Chemistry for Life Sciences
Practice Questions
Course Introduction
This course provides an introduction to the fundamental concepts of chemistry with a focus on their application in the life sciences. Topics include atomic and molecular structure, chemical bonding, thermodynamics, acids and bases, and chemical kinetics, with particular emphasis on biological molecules and reactions relevant to living systems. Students will develop a strong foundation in chemical principles while exploring the role of chemistry in areas such as cellular metabolism, genetic information, and physiological processes, preparing them for advanced studies in biology, medicine, and related fields.
Recommended Textbook
Chemistry The Molecular Nature of Matter and Change 7th Edition by Silberberg
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Page 2
Chapter 1: Keys to the Study of Chemistry
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Sample Questions
Q1) The SI prefix mega- (M) means
A) 10<sup>-6</sup>.
B) 10<sup>-3</sup>.
C) 10<sup>3</sup>.
D) 10<sup>6</sup>.
E) 10<sup>9</sup>.
Answer: D
Q2) The numerical value of any temperature expressed in Celsius is always different from the numerical value of the same temperature in Fahrenheit.
A)True
B)False
Answer: False
Q3) When a wooden match burns in air, chemical potential energy is converted to kinetic energy.
A)True
B)False
Answer: True
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Chapter 2: The Components of Matter
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Sample Questions
Q1) Millikan's oil-drop experiment
A) established the charge on an electron.
B) showed that all oil drops carried the same charge.
C) provided support for the nuclear model of the atom.
D) suggested that some oil drops carried fractional numbers of electrons.
E) suggested the presence of a neutral particle in the atom.
Answer: A
Q2) Ionic compounds may carry a net positive or negative charge.
A)True
B)False
Answer: False
Q3) Sodium peroxide is an oxidizer used to bleach animal and vegetable fibers. What is its formula?
A) NaO
B) NaO<sub>2</sub>
C) Na<sub>2</sub>O<sub>2</sub>
D) Na<sub>2</sub>O
E) NaH<sub>2</sub>O<sub>2</sub>
Answer: C
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4
Chapter 3: Stoichiometry of Formulas and Equations
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Sample Questions
Q1) Phosphine, an extremely poisonous and highly reactive gas, will react with oxygen to form tetraphosphorus decaoxide and water. PH<sub>3</sub>(g) + O<sub>2</sub>(g) \(\to\) P<sub>4</sub>O<sub>10</sub>(s) + H<sub>2</sub>O(g) [unbalanced]
Calculate the mass of P<sub>4</sub>O<sub>10</sub>(s) formed when 225 g of PH<sub>3 </sub>reacts with excess oxygen.
A) 1880 g
B) 940. g
C) 900. g
D) 470 g
E) 56.3 g
Answer: D
Q2) One mole of O<sub>2</sub> has a mass of 16.0 g.
A)True
B)False
Answer: False
Q3) Analysis of a white solid produced in a reaction between chlorine and phosphorus showed that it contained 77.44% chlorine and 22.56% phosphorus. What is its empirical formula?
Answer: PCl<sub>3</sub>
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Chapter 4: The Major Classes of Chemical Reactions
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Sample Questions
Q1) Potassium carbonate, K<sub>2</sub>CO<sub>3</sub>, sodium iodide, NaI, magnesium chloride, MgCl<sub>2</sub>, methanol, CH<sub>3</sub>OH, and ammonium chloride, NH<sub>4</sub>Cl, are soluble in water. Which produces the largest number of dissolved particles per mole of dissolved solute?
A) K<sub>2</sub>CO<sub>3</sub>
B) NaI
C) MgCl<sub>2</sub>
D) CH<sub>3</sub>OH
E) NH<sub>4</sub>Cl
Q2) Which one of the following substances is the best electrolyte?
A) CO
B) CH<sub>3</sub>Cl
C) CH<sub>4</sub>
D) C<sub>2</sub>H<sub>5</sub>OH
E) HCl
Q3) The combustion of an element is always a combination reaction.
A)True
B)False
Q4) Complete and balance the equation for the following acid-base reaction. Ca(OH)<sub>2</sub>(aq) + HCl(aq) \(\to\)
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Chapter 5: Gases and the Kinetic-Molecular Theory
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Sample Questions
Q1) According to the postulates of kinetic-molecular theory, the average kinetic energy of gas molecules is proportional to the absolute temperature.
A)True
B)False
Q2) A sample of nitrogen gas is confined to a 14.0 L container at 375 torr and 37.0°C. How many moles of nitrogen are in the container?
A) 0.271 mol
B) 2.27 mol
C) 3.69 mol
D) 206 mol
E) 227 mol
Q3) The ideal gas law tends to become inaccurate when
A) the pressure is lowered and molecular interactions become significant.
B) the pressure is raised and the temperature is lowered.
C) the temperature is raised above the temperature of STP.
D) large gas samples are involved.
E) the volume expands beyond the standard molar volume.
Q4) Briefly state the conditions corresponding to STP (standard temperature and pressure).
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Chapter 6: Thermochemistry: Energy Flow and Chemical Change
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Sample Questions
Q1) For all processes, both q and \(\Delta\)E will have the same sign.
A)True
B)False
Q2) Benzene is a starting material in the synthesis of nylon fibers and polystyrene (styrofoam). Its specific heat capacity is 1.74 J/(g·K). If 16.7 kJ of energy is absorbed by a 225-g sample of benzene at 20.0°C, what is its final temperature?
