Preparatory Chemistry Pre-Test Questions - 3452 Verified Questions

Preparatory Chemistry

Pre-Test Questions

Course Introduction

Preparatory Chemistry is designed to provide students with a foundational understanding of basic chemical principles and concepts, serving as a bridge to more advanced chemistry courses. The course covers fundamental topics such as atomic structure, chemical bonding, stoichiometry, states of matter, and an introduction to chemical reactions. Emphasis is placed on developing problem-solving skills, laboratory techniques, and the application of scientific reasoning. This introductory course is ideal for students with limited prior experience in chemistry or those seeking to strengthen their background before enrolling in general or organic chemistry.

Recommended Textbook Chemistry A Molecular Approach 4th Edition by Nivaldo J. Tro

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25 Chapters

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Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) Water is an example of

A) a compound.

B) an element.

C) a heterogeneous mixture.

D) a homogeneous mixture.

Answer: A

Q2) The statement,"In a chemical reaction,matter is neither created nor destroyed" is called

A) the Law of Conservation of Mass.

B) Dalton's Atomic Theory.

C) the Scientific Method.

D) the Law of Multiple Proportions.

E) the Law of Definite Proportions.

Answer: A

Q3) Convert 3.3 m to meters.

A) 3.3 × 10^-9 m

B) 3.3 × 10^-6 m

C) 3.3 × 10^-3 m

D) 3.3 × 10⁶ m

Answer: B

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) What element is defined by the following information? p⁺ = 11 n° = 12

e^- = 11

A) sodium

B) vanadium

C) magnesium

D) titanium

E) selenium

Answer: A

Q2) Which of the following is equal to exactly Avogadro's number of atoms?

A) 8.00 grams of oxygen

B) 35.38 grams of zinc

C) 12.01 grams of carbon

D) 11.99 grams of sodium

E) 34.00 grams of selenium

Answer: C

Q3) The atomic number is equal to the number of ________.

Answer: protons

Q4) What group of elements in the periodic table are the most unreactive and why?

Answer: The noble gases are the most unreactive since they do not combine with other elements to form compounds.

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) Write the name for Mg₃(PO₄)₂.

A) magnesium(III) phosphite

B) magnesium(II) phosphite

C) magnesium phosphate

D) trimagnesium phosphorustetraoxide

E) magnesium phosphite

Answer: C

Q2) Calculate the mass percent composition of sulfur in Al₂(SO₄)₃.

A) 28.12%

B) 9.372%

C) 42.73%

D) 21.38%

E) 35.97%

Answer: A

Q3) Calculate the mass percent composition of oxygen in Al₂(SO₄)₃.

Answer: 56.11%

Q4) Give the name for HNO₂.

Answer: nitrous acid

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) Determine the reducing agent in the following reaction. 2 Li(s)+ Fe(C₂H₃O₂)₂(aq) 2 LiC₂H₃O₂(aq)+ Fe(s)

A) O

B) H

C) C

D) Fe

E) Li

Q2) A student dissolved 5.00 g of Co(NO₃)₂ in enough water to make 100.mL of stock solution.He took 4.00 mL of the stock solution and then diluted it with water to give 275.mL of a final solution.How many grams of NO₃^- ion are in the final solution?

A) 0.0247 g

B) 0.0493 g

C) 0.0678 g

D) 0.136 g

Q3) Define a limiting reagent.

Q4) Determine the oxidation state of Mn in KMnO₄.

Q5) Define a spectator ion.

Q6) What causes a precipitation reaction to occur between two soluble compounds?

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Chapter 5: Gases

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Sample Questions

Q1) Identify the gas particle that travels the fastest.

A) H₂

B) O₂

C) Ar

D) Cl₂

E) Kr

Q2) A 1.000 kg sample of nitroglycerine,C₃H₅N₃O₉,explodes and releases gases with a temperature of 1985°C at 1.000 atm.What is the volume of gas produced? 4 C₃H₅N₃O₉(s) 12 CO₂(g)+ 10 H₂O(g)+ 6 N₂(g)+ O₂(g)

A) 816.4 L

B) 3878 L

C) 5203 L

D) 5916 L

Q3) Why doesn't Dalton's Law of Partial Pressures depend on the identity of the gases present?

