Introductory Chemistry Final Exam - 3452 Verified Questions

Introductory Chemistry Final Exam

Course Introduction

Introductory Chemistry provides a foundational understanding of the principles and concepts essential to the study of chemistry. This course covers topics such as atomic and molecular structure, chemical bonding, stoichiometry, states of matter, and basic chemical reactions. Students will explore the properties of elements and compounds, the use of the periodic table, and the principles of chemical equations, with an emphasis on problem-solving and laboratory techniques. Designed for students beginning their study of chemistry, the course aims to build scientific reasoning skills and prepare learners for more advanced chemistry coursework.

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

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Page 2

Chapter 1: Matter, measurement, and Problem Solving

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Sample Questions

Q1) Define the law of the conservation of energy. Answer: Energy is neither created nor destroyed.

Q2) Because of the high heat and low humidity in the summer in Death Valley,California,a visitor requires about one quart of water for every two miles traveled on foot.Calculate the approximate number of liters required for a person to walk 15 kilometers in Death Valley.

A) 4.4 L

B) 18 L

C) 46 L

D) 70 L

Answer: A

Q3) Define an atom.

Answer: An atom is the submicroscopic particle that constitutes the fundamental building block of ordinary matter.

Q4) Identify a liquid.

A) definite volume and definite shape

B) definite volume and no definite shape

C) definite shape and no definite volume

D) no definite shape and no definite volume

Answer: B

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Chapter 2: Atoms and Elements

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Q1) Which of the following does NOT describe a nonmetal?

A) tends to gain electrons

B) found in the upper right hand corner of the periodic table

C) poor conductor of electricity

D) nonmetals are generally unreactive

E) poor conductor of heat

Answer: D

Q2) Which of the following elements is in Group 4?

A) Hf

B) Ca

C) Bi

D) Ne

E) Cd

Answer: A

Q3) Which element has the chemical symbol,Rn?

A) radium

B) radon

C) rhenium

D) rhodium

Answer: B

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Chapter 3: Molecules,compounds,and Chemical Equations

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Sample Questions

Q1) How much calcium is contained in 35.0 g of calcium fluoride?

A) 1.99 g

B) 7.70 g

C) 15.0 g

D) 18.0 g

Answer: D

Q2) Determine the number of moles of oxygen in 3.20 moles of sodium hydrogen phosphate.

A) 6.40 moles of oxygen

B) 3.20 moles of oxygen

C) 9.60 moles of oxygen

D) 1.60 moles of oxygen

E) 12.80 moles of oxygen

Answer: E

Q3) Write a possible molecular formula for C₄H₄O.

A) C₂₄H₂₄O₆

B) C₁₂H₁₂O₂

C) C₂H₂O

D) C₈H₈O

Answer: A

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Chapter 4: Chemical Quantities and Aqueous Reactions

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Sample Questions

Q1) What volume of a 0.540 M NaOH solution contains 12.5 g of NaOH?

A) 0.169 L

B) 0.579 L

C) 1.73 L

D) 5.92 L

Q2) Dinitrogen monoxide gas decomposes to form nitrogen gas and oxygen gas.How many grams of oxygen are formed when 10.0 g of dinitrogen monoxide decomposes?

A) 0.275 g

B) 3.64 g

C) 7.27 g

D) 14.5 g

Q3) Balance the chemical equation given below,and determine the number of grams of MgO that are needed to produce 20.0 g of Fe₂O₃. _____ MgO(s)+ _____ Fe(s) _____ Fe₂O₃(s)+ _____ Mg(s)

A) 0.0660 g

B) 1.68 g

C) 5.05 g

D) 15.1 g

Q4) Describe the difference between complete ionic and net ionic equations.

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Chapter 5: Gases

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Sample Questions

Q1) Give the definition for diffusion.

