) 4.02 g

C) 4.07 g

D) 8.04 g

Q2) Give the major gas in dry air.

Q3) What is the volume of 9.783 × 10²³ atoms of Kr at 9.25 atm and 512K?

A) 7.38 L

B) 3.69 L

C) 1.85 L

D) 15.4 L

E) 30.8 L

Q4) Which of the following gases has the highest average speed at 400K?

A) N₂

B) O₂ 

C) F₂

D) Cl₂

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Chapter 12: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Define freezing.

A) the phase transition from solid to gas

B) the phase transition from gas to solid

C) the phase transition from gas to liquid

D) the phase transition from liquid to gas

E) the phase transition from liquid to solid

Q2) Choose the pair of substances that are most likely to form a homogeneous solution.

A) K I and Hg

B) Li Cl and C₆H₁₄

C) C₃H₈ and C₂H₅OH

D) F₂ and PF₃

E) NH₃ and CH₃OH

Q3) Which is expected to have the largest dispersion forces?

A) C₃H₈

B) C₁₂H₂₆

C) F₂

D) BeCl₂

Q4) Define boiling point of a liquid.

Q5) Define viscosity.

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Chapter 13: Crystalline Solids and Modern Materials

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Sample Questions

Q1) Identify the type of solid for diamond.

A) metallic atomic solid

B) ionic solid

C) nonbonding atomic solid

D) molecular solid

E) networking atomic solid

Q2) When an X-ray beam with a wavelength of 180 pm strikes the surface of a crystal,it produces a maximum reflection at an angle of 54.0°.If n=1,what is the separation between layers of atoms in the crystal?

A) 151 nm

B) 38.9 mm

C) 111 nm

D) 55.3 nm

E) 83.5 nm

Q3) Identify the copolymer.

A) nylon 6,6

B) polyvinyl chloride

C) polyethylene

D) polypropylene

E) polystyrene

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Chapter 14: Solutions

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Sample Questions

Q1) A solid can be purified through what technique?

A) recrystallization

B) dilution

C) dissolution

D) boiling

E) melting

Q2) Calculate the freezing point of a solution of 40.0 g methyl salicylate,C₇H₆O₂,dissolved in 800.g of benzene,C₆H₆.K_f for benzene is 5.10°C/m and the freezing point is 5.50°C for benzene.

A) - 2.09°C

B) 2.09°C

C) 3.41°C

D) 7.59°C

Q3) Give the reason that antifreeze is added to a car radiator.

A) The freezing point is lowered and the boiling point is elevated.

B) The freezing point is elevated and the boiling point is lowered.

C) The freezing point and the boiling point are elevated.

D) The freezing point and the boiling point are lowered.

E) None of the above.

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Chapter 15: Chemical Kinetics

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Sample Questions

Q1) What is the overall order of the following reaction,given the rate law?

2NO(g)+ H₂(g) N₂(g)+ 2H₂O(g)Rate = k[NO]²[H₂]

A) 1st order

B) 7th order

C) 3rd order

D) 4th order

E) 0th order

Q2) A reaction is found to have an activation energy of 38.0 kJ/mol.If the rate constant for this reaction is 1.60 × 10² M^-1s^-1 at 249 K,what is the rate constant at 436 K?

A) 2.38 × 10⁵M^-1s^-1

B) 1.26 × 10³M^-1s^-1

C) 7.94 × 10⁴ M^-1s^-1

D) 4.20 × 10⁵ M^-1s^-1

E) 3.80 × 10⁴M^-1s^-1

Q3) Define activation energy.

Q4) What is a catalyst and what function does it serve?

Q5) What happens to the concentration of reactants and products during a chemical reaction?

Page 17

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Chapter 16: Chemical Equilibrium

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Q1) Consider the following reaction at equilibrium.What effect will adding more H₂S have on the system?

2 H₂S(g)+ 3 O₂(g) 2 H₂O(g)+ 2 SO₂(g)

A) The reaction will shift to the left.

B) No change will be observed.

C) The equilibrium constant will decrease.

D) The equilibrium constant will increase.

E) The reaction will shift in the direction of products.

Q2) Define Le Chatelier's Principle.

Q3) The reaction below has a K_p value of 41.What is the value of K_c for this reaction at 400 K?

N₂(g)+ 3 H₂(g) 2 NH₃(g)

A) 3.8 × 10^-2

B) 4.4 × 10⁴

C) 26

D) 2.3 × 10^-5

E) 41

Q4) How is the reaction quotient different from an equilibrium constant for a given reaction?

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Chapter 17: Acids and Bases

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Q1) Place the following in order of increasing acid strength. HBrO₂ HBrO₃ HBrO HBrO₄

A) HBrO₂ < HBrO₄ < HBrO < HBrO₃

B) HBrO < HBrO₂< HBrO₃< HBrO₄

C) HBrO₂ < HBrO₃ < HBrO₄< HBrO

D) HBrO₄ < HBrO₂ < HBrO₃< HBrO

E) HBrO < HBrO₄< HBrO₃< HBrO₂

Q2) Determine the pH of a 0.232 M Ca(OH)₂ solution at 25°C.

