Principles of Chemistry Exam Questions - 2556 Verified Questions

Course Introduction

Principles of Chemistry

Exam Questions

Principles of Chemistry provides a foundational understanding of the chemical concepts and principles that govern matter and its transformations. This course covers topics such as atomic and molecular structure, chemical bonding, stoichiometry, thermodynamics, chemical equilibria, acids and bases, and basic kinetics. Emphasis is placed on developing problem-solving skills, understanding laboratory techniques, and applying theoretical concepts to real-world scenarios. Through lectures, laboratory experiments, and interactive activities, students gain the essential knowledge required for advanced study in chemistry and related scientific fields.

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General Chemistry The Essential Concepts 7th Edition by Raymond Chang

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22 Chapters

2556 Verified Questions

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Page 2

Chapter 1: Introduction

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Sample Questions

Q1) Classify the following as either a physical or a chemical property: Ice melts at 0^oC.

Answer: Physical

Q2) One of the common intravenous fluids, called physiological saline, is a homogeneous mixture of NaCl in water. In this mixture, 0.89% of the mass is contributed by the NaCl. What mass of NaCl is found in 450. mL of physiological saline? (Given: density of physiological saline = 1.005 g/cm³)

A)2.0 g

B)4.0 g

C)5.1 g

D)508 g

E)400 g

Answer: B

Q3) Give an example of an intensive property.

Answer: (Answers will vary.)Temperature, density, melting point, boiling point

Q4) An automobile engine has a piston displacement of 1,600 cm³. Express this volume in cubic inches. (1 in = 2.54 cm)

Answer: 98 in³

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Page 3

Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) Give the formula for carbon disulfide.

Answer: CS₂

Q2) Which is the correct formula for copper(II)phosphate?

A)Cu₂PO₄

B)Cu₃(PO₄)₂

C)Cu₂PO₃

D)Cu(PO₄)₂

E)Cu(PO₃)₂

Answer: B

Q3) When a beam of alpha particles passes between two electrically charged plates, the beam is deflected toward the positive plate.

A)True

B)False

Answer: False

Q4) In a cathode ray tube

A)electrons pass from the anode to the cathode.

B)electrons pass from the cathode to the anode.

C)protons pass from the anode to the cathode.

D)protons pass from the cathode to the anode.

Answer: B

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Chapter 3: Stoichiometry

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Sample Questions

Q1) What is the minimum mass of sulfur dioxide necessary to produce 200. g of sulfuric acid in the following reaction?

2SO₂ + O₂ + 2H₂O \(\rarr\) 2H₂SO₄

Answer: 131 g

Q2) Balance the following chemical equation:

C + Fe₂O₃ \(\rarr\) Fe + CO

Answer: 3C + Fe₂O₃ \(\rarr\) 2Fe + 3CO

Q3) A 1.375 g sample of mannitol, a sugar found in seaweed, is burned completely in oxygen to give 1.993 g of carbon dioxide and 0.9519 g of water. The empirical formula of mannitol is A)CHO.

B)CH₇O₃.

C)C₃H₂O.

D)C₃H₇O₃.

E)CH₂O.

Answer: D

Q4) Calculate the molecular mass, in g/mol, of C₆H₁₂O₆.

Answer: 180.2 g/mol

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Chapter 4: Reactions in Aqueous Solution

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Sample Questions

Q1) Batteries in our cars generate electricity by the following chemical reaction. Pb + PbO₂ + 2H₂SO₄ \(\rarr\) 2PbSO₄ + 2H₂O

Which substance is oxidized in this process?

Q2) Write balanced molecular and net ionic equations for the reaction that would occur between Al(s)and Co(NO₃)₂(aq). Be sure to include the correct states in your final equations. If no reaction is expected, write "no reaction."

Q3) What element is oxidized in the following chemical reaction? NiO₂ + Cd + 2H₂O \(\rarr\) Ni(OH)₂ + Cd(OH)₂

A)Ni

B)Cd

C)O

D)H

E)This is not a redox reaction.

Q4) During a titration the following data were collected. A 10. mL portion of an unknown monoprotic acid solution was titrated with 1.0 M NaOH; 40. mL of the base were required to neutralize the sample. How many moles of acid are present in 2.0 liters of this unknown solution?