Q4) Define pressure.

Q5) Define effusion.

Q6) Why does HOT air rise?

Page 7

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Chapter 6: Thermochemistry

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Sample Questions

Q1) A student is preparing to perform a series of calorimetry experiments.She first wishes to determine the calorimeter constant (C_cal)for her coffee cup calorimeter.She pours a 50.0 mL sample of water at 345 K into the calorimeter containing a 50.0 mL sample of water at 298 K.She carefully records the final temperature of the water as 317 K.What is the value of C_calfor the calorimeter?

A) 19 J/K

B) 28 J/K

C) 99 J/K

D) 21 J/K

E) 76 J/K

Q2) Calculate the amount of heat (in kJ)necessary to raise the temperature of 51.8 g benzene by 52.6 K.The specific heat capacity of benzene is 1.05 J/g°C.

A) 1.81 kJ

B) 16.6 kJ

C) 2.59 kJ

D) 2.86 kJ

E) 3.85 kJ

Q3) Give the temperature and pressure for the standard state for a liquid.

Q4) Where does the energy absorbed during an endothermic reaction go?

Page 8

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) What is the maximum number of s orbitals that are possible?

A) 1

B) 2

C) 7

D) 0

E) 12

Q2) How much energy (in kJ)is required to ionize 2.52 moles of hydrogen atoms?

A) 1.65 × 10³ kJ

B) 1.52 × 10³ kJ

C) 5.49 × 10³ kJ

D) 3.31 × 10³ kJ

E) 5.89 × 10³ kJ

Q3) Give the numbers for m_l for an s orbital.

A) 0

B) -2, -1, 0, +1, +2

C) +1, +2, +3

D) 0, +1, +2, +3

E) +1

Q4) How many orbitals are contained in the n = 2 level? Give thel and m_l values of each of them.

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Place the following in order of decreasing IE₁. Cs Mg Ar

A) Cs > Mg > Ar

B) Mg > Ar > Cs

C) Ar > Mg > Cs

D) Cs > Ar > Mg

E) Mg > Cs > Ar

Q2) Why do successive ionization energies increase?

Q3) Identify the d-block element.

A) Th

B) W

C) Na

D) Te

E) None of the above

Q4) Give the ground state electron configuration for Br^-.

A) [Ar]4s²3d¹⁰4p⁵

B) [Ar]4s²3d¹⁰4p³

C) [Ar]4s²4p⁶

D) [Ar]4s²3d¹⁰4p⁶

E) [Ar]4s²3d⁸4p⁶

Q5) Give the ground state electron configuration for Cd²⁺.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Give the number of pairs of valence electrons for BCl₃.

A) 16

B) 8

C) 14

D) 10

E) 12

Q2) Identify the ionic compound that is used as an image enhancer in x-rays.

A) potassium permanganate

B) barium sulfate

C) lithium carbonate

D) sodium chloride

E) sodium bicarbonate

Q3) List the least electronegative atom.

Q4) Place the following in order of decreasing magnitude of lattice energy. NaF RbBr KCl

A) RbBr > NaF > KCl

B) NaF > KCl > RbBr

C) KCl > NaF > RbBr

D) NaF > RbBr > KCl

E) RbBr > KCl > NaF

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) The bond angle in H₂Se is A) 107°.

B) <104.5°.

C) 109.5°.

D) 190°.

E) 145°.

Q2) Give the electron geometry,molecular geometry,and hybridization for both carbons in CH₃COOH.

Q3) The hybrid orbital set used by the central atom in SeF₄ is A) sp

B) sp².

C) sp³.

D) sp³d.

E) sp³d².

Q4) Describe a pi bond.

A) side by side overlap of p orbitals

B) end to end overlap of p orbitals

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

D) overlap of two s orbitals

E) p orbital overlapping with a d orbital

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) In a liquid,the energy required to increase the surface of the area by a unit amount is called

A) viscosity.

B) surface tension.

C) dipole-dipole force.

D) hydrogen bonding.

E) capillary action.

Q2) Choose the substance with the highest vapor pressure at a given temperature.