A) gas molecules mix equally

B) gas molecules spread out in response to a concentration gradient

C) gas molecules escape from a container into a vacuum through a small hole

D) average distance between collisions

E) gas molecules mix unequally

Q2) A syringe contains 0.65 moles of He gas that occupy 900.0 mL.What volume (in L)of gas will the syringe hold if 0.35 moles of Ne is added?

A) 0.48 L

B) 1.7 L

C) 1.4 L

D) 2.6 L

E) 0.80 L

Q3) A gas bottle contains 0.450 mol of gas at 730 mm Hg pressure.If the final pressure is 1.15 atm,how many moles of gas were added to the bottle?

A) 0.0471 mol

B) 0.0888 mol

C) 0.497 mol

D) 0.539 mol

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Identify the energy source with the highest negative H°_rxn.

A) 1 mole ethanol

B) 1 mole methane

C) 1 mole carbon

D) 1 mole petroleum

E) All have the same value.

Q2) When 1.365 g of anthracene,C₁₄H₁₀,is combusted in a bomb calorimeter that has a water jacket containing 500.0 g of water,the temperature of the water increases by 25.89°C.Assuming that the specific heat of water is 4.18 J/(g °C),and that the heat absorption by the calorimeter is negligible,estimate the enthalpy of combustion per mole of anthracene.

A) +39.7 kJ/mol

B) -39.7 kJ/mol

C) -7060 kJ/mol

D) -8120 kJ/mol

Q3) For E_systo always be -,what must be TRUE?

A) q = w

B) +q > -w

C) +w > -q

D) -w > +q

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Place the following types of electromagnetic radiation in order of decreasing energy. visible light radio waves infrared light

A) radio waves > infrared light > visible light

B) visible light > infrared light > radio waves

C) radio waves > visible light > infrared light

D) visible light > radio waves > infrared light

E) infrared light > radio waves > visible light

Q2) Which of the following statements is TRUE?

A) We can sometimes know the exact location and speed of an electron at the same time.

B) All orbitals in a given atom are roughly the same size.

C) Since electrons have mass, we must always consider them to have particle properties and never wavelike properties.

D) Atoms are roughly spherical because when all of the different shaped orbitals are overlapped, they take on a spherical shape.

E) All of the above are true.

Q3) What is the photoelectric effect?

Q4) Give an example of a d orbital.

Q5) Describe how a neon light works.

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) List the noble gas that has the highest ionization energy.

Q2) Give the number of valence electrons for Te.

A) 1

B) 7

C) 3

D) 5

E) 6

Q3) Describe the reaction of the alkali metals with non-metals.

A) inert

B) vigorous

C) mild reaction

D) forms water

E) dissolves

Q4) Of the following,which element has the highest first ionization energy?

A) beryllium

B) hydrogen

C) boron

D) lithium

Q5) Define paramagnetic.

Q6) Give the number of valence electrons for Mg²⁺.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Use the bond energies provided to estimate H°_rxn for the reaction below. PCl₃(g)+ Cl₂(g) PCl₅(l) H°_rxn= ?

Bond Bond Energy (kJ/mol)

Cl-Cl 243

P-Cl 331

A) -243 kJ

B) -419 kJ

C) -662 kJ

D) -67 kJ

E) -905 kJ

Q2) The compound BrCl contains

A) ionic bonds.

B) nonpolar covalent bonds.

C) polar covalent bonds with partial negative charges on the Cl atoms.

D) polar covalent bonds with partial negative charges on the Br atoms.

Q3) Draw the Lewis Dot structure for Al³⁺.

Q4) List the least electronegative atom.

Q5) Draw the Lewis structure for BrO₃^-.Make sure to include any important resonance structures.

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence

Bond Theory,

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Sample Questions

Q1) Determine the electron geometry (eg)and molecular geometry (mg)of the underlined atom CH₃OCH₃.