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Chapter 5: Gases

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Sample Questions

Q1) Calculate the number of moles of gas contained in a 10.0 L tank at 22°C and 105 atm. (R = 0.08206 L.atm/K.mol)

A)1.71 * 10^-3 mol

B)0.0231 mol

C)1.03 mol

D)43.4 mol

E)582 mol

Q2) On a spring morning (20°C)you fill your tires to a pressure of 2.25 atmospheres. As you ride along, the tire heats up to 45°C from the friction on the road. What is the pressure in your tires now?

Q3) What volume of H₂ is formed at STP when 6.0 g of Al is treated with excess NaOH?

2NaOH + 2Al + 6H₂O \(\rarr\) 2NaAl(OH)₄ + 3H₂(g)

Q4) Gasoline (which can be considered to be octane, C₈H₁₈)burns in oxygen to produce carbon dioxide and water. What volume of carbon dioxide at STP is generated as a result of the combustion of 1.0 gal of gasoline?

(The density of gasoline is 0.81 g/mL. 1 gal = 3.78 L)

Q5) What is standard temperature and standard pressure?

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Chapter 6: Energy Relationships in Chemical Reactions

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Sample Questions

Q1) According to the first law of thermodynamics:

A)Energy is neither lost nor gained in any energy transformations.

B)Perpetual motion is possible.

C)Energy is conserved in quality but not in quantity.

D)Energy is being created as time passes. We have more energy in the universe now than when time began.

Q2) Calculate the amount of heat necessary to raise the temperature of 12.0 g of water from 15.4°C to 93.0°C. The specific heat of water = 4.18 J/g·°C.

A)0.027 J

B)324 J

C)389 J

D)931 J

E)3,890 J

Q3) The heat of combustion of propane, C₃H₈, 2220 kJ/mol. The specific heat of copper is 0.385 J/g.°C. How many grams of propane must be burned to raise the temperature of a 10.0 kg block of copper from 25.0°C to 65.0°C, assuming none of the heat is lost to the surroundings?

Q4) What is the standard enthalpy of formation of H₂(g)at 25°C?

Q5) Define specific heat.

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Chapter 7: The Electronic Structure of Atoms

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Sample Questions

Q1) A neon atom in its ground state will be diamagnetic.

A)True

B)False

Q2) When the electron in a hydrogen atom falls from its first excited energy level to the ground state energy level, a photon with wavelength \(\lambda\) is emitted. A proton having this same wavelength would have a velocity of

A)3.87 m/s.

B)5990 m/s.

C)1.21 * 10^-7 m/s.

D)3.26 m/s.

E)5.99 m/s.

Q3) What is the ground-state electron configuration for chlorine?

Q4) The colors of the visible spectrum are blue, green, orange, red, violet, and yellow. Of these colors, ______ has the least energy.

Q5) The frequency of the emitted light from a cesium atom is an intensive property.

A)True

B)False

Q6) What is the total number of electrons possible in the 6s orbital?

Q7) What is the electron configuration of calcium?

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Chapter 8: The Periodic Table

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Sample Questions

Q1) If the radius of atom X is greater than the radius of atom Y, then it is also likely that

A)X has a larger electron affinity than Y does.

B)X has a larger effective nuclear charge than Y does.

C)X has greater metallic character than Y does.

D)X has a larger first ionization energy than Y does.

E)X is a poorer conductor of electricity than Y when in the solid state.

Q2) Write the ground-state electron configuration for I^-.

Q3) Which pair of elements from different groups resemble each other the most in their chemical properties?

A)Be and B

B)Al and Si

C)Li and Be

D)Al and Be

E)Be and C

Q4) Write the ground-state electron configuration for Ni²⁺.

Q5) Which species has the greater radius, an I^- ion or an I atom? Briefly explain your choice of answer.

Q6) Write the ground-state electron configuration for Br^-.

Q7) Write the ground-state electron configuration for Ca²⁺.

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Chapter 9: Chemical Bonding I: the Covalent Bond

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Sample Questions

Q1) How many covalent bonds will be drawn to phosphorous in PO₄^3- for the dot structure that expands the octet to minimize formal charge and if necessary places negative formal charges on the most electronegative atom(s)?

A)4

B)5

C)6

D)7

E)8

Q2) The electron dot structure for AsCl₃ shows

A)a total of 84 electron dots.