A) SiS₂

B) RbBr

C) CH₃OCH₂CH₃

D) BH₃

E) SbH₃

Q3) Choose the substance with the lowest viscosity.

A) F₃CCF₃

B) F₂CHCH₂F

C) F₂CHCHF₂

D) FCH₂CH₂F

E) F₃CCHF₂

Q4) Why is the H_vap higher than H_fusfor a given compound?

Page 13

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) What is the edge length of a face-centered cubic unit cell made up of atoms having a radius of 125 pm?

A) 177 pm

B) 354 pm

C) 500 pm

D) 1000 pm

E) 800 pm

Q2) A common use of nylon 6,6 is

A) tire cord.

B) fishing line.

C) foam cups.

D) spray-on insulation.

E) films.

Q3) Identify the type of solid for argon.

A) metallic atomic solid

B) ionic solid

C) nonbonding atomic solid

D) molecular solid

E) networking atomic solid

Q4) Give the edge length in terms of r for a simple cubic cell.

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Chapter 13: Solutions

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Sample Questions

Q1) A solution is 2.25% by weight KHCO₃.How many grams of KHCO₃ are in 150.0 g of solution?

A) 1.50 g

B) 3.38 g

C) 66.7 g

D) 225 g

Q2) Define colloid.

Q3) A solution is prepared by dissolving 16.2 g of benzene (C₆H₆)in 282 g of carbon tetrachloride (CCl₄).The concentration of benzene in this solution is ________ molal.The molar masses of C₆H₆ and CCl₄ are 78.1 g/mol and 154 g/mol,respectively.

A) 7.36 × 10^-⁴

B) 0.736

C) 0.102

D) 0.0543

E) 5.43

Q4) Calculate the osmotic pressure,in torr,of a solution containing 3.00 mg of a sugar (342 g/mole)in 15.0 mL of water at 25°C.

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) The first-order decomposition of cyclopropane has a rate constant of 6.7 × 10^-4 s^-1.If the initial concentration of cyclopropane is 1.33 M,what is the concentration of cyclopropane after 644 s?

A) 0.43 M

B) 0.15 M

C) 0.94 M

D) 0.86 M

E) 0.67 M

Q2) The rate constant,k,for a first-order reaction is equal to 4.2 × 10^-4 s^-1.What is the half-life for the reaction?

A) 2.9 × 10^-4 s

B) 1.2 × 10³s

C) 1.7 × 10³ s

D) 2.4 × 10³ s

Q3) What function do enzymes serve?

Q4) What is the difference between average reaction rate and instantaneous reaction rate?

Q5) Define the frequency factor.

Q6) Define half-life.

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction at equilibrium.What effect will increasing the pressure of the reaction mixture have on the system? CuS(s)+ O₂(g) Cu(s)+ SO₂(g)

A) The equilibrium constant will decrease.

B) No effect will be observed.

C) The reaction will shift to the right in the direction of products.

D) The equilibrium constant will increase.

E) The reaction will shift to the left in the direction of reactants.

Q2) What is n for the following equation in relating K_c to K_p?

N₂(g)+ 3 F₂(g) 2 NF₃(g)

A) 4

B) -3

C) -2

D) 2

E) 0

Q3) Explain dynamic equilibrium.Use the generic reaction A(g) B(g)to explain.

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

Q5) Why is an equilibrium constant unitless?

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Chapter 16: Acids and Bases

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Q1) What is the selenide ion concentration [Se^2-] for a 0.100 M H₂Se solution that has the stepwise dissociation constants of K_a1 = 1.3 × 10^-4 and K_a2 = 1.0 × 10^-11?

A) 3.6 × 10^-3 M

B) 1.3 × 10^-4 M

C) 1.3 × 10^-5 M

D) 1.0 × 10^-11 M

Q2) Calculate the pH of a solution that contains 2.4 × 10^-⁵ M H₃O at 25°C.

A) 2.40

B) 9.38

C) 4.62

D) 8.38

E) 5.62

Q3) List the major source of energy in the United States.

Q4) Describe a molecule that can be a Lewis acid.

Q5) List the compound that is formed from the reaction of carbon dioxide with water.