A) 12.73

B) 13.37

C) 13.67

D) 0.333

E) 0.634

Q3) Which one of the following will form an acidic solution in water?

A) NH₄Cl

B) KF

C) KI

D) KNO₃

E) None of the above solutions will be acidic.

Q4) What is the autoionization of water?

Page 19

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Chapter 18: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Identify the indicator that can be used at the lowest pH.

A) 2,4-dinitrophenol

B) thymol blue

C) crystal violet

D) thymolphthalein

E) methyl red

Q2) The molar solubility of CaF₂ is 2.15 × 10^-4 M in pure water.Calculate the K_sp for CaF₂.

A) 1.63 × 10^-12

B) 8.05 × 10^-9

C) 3.97 × 10^-11

D) 4.47 × 10^-12

E) 5.31 × 10^-10

Q3) Identify the compound that is acid-insoluble.

A) PbCl₂

B) As₂S₃

C) FeS

D) Ca₃(PO₄)₂

E) LiCl

Q4) Define a buffer.

Page 20

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Chapter 19: Free Energy and Thermodynamics

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Sample Questions

Q1) Which of the following is NOT true for G_rxn?

A) If G°_rxn > 0, the reaction is spontaneous in the forward direction.

B) If Q = 1, then G_rxn = G°_rxn.

C) G°_rxn = K

D) If G°_rxn > 0, the reaction is spontaneous in the reverse direction.

E) Under equilibrium conditions, G_rxn = 0.

Q2) Why is heating your home with gas more efficient than heating it with electricity?

Q3) If a reaction has a value of-16.5 kJ/mol for H° and a value of -150 J/mol K for S°,what is the value of the equilibrium constant at 25°C?

A) 9.42 × 10³

B) 1.14 x10^-5

C) 6.52 × 10^-2

D) 8.27

Q4) Define the third law of thermodynamics.

Q5) How is a nonspontaneous process made spontaneous?

Q6) What is "free" energy?

Give a fictitious example.

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Chapter 20: Electrochemistry

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Q1) How many grams of chromium metal are plated out when a constant current of 8.00 A is passed through an aqueous solution containing Cr³⁺ ions for 320.minutes?

A) 27.6 g

B) 49.2 g

C) 82.4 g

D) 248 g

Q2) Why is sugar water not a good conductor of current?

Q3) Predict the species that will be reduced first if the following mixture of molten salts undergoes electrolysis. Na⁺,Ca²⁺,Cl ,Br ,F

A) Na

B) Cl

C) Ca²⁺

D) Br

E) F

Q4) What is the difference between a voltaic cell and an electrolytic cell?

Q5) Give an example of an inert electrode.

Q6) Explain the significance of the standard hydrogen electrode (SHE)in the tabulation of standard reduction potentials of other species.

Page 22

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Chapter 21: Radioactivity and Nuclear Chemistry

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Q1) Atoms with Z > ________ are radioactive and decay in one or more steps involving mostly alpha and beta decay.

A) 60

B) 110

C) 83

D) 160

E) 30

Q2) Calculate the mass defect in Fe-56 if the mass of an Fe-56 nucleus is 55.921 amu.The mass of a proton is 1.00728 amu and the mass of a neutron is 1.008665 amu.

A) 0.528 amu

B) 3.507 amu

C) 0.564 amu

D) 1.056 amu

E) 0.079 amu

Q3) How does a dosimeter measure exposure to radioactivity?

Q4) Why is an alpha emitter much more harmful if is ingested than when applied to the skin?

Q5) Explain the concept of "magic numbers."

Q6) Describe what is meant by the term "valley of stability"?

Page 23

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Chapter 22: Organic Chemistry

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Q1) Write the balanced chemical equation that represents the addition of Cl₂ to CH₃CH=CH₂.

A) CH₃CH=CH₂+Cl₂ CH₃CHClCH₃ + HCl

B) CH₃CH=CH₂+Cl₂ CH₃CCl=CHCl₊H₂

C) CH₃CH=CH₂+Cl₂ CH₃CHClCH₂Cl

D) CH₃CH=CH₂+2 Cl₂ CH₃CHCl₂ + CH₂Cl₂

E) CH₃CH=CH₂+2 Cl₂ CH₃CCl₂CHCl₂+ H₂

Q2) Which of the following alkane names is correct?

A) 2,3-dibutylpropane

B) 3-ethyl-4- methylbutane

C) 4-ethyl-5- butyl heptene

D) 3,3-dibutylpentane

E) None of the above names are correct.

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Chapter 23: Transition Metals and Coordination Compounds

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Q1) Choose the bidentate ligand from the substances below.

A) oxalate ion

B) EDTA

C) cyanide ion

D) water

E) carbon monoxide

Q2) Why is +2 a common oxidation state for transition elements?

Q3) Choose the electron configuration for Mo⁵⁺.

A) [Kr]5s¹4d³

B) [Kr]5s²4d²

C) [Kr]5s²4d¹

D) [Kr]4d¹

E) [Kr]5s²4d⁶

Q4) Identify the transition metal that is used in fat and carbohydrate synthesis in the human body.

A) sodium

B) zinc

C) copper

D) manganese

E) chromium

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