A) eg = bent, mg = tetrahedral

B) eg = linear, eg = bent

C) eg = tetrahedral, mg = bent

D) eg = trigonal bipyramidal, mg = trigonal planar

E) eg = octahedral, mg = square planar

Q2) Consider the following compound.How many sigma and pi bonds does it contain? CH₃CHCHCO₂H

A) 9 sigma, 4 pi

B) 11 sigma, 2 pi

C) 9 sigma, 0 pi

D) 13 sigma, 2 pi

E) 6 sigma, 2 pi

Q3) Give the molecular geometry and number of electron groups for BrF₅.

A) square planar, 6 electron groups

B) square pyramidal, 6 electron groups

C) T-shaped, 5 electron groups

D) octahedral, 6 electron groups

E) seesaw, 5 electron groups

Page 12

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Chapter 11: Liquids,solids,and Intermolecular Forces

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Sample Questions

Q1) Define condensation.

A) A liquid becomes a gas.

B) A gas becomes a liquid.

C) A gas becomes a solid.

D) A liquid becomes a solid.

E) A solid becomes a liquid.

Q2) Capillary action occurs because

A) cohesive forces are greater than adhesive forces.

B) adhesive forces equal adhesive forces.

C) adhesive forces are greater than cohesive forces.

D) surface tension is high.

E) surface tension is low.

Q3) 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

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

Q5) Define volatile.

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Chapter 12: Solids and Modern Materials

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Sample Questions

Q1) Nickel has a face-centered cubic structure and has a density of 8.90 g/cm³.What is its atomic radius?

A) 125 pm

B) 249 pm

C) 353 pm

D) 997 pm

Q2) Identify the largest unreinforced concrete dome in existence.

A) Pantheon

B) Hagia Sophia

C) Orbetello aircraft hangar

D) Ream's turtle

E) Turin's exposition hall

Q3) Identify the compound that has the zinc blende structure.

A) silver iodide

B) barium chloride

C) potassium bromide

D) cesium chloride

E) water

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

Q5) Describe the difference between the conduction band and the valence band.

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Chapter 13: Solutions

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Sample Questions

Q1) What is the freezing point of a solution of 7.15 g MgCl₂ in 100 g of water?

K_f for water is 1.86°C/m.

A) -0.140°C

B) -1.40°C

C) -2.80°C

D) -4.18°C

Q2) Dynamic equilibrium can be defined as

A) rate of dissolution = rate of recrystallization.

B) rate of dissolution < rate of recrystallization.

C) rate of dissolution > rate of recrystallization.

D) rate of bubbling > rate of dissolving.

E) rate of condensing > rate of bubbling.

Q3) Calculate the molality of a solution formed by dissolving 34.8 g of LiI in 500.0 mL of water.

A) 0.254 m

B) 0.130 m

C) 0.696 m

D) 0.260 m

E) 0.520 m

Q4) Define colloid.

Page 15

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Chapter 14: Chemical Kinetics

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Sample Questions

Q1) What data should be plotted to show that experimental concentration data fits a second-order reaction?

A) ln[reactant] vs. time

B) [reactant] vs. time

C) ln(k) vs. 1/T

D) 1/[reactant] vs. time

E) ln(k) vs. E_a

Q2) The rate constant for a zero-order reaction is 0.54 s^-1.What is the half-life of this reaction if the initial concentration is 0.33 M?

A) 0.089 s

B) 1.8 s

C) 0.31 s

D) 5.6 s

E) 1.3 s

Q3) What is the difference between average reaction rate and instantaneous reaction rate?

Q4) What function do enzymes serve?

Q5) Explain what the exponential factor in the Arrhenius equation represents.

Q6) What is a catalyst and what function does it serve?

Page 16

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Chapter 15: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction and its equilibrium constant: I₂(g) 2 I(g)K_p = 0.209

A reaction mixture contains 0.72 atm I₂ and 1.80 atm I.Which of the following statements is TRUE concerning this system?

A) The reaction will shift in the direction of reactants.

B) The reaction quotient will increase.

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

D) The equilibrium constant will decrease.

E) The system is at equilibrium.

Q2) Consider the following reaction at equilibrium.What effect will reducing the volume 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.