Q6) What does the term amphoteric mean?

Q7) Describe the relationship between molecular structure and acid strength.

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Chapter 17: Aqueous Ionic Equilibrium

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Q1) If the pKa of HCHO₂ is 3.74 and the pH of an HCHO₂/NaCHO₂solution is 3.11,which of the following is TRUE?

A) [HCHO₂] < [NaCHO₂]

B) [HCHO₂] = [NaCHO₂]

C) [HCHO₂] << [NaCHO₂]

D) [HCHO₂] > [NaCHO₂]

E) It is not possible to make a buffer of this pH from HCHO₂ and NaCHO₂.

Q2) A 900.0 mL sample of 0.18 M HClO₄ is titrated with 0.27 M LiOH.Determine the pH of the solution after the addition of 300.0 mL of LiOH (this is the equivalence point).

A) 0.97

B) 13.03

C) 2.76

D) 11.24

E) 7.00

Q3) Identify the compound that is in stalactites and stalagmites.

Q4) Give the name of the compound that is in antifreeze and is toxic to pets.

Q5) Define a buffer.

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) Which of the following statements is TRUE?

A) Entropy is not a state function.

B) Endothermic processes decrease the entropy of the surroundings, at constant T and P.

C) Endothermic processes are never spontaneous.

D) Exothermic processes are always spontaneous.

E) None of the above is true.

Q2) Use the free energies of formation given below to calculate the equilibrium constant (K)for the following reaction at 298 K. 2 HNO₃(aq)+ NO(g) 3 NO₂(g)+ H₂O(l)K= ?

G°_f (kJ/mol)-110.9 87.6 51.3 -237.1

A) 8.71 × 10⁸

B) 0.980

C) 1.15 × 10^-9

D) 1.02

E) 5.11 × 10^-4

Q3) Define allotrope.

Q4) Why is heating your home with gas more efficient than heating it with electricity?

Q5) How is a nonspontaneous process made spontaneous?

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Chapter 19: Electrochemistry

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Q1) Use the tabulated half-cell potentials to calculate the equilibrium constant (K)for the following balanced redox reaction at 25°C. 3 I₂(s)+ 2 Fe(s) 2 Fe³⁺(aq)+ 6 I (aq)

A) 3.5 × 10^-59

B) 1.1 × 10¹⁷

C) 2.4 × 10⁵⁸

D) 8.9 × 10^-18

E) 1.7 × 10²⁹

Q2) What is the shorthand notation that represents the following galvanic cell reaction?

Pb(s)+ Cu(NO₃)₂(aq) Pb(NO₃)₂(aq)+ Cu(s)

A) Pb(s) Pb²⁺(aq) Cu²⁺(aq) Cu(s)

B) Cu(s) Cu²⁺(aq) Pb²⁺(aq) Pb(s)

C) Pb(s) NO₃^-(aq) NO₃^-(aq) Cu(s) D) Cu(s) Cu(NO₃)₂(aq)

Pb(NO₃)₂(aq) Pb(s)

Q3) What is the difference between a voltaic cell and an electrolytic cell?

Q4) Why,if we multiply a reaction by 2,don't we multiply its E°_red by 2?

Q5) Give an example of an inert electrode.

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Which of the following nuclides are most likely to decay via beta decay?

A) I-131

B) Ar-40

C) F-18

D) Zr-90

E) Pb-206

Q2) Determine the binding energy per nucleon of an Mg-24 nucleus.The Mg-24 nucleus has a mass of 23.985045.A proton has a mass of 1.00728 amu,a neutron has a mass of 1.008665 amu,and 1 amu is equivalent to 931.5 MeV of energy.

A) 0.3050 MeV

B) 8.83 MeV

C) 0.113 MeV

D) 192 MeV

E) 4.41 MeV

Q3) Give the goal of the Manhattan project.