B)three single bonds and 10 lone pairs.

C)two single bonds, one double bond, and 9 lone pairs.

D)one single bond, two double bonds, and 8 lone pairs.

E)three single bonds and one lone pair.

Q3) The Si - Cl bond has less ionic character than the C - Cl bond.

A)True

B)False

Q4) Write the chemical equation for which the enthalpy of reaction is the lattice energy of KCl(s).

Page 11

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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization

of Atomic Orbitals

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Sample Questions

Q1) The hybridization of B in the BF₃ molecule is sp³.

A)True

B)False

Q2) What is the hybridization of As in the AsF₄^- ion?

A)sp

B)sp²

C)sp³

D)sp³d

E)sp³d²

Q3) Consider the species N₂^-, N₂, and N₂⁺. Which of these species will be paramagnetic?

A)N₂ and N₂^-

B)N₂⁺ and N₂

C)N₂⁺ and N₂^-

D)only N₂^-

E)none are paramagnetic

Q4) Indicate the number of \(\pi\)-bonds in C₂H₆.

Q5) Indicate the number of \(\pi\)-bonds in C₂H₄.

Page 12

Q6) According to the VSEPR theory, will the molecule PF₅ be polar or nonpolar?

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Chapter 11: Introduction to Organic Chemistry

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Sample Questions

Q1) Which of these molecules is unsaturated?

A)CH₄

B)C₂H₆

C)C₄H₆

D)C₅H₁₂

E)C₆H₁₄

Q2) How many structural isomers are there of C₄H₁₀?

A)4

B)6

C)2

D)8

E)10

Q3) The reaction of ethylene and water yields

A)an aldehyde.

B)an ester.

C)an alcohol.

D)an ether.

E)an organic acid.

Q4) A compound with the formula C₆H₁₂ may or may not be a saturated hydrocarbon. Explain.

Page 14

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Chapter 12: Intermolecular Forces and Liquids and Solids

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Sample Questions

Q1) Indicate all the types of intermolecular forces of attraction in HCl(g).

Q2) Indicate all the types of intermolecular forces of attraction in F₂(l).

Q3) Methane has a heat of fusion of 0.84 kJ/mol and a heat of vaporization of 9.2 kJ/mol. Estimate the value for the heat of sublimation.

Q4) Identify the dominant (strongest)type of intermolecular force present in Cl₂(l).

Q5) For which of the following species are the intermolecular interactions entirely due to dispersion forces?

A)C₂H₆

B)CH₃OCH₃

C)NO₂

D)H₂S

E)CaNO₃

Q6) Boron nitride, BN₃, melts at approximately at 3,000°C under high pressure. This material is almost as hard as diamond. What kind of crystal is this?

Q7) Indicate all the types of intermolecular forces of attraction in CF₄(g).

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Page 15

Chapter 13: Physical Properties of Solutions

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Sample Questions

Q1) Plasma is the fluid portion of blood. The concentration of acetylsalicylic acid (aspirin, C₉H₈O₄, molar mass = 180. g/mol)in your plasma is found to be 2.99 * 10^-4 M after you take two tablets of aspirin. If the volume of your plasma is 5.85 L, what is the molality, m, of aspirin in your blood? (Assume density of plasma is 1.00 g/mL.)

Q2) A saturated solution

A)contains more solute than solvent.

B)contains more solvent than solute.

C)contains equal moles of solute and solvent.

D)contains the maximum amount of solute that will dissolve in that solvent at that temperature.

E)contains a solvent with only sigma bonds and no pi bonds (i.e. only single bonds, with no double or triple bonds).

Q3) Which of the following has the greater molal concentration (molality)?

A)1.0 m KNO₃

B)1.0 M KNO₃

C)Both have same molality.

Q4) Calculate the mass of solute in the following solution: 50.0 mL of 0.0300 M C₁₂H₂₂O₁₁.

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The rate law for the reaction 2NO₂ + O₃ \(\rarr\) N₂O₅ + O₂ is rate = k[NO₂][O₃]. Which one of the following mechanisms is consistent with this rate law?