Q3) Can the K_p and K_c for a reaction ever have the same value? Why or why not?

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Chapter 16: Acids and Bases

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Q1) Determine the ammonia concentration of an aqueous solution that has a pH of 11.00.The equation for the dissociation of NH₃ (K_b= 1.8 × 10^-5)is NH₃(aq)+ H₂O(l) NH₄⁺(aq)+ OH^-(aq)

A) 3.0 M

B) 0.056 M

C) 1.8 × 10^-² M

D) 1.0 × 10^-3 M

Q2) What is the conjugate acid of HCO₃ ?

A) H₃O⁺

B) H₂O

C) CO₃^2-

D) OH

E) H₂CO₃

Q3) An Arrhenius acid

A) donates an electron pair.

B) is a H⁺ donor.

C) is a H⁺ acceptor.

D) produces OH^-in aqueous solutions.

E) produces H⁺ in aqueous solutions.

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Chapter 17: Aqueous Ionic Equilibrium

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Q1) What is the pH of a solution prepared by mixing 50.00 mL of 0.10 M NH₃ with 5.00 mL of 0.10 M NH₄Cl? Assume that the volume of the solutions are additive and that K_b = 1.8 × 10^-5for NH₃.

A) 8.25

B) 9.28

C) 10.26

D) 11.13

Q2) A complex ion contains

A) a central nonmetal ion bound to one or more ligands.

B) an anion bound to water.

C) water bound to one or more ligands.

D) a central metal ion bound to one or more ligands.

E) ammonia bound to water.

Q3) Identify a good buffer.

A) small amounts of both a weak acid and its conjugate acid

B) significant amounts of both a strong base and a strong acid

C) small amounts of both a strong acid and a conjugate acid

D) significant amounts of both a weak acid and a strong acid

E) significant amounts of both a weak acid and its conjugate base

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Chapter 18: Free Energy and Thermodynamics

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Sample Questions

Q1) Consider a reaction that has a negative H and a negative S.Which of the following statements is TRUE?

A) This reaction will be spontaneous only at low temperatures.

B) This reaction will be spontaneous at all temperatures.

C) This reaction will be nonspontaneous at all temperatures.

D) This reaction will be nonspontaneous only at low temperatures.

E) It is not possible to determine without more information.

Q2) For a given compound,list the decreasing order of entropy for a liquid,solid,and gas.

A) solid > gas > liquid

B) liquid > solid > gas

C) gas > liquid > solid

D) gas > solid > liquid

E) solid > liquid > gas

Q3) Place the following in order of increasing entropy at 298 K. Ne Xe He Ar Kr

A) He < Kr < Ne < Ar < Xe

B) Xe < Kr < Ar < Ne < He

C) Ar < He < Ar < Ne < Kr

D) Ar < Ne < Xe < Kr < He

E) He < Ne < Ar < Kr < Xe

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Chapter 19: Electrochemistry

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Sample Questions

Q1) What statement is NOT true about standard electrode potentials?

A) E°_cell^{is positive for spontaneous reactions.}

B) Electrons will flow from more negative electrode to more positive electrode.

C) The electrode potential of the standard hydrogen electrode is exactly zero.

D) E°_cellis the difference in voltage between the anode and the cathode.

E) The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative to the standard hydrogen electrode and therefore has a positive E°.

Q2) Use the tabulated half-cell potentials to calculate G° for the following balanced redox reaction. Pb²⁺(aq)+ Cu(s) Pb(s)+ Cu²⁺(aq)

A) -41 kJ

B) -0.47 kJ

C) +46 kJ

D) +91 kJ

E) -21 kJ

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) What is electrolysis?

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Chapter 20: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Describe what is meant by the term "valley of stability"?

Q2) Write a nuclear equation to describe the neutron induced fission of U-235 to form Xe-134 and Sr-100.Determine how many neutrons are produced in the reaction.