A) to build an hydrogen bomb

B) to build the first particle accelerator at Cornell

C) to build an atomic bomb

D) to build the first nuclear reactor to generate electricity for Manhattan

E) to set up the electrification of Manhattan using DC current

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Chapter 21: Organic Chemistry

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Q1) Give the organic product for the following reaction in the presence of sulfuric acid. CH₃CH₂COOH + CH₃CH₂CH₂OH

A) CH₃CH₂COOCH₂CH₃ B)

CH₃CH₂COOCH₂CH₂CH₃ C)

CH₃CH₂CH₂COOCH₂CH₃ D)

CH₃CH₂COCH₂CH₂CH₃

E) CH₃CH₂OCH₂CH₂CH₃

Q2) Write the balanced chemical equation for the addition of HBr to CH₂=CHCH₂CH₃.

A) CH₂=CHCH₂CH₃ + HBr CH₃BrCH₂CH₂CH₃

B) CH₂=CHCH₂CH₃ + 2 HBr 2 CH₂BrCH₃

C) CH₂=CHCH₂CH₃ + 2 HBr CH₃Br + CH₂BrCH₂CH₃

D) CH₂=CHCH₂CH₃ + 4 HBr 4 CH₃Br

E) CH₂=CHCH₂CH₃ + HBr CH₃CHBrCH₂CH₃

Page 23

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Chapter 22: Biochemistry

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Q1) Which of the following has a tertiary protein structure?

A) ketone

B) amide

C) ribonucleotide

D) amino acid

E) fibrous protein

Q2) Identify the type of bond between the two strands of DNA.

A) hydrogen bonds

B) London dispersion forces

C) dipole-dipole bonds

D) amide bonds

E) water bonds

Q3) Identify the sugar in DNA.

A) ribose

B) deoxyribose

C) fructose

D) glucose

E) galactose

Q4) Why are lipids well-suited as structural components of cell membranes?

Q5) How do cis- and trans-fats differ?

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Chapter 23: Chemistry of the Nonmetals

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Q1) Determine the formula for an aluminosilicate where one fourth of the silicons are replaced by an aluminum atom.The charge is balanced by sodium ions.

A) Na₂(AlO₂)₂(SiO₂)

B) Na₂(AlO₂)₂(SiO₂)₃

C) Na(AlO₂)₂(SiO₂)

D) Na₂(AlO₂)₂(SiO₂)₂

E) Na(AlO₂)(SiO₂)₃

Q2) Which of the following is TRUE?

A) Phosphoric acid is used to produce the "fizziness" in soda.

B) Sodium phosphate is an additive in baked goods to give them a longer shelf life.

C) Phosphorus can range in oxidation state from -3 to +5.

D) The phosphorus in the phosphite ion has an oxidation state of - 3

E) None of the above is true.

Q3) Determine the oxidation state of chlorine in ClO^-.

A) 0

B) +1

C) -2

D) -1

E) +2

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Chapter 24: Metals and Metallurgy

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Q1) Galvanizing is when an object is dipped into a molten bath of A) magnesium.

B) lithium.

C) zinc.

D) sodium.

E) potassium.

Q2) Calcination is

A) heating an ore, in the presence of hydrogen or another substance, in order to cause a chemical reaction that drives off newly formed volatile compounds.

B) heating an ore, in the presence of helium or another substance, in order to purify the ore and obtain the liquid metal.

C) selectively dissolving a metal in solution to separate it from its ore.

D) the process of heating an ore to drive off volatile compounds.

E) forming metal parts using steam and large crystals of metal.

Q3) Why is zinc used to coat steel objects?

Q4) What is an advantage of hydrometallurgical processes over pyrometallurgical processes?

Q5) What is the difference between ferromagnetism and paramagnetism?

Q6) Why does a "two-phase" structure occur in a substitutional alloy?

Page 26

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) What is the highest possible oxidation state for cobalt?

A) +1

B) +7

C) +6

D) +4

E) +5

Q2) What is the Lanthanide contraction?

Q3) Identify the transition metal that is used in fat and carbohydrate synthesis in the human body.

A) potassium

B) zinc

C) copper

D) manganese

E) chromium

Q4) Which of the following species is diamagnetic?

A) an isolated, gas-phase Mn²⁺ ion

B) a high-spin octahedral Fe²⁺ complex

C) an isolated, gas-phase Cu²⁺ ion

D) a low-spin octahedral Co³⁺ complex

Q5) Why is +2 a common oxidation state for transition elements?

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