A)NO₂+ NO₂ \(\rarr\) N₂O₄ (fast) N₂O₄ + O₃ \(\rarr\)N₂O₅ + O₂ (slow)

B)NO₂ + O₃ \(\rarr\) NO₅ (fast) NO₅ + NO₅ \(\rarr\)N₂O₅ + ⁵/₂O₂ (slow)

C)NO₂ + O₃ \(\rarr\) NO₃ + O₂ (slow) NO₃ + NO₂ \(\rarr\) N₂O₅ (fast)

D)NO₂ + NO₂ \(\rarr\) N₂O₂ + O₂ (slow) N₂O₂ + O₃ \(\rarr\) N₂O₅ (fast)

Q2) An experimental drug, D, is known to decompose in the blood stream. Tripling the concentration of the drug increases the decomposition rate by a factor of nine. Write the rate law for decomposition of D and give the units of the rate constant.

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) For the common allotropes of carbon (graphite and diamond), C(gr)\(\rarr\)C(dia)with equilibrium constant K = 0.32. The molar volumes of graphite and diamond are, respectively, 5.30 cm³/mol and 3.42 cm³/mol; \(\Delta\)H_f of diamond is 1.90 kJ/mol. This data suggests that the formation of diamond is favored at

A)low temperatures and low pressures.

B)high temperatures and low pressures.

C)low temperatures and high pressures.

D)high temperatures and high pressures.

Q2) Equilibrium is established for the reaction 2X(s)+ Y(g)\(\rarr\)2Z(g)at 500K, K_c = 100. Determine the concentration of Z in equilibrium with 0.2 mol X and 0.50 M Y at 500K.

A)3.2 M

B)3.5 M

C)4.5 M

D)7.1 M

E)None of these.

Q3) What conditions are used in the Haber process to enhance the yield of ammonia? Explain why each condition affects the yield in terms of the Le Châtelier principle.

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Chapter 16: Acids and Bases

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Sample Questions

Q1) Al(OH)₃ is an amphoteric hydroxide. Write a balanced ionic equation to show its reaction with HNO₃.

Q2) Which one of these statements about strong acids is true?

A)All strong acids have H atoms bonded to electronegative oxygen atoms.

B)Strong acids are 100% ionized in water.

C)The conjugate base of a strong acid is itself a strong base.

D)Strong acids are very concentrated acids.

E)Strong acids produce solutions with a higher pH than weak acids.

Q3) Identify the conjugate acid-base pairs in the reaction HSO₄^- + HF \(\rarr\) H₂SO₄ + F^-

One conjugate acid-base pair is _______________; the other acid-base pair is ______________.

Q4) Which of these lists of molecules is arranged in order of increasing acid strength?

A)H₂S < H₂O < H₂Se

B)H₂O < H₂S < H₂Se

C)H₂Se < H₂O < H₂S

D)H₂S < H₂Se < H₂O

E)H₂O < H₂Se < H₂S

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Chapter 17: Acid-Base Equilibria and Solubility Equilibria

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Sample Questions

Q1) Will a precipitate of AgCl form when 0.050 mol NaCl(s)and 0.050 mol AgNO₃(s)are dissolved in 500. mL of 3.0 M NH₃? [K_fforAg(NH₃)₂⁺is 1.5 * 10⁷; K_sp(AgCl)= 1.6 * 10^-10]

Q2) Will a precipitate of magnesium fluoride form when 200. mL of 1.9 * 10^-3 M MgCl₂ are added to 300. mL of 1.4 * 10^-2 M NaF?

[K_sp (MgF₂)= 6.9 * 10^-9]

A)Yes, Q > K_sp

B)No, Q < K_sp

C)No, Q = K_sp

D)Yes, Q < K_sp

Q3) A saturated sodium carbonate solution at 0°C contains 7.1 g of dissolved sodium carbonate per 100. mL of solution. The solubility product constant for sodium carbonate at this temperature is A)1.2.

B)0.30.

C)3.0 * 10^-4.

D)0.90.

E)1.2 * 10^-3.