A) 4

B) 3

C) 1

D) 0

E) 2

Q3) A freshly prepared sample of curium-243 undergoes 3312 disintegrations per second.After 6.00 yr,the activity of the sample declines to 2755 disintegrations per second.The half-life of curium-243 is ________ yr.

A) 4.99

B) 32.6

C) 7.21

D) 0.765

E) 22.6

Q4) List two problems associated with nuclear power.

Q5) Define radioactivity.

Q6) What is the "mass defect"?

22

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Chapter 21: Organic Chemistry

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Q1) Identify the alkane with the highest boiling point.

A) pentane

B) butane

C) ethane

D) methane

E) propane

Q2) A carboxylic acid reacts with itself to form

A) an aldehyde.

B) an amine.

C) an ester.

D) a thiol.

E) an anhydride.

Q3) Identify the formula for an alkane.

A) C_nH_2n+2

B) C_nH_2n-2

C) C_nH_2n

D) C_nH_n

E) C_nH_2n+4

Q4) What does the term "racemic mixture" mean?

Q5) Why doesn't benzene typically undergo addition reactions like the alkenes do?

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Chapter 22: Biochemistry

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Q1) Which of the following is a type of nucleic acid?

A) DNA

B) amino acid

C) tripeptide

D) carbohydrate

E) lipid

Q2) Identify the bonds that connect the protein structures in a pleated sheet.

A) ester linkages

B) sulfide linkages

C) peptide bonds

D) hydrogen bonds

E) glycosidic linkages

Q3) Identify the disaccharide.

A) estradiol

B) cellulose

C) galactose

D) triglyceride

E) sucrose

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) Describe the major production method for obtaining oxygen.

Q2) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

Q3) Which of the following statements is TRUE?

A) Boron is a large component of Earth's crust.

B) Boron is usually found in its elemental state in the Earth's crust.

C) Due to boron's small size and low electronegativity, it behaves as a semimetal instead of a metal.

D) Boron atoms don't typically bond to one another due to its small atomic radius.

E) There are at least 5 allotropes of elemental boron.

Q4) Why are phosphate compounds added to detergents?

Q5) Identify a nonmetal.

A) P

B) Ca

C) Pd

D) Ni

E) Ge

Q6) Draw the Lewis structure for IF₄ and determine the geometry of the ion.

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Chapter 24: Metals and Metallurgy

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Q1) In an interstitial alloy,

A) one nonmetal substitutes for another nonmetal in a crystal structure.

B) small, usually nonmetallic atoms fit in between the metallic atoms of a crystal.

C) small metallic atoms fit in between the nonmetallic atoms of a crystal

D) one metal substitutes for another metal in a crystal structure.

E) one metal substitutes for a nonmetal in a crystal structure.

Q2) Why can't chromium and nickel form a miscible solid solution over the entire composition range?

A) The lever rule says that these two metals can't coexist in an alloy.

B) The two metals crystallize into different cubic structures in their pure forms.

C) The tetrahedral holes in the nickel crystal are not large enough to accommodate the chromium.

D) The octahedral holes in the chromium crystal are not large enough to accommodate the nickel.

E) None of the above is true.

Q3) What is the difference between ferromagnetism and paramagnetism?

Q4) Describe the difference between a substitutional alloy and an interstitial alloy.

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Chapter 25: Transition Metals and Coordination Compounds

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Q1) Choose the electron configuration for Nb³⁺.

A) [Kr]5s²4d²

B) [Kr]5s²

C) [Kr]4d²

D) [Ar]4s²3d¹⁰4p⁶

E) [Kr]5s²4d⁴

Q2) What type of geometry (according to valence bond theory)does V exhibit in the complex ion: [V(NH₃)₄]²⁺?

A) square bipyramidal

B) see-saw

C) trigonal pyramidal

D) bent

E) square planar

Q3) Identify the isomers that have ligands with different spatial arrangements about the metal ions.

A) linkage isomers

B) geometric isomers

C) coordination isomers

D) optical isomers

E) structural isomers

Page 27

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