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Page 20

Chapter 18: Thermodynamics

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Sample Questions

Q1) H₂O₂(l) \(\rarr\) H₂O(l)+ ¹/₂O₂(g)

A)1 < 3 < 2

B)2 < 3 < 1

C)2 < 1 < 3

D)3 < 2 < 1

E)3 < 1 < 2

Q2) Find the temperature at which K_p = 4.00 for the reaction N₂O₄(g)\(\rarr\) 2NO₂(g). [Given: at 25°C, for NO₂(g), \(\Delta\)H°_f = 33.85 kJ/mol, S° = 240.46 J/mol·K; for N₂O₄(g), \(\Delta\)H°_f = 9.66 kJ/mol, S° = 304.3 J/mol·K; assume that \(\Delta\)H° and \(\Delta\)S° are independent of temperature.]

A)197 °C

B)56 °C

C)36 °C

D)79 °C

E)476°C

Q3) Predict the sign of \(\Delta\)S for the process N₂(g, 10 atm) \(\rarr\) N₂(g, 1atm).

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Page 21

Chapter 19: Redox Reactions and Electrochemistry

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Q1) Aluminum metal is formed by the electrolysis of Al₂O₃in molten cryolite. How many minutes are required to form 10.0 g of Al using a current of 30 A?

Q2) Will H₂(g)form when Sn is placed in 1.0 M HCl?

Q3) Complete and balance the following redox reaction under acidic conditions: ClO₂^-(aq)\(\rarr\) ClO₂(g)+ Cl^-(aq)

Q4) Complete and balance the following redox equation using the set of smallest whole-numbers coefficients. What is the sum of the coefficients?

HI + HNO₃ \(\rarr\) I₂ + NO (acidic solution)

A)5

B)7

C)14

D)17

E)None of these.

Q5) When a solution of a certain gadolinium salt is electrolyzed with a current of 1.0 A for 2.0 h, 0.025 mol of Gd metal forms. Calculate the charge on the gadolinium ion in the salt.

Q6) How many moles of silver metal are produced in 1.2 hours by the electrolysis of AgNO₃(aq)using a current of 6.0 A?

Page 22

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Chapter 20: The Chemistry of Coordination Compounds

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Q1) Which of these ligands produces the weakest crystal field?

A)CN^-

B)I^-

C)OH^-

D)H₂O

E)NH₃

Q2) Predict the number of unpaired electrons in the [Cr(en)₃]²⁺ ion.

Q3) In the complex ion [ML₆]ⁿ⁺, Mⁿ⁺ has five d electrons and L is a strong field ligand. According to crystal field theory, the magnetic properties of the complex ion correspond to how many unpaired electrons?

A)0

B)1

C)2

D)3

E)5

Q4) What is the coordination number of cobalt in [Co(NH₃)₆]Cl₃?

Q5) What is the oxidation number of Fe in [Fe(CN)₆]^4-?

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Chapter 21: Nuclear Chemistry

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Q1) On a neutron number versus proton number plot, the nucleus of oxygen-15 lies below the belt of stability. Write an equation for radioactive decay of oxygen-15.

Q2) A sample of a radioisotope shows an activity of 999 disintegrations per minute due to beta decay. If after 1.10 years the activity is 952 disintegrations per minute, what is the half-life of this radioisotope?

A)4.38 * 10^-2 yr

B)11.4 yr

C)0.25 yr

D)15.8 yr

E)9.1 yr

Q3) Tritium is a radioisotope of hydrogen having a half-life of 12.3 years. If you initially had 1.0 * 10^-7 mole of tritium, calculate the decay rate of the sample.

Q4) When a ⁸⁷Br nucleus emits a beta particle, the nuclear species that results is ___________.

Q5) Tritium is a radioisotope of hydrogen having a half-life of 12.3 years. If you initially had 1.0 * 10^-7 mole of tritium, how many moles of tritium would be left after 78 years?

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Chapter 22: Organic Polymerssynthetic and Natural

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Q1) Which of these molecules is a product of the hydrolysis of DNA?

A)acetic acid

B)glucose

C)adenine

D)ribose

E)water

Q2) Which choice lists both the sugar and the nitrogen base that are a part of RNA but are not a part of DNA?

A)deoxyribose and thymine

B)ribose and deoxyribose

C)ribose and uracil

D)uracil and thymine

E)ribose and thymine

Q3) Which choice contains all three molecular units found in nucleotides?

A)phosphate, sugar, amino acid

B)amino acid, nitrogen-containing base, sugar

C)carboxylic acid, sugar, protein

D)phosphate, nitrogen-containing base, sugar

E)sugar, amino acid, protein

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