A) -22.7°C
B) 36.7°C
C) 42.7°C
D) 62.7°C
E) none of the above
Q3) Although internal energy (E) is more fundamental and conceptually easier than enthalpy (H), in most chemical applications \(\Delta\)H is more relevant and useful than \(\Delta\)E. Why?
Q4) The only way in which a system can do work on the surroundings is by expansion against the external pressure.
A)True
B)False
Page 8
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Chapter 7: Quantum Theory and Atomic Structure
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Sample Questions
Q1) Which of the following is a correct set of quantum numbers for an electron in a 3d orbital?
A) n = 3, l = 0, m<sub>l</sub> = -1
B) n = 3, l = 1, m<sub>l</sub> = +3
C) n = 3, l = 2, m<sub>l</sub><sub> </sub>= 3
D) n = 3, l = 3, m<sub>l</sub> = +2
E) n = 3, l = 2, m<sub>l</sub><sub> </sub>= -2
Q2) The orientation in space of an atomic orbital is associated with
A) the principal quantum number (n).
B) the angular momentum quantum number (l).
C) the magnetic quantum number (m<sub>l</sub>).
D) the spin quantum number (m<sub>s</sub>).
E) none of the above.
Q3) Use the Bohr equation to calculate the energy of
a. the largest energy absorption or emission process involving the n = 2 state of the hydrogen atom.
b. the smallest energy absorption or emission process involving the n = 2 state of the hydrogen atom.
Q4) What is the speed of an electron in m/s if its wavelength is 0.155 nm?
Q5) The following combinations of quantum numbers are not allowed.
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Chapter 8: Electron Configuration and Chemical Periodicity
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Sample Questions
Q1) Select the most basic compound from the following:
A) Bi<sub>2</sub>O<sub>3</sub>
B) SiO<sub>2</sub>
C) Cs<sub>2</sub>O
D) Na<sub>2</sub>O
E) H<sub>2</sub>O
Q2) Which of the following elements has the largest first ionization energy?
A) Na
B) Cl
C) Ca
D) Te
E) Br
Q3) How many valence electrons are there in an atom with the electron configuration [noble gas]ns<sup>2</sup>(n - 1)d<sup>10</sup>np<sup>3</sup>?
A) 2
B) 3
C) 5
D) 10
E) 15
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Chapter 9: Models of Chemical Bonding
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Sample Questions
Q1) Covalent bonding typically occurs when a ______________ bonds with a ______________.
Q2) In covalent bond formation, the potential energy reaches a maximum when the internuclear distance is equal to the bond length.
A)True
B)False
Q3) Which one of the following properties is least characteristic of substances composed of small, covalently-bonded molecules?
A) low melting point
B) low boiling point
C) weak bonds
D) poor electrical conductor when solid
E) poor electrical conductor when molten
Q4) The electrostatic energy of two charged particles is inversely proportional to the square of the distance between them.
A)True
B)False
Q5) Most of the copper sold in major metal markets is highly purified, typically to 99.99%. Why is this?
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Chapter 10: The Shapes of Molecules
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Sample Questions
Q1) According to VSEPR theory, a molecule with the general formula AX<sub>5</sub>E will have a ______ molecular shape.
A) tetrahedral
B) trigonal bipyramidal
C) square pyramidal
D) octahedral
E) see-saw
Q2) In order for a non-cyclic triatomic molecule to be bent, VSEPR theory requires that there must be two lone pairs on the central atom.
A)True
B)False
Q3) 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 (HCO<sub>2</sub><sup>-</sup>, carbon is the central atom) to illustrate.
Q4) Bond angles of 180° only occur around atoms which display linear molecular geometry.
A)True
B)False
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Chapter 11: Theories of Covalent Bonding
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Sample Questions
Q1) In applying molecular orbital theory, the bond order is calculated in the same way as with Lewis structures.
A)True
B)False
Q2) A molecule with the formula AX<sub>3</sub>E uses _________ to form its bonds.
A) s and p atomic orbitals
B) sp<sup>3</sup> hybrid orbitals
C) sp<sup>2</sup> hybrid orbitals
D) sp hybrid orbitals
E) sp<sup>3</sup>d<sup>2</sup> hybrid orbitals
Q3) According to molecular orbital theory, all diatomic molecules with an even number of electrons will be diamagnetic.
A)True
B)False
Q4) Hybrid orbitals of the sp<sup>3</sup>d type occur in sets of four.
A)True
B)False
Q5) What type of hybridization is needed to explain why ethyne, C<sub>2</sub>H<sub>2</sub>, is linear?
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Chapter 12: Intermolecular Forces: Liquids, Solids, and Phase Changes
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Sample Questions
Q1) Which of the following pairs of molecules can form hydrogen bonds between them?
A) HCl and HI
B) CH<sub>3</sub>OH and NH<sub>3</sub>
C) CH<sub>4</sub> and H<sub>2</sub>O
D) SO<sub>2</sub> and CH<sub>2</sub>O
E) H<sub>2</sub> and O<sub>2</sub>
Q2) Which of the following liquids is likely to have the highest surface tension?
A) Br<sub>2</sub>
B) C<sub>8</sub>H<sub>18</sub>
C) CH<sub>3</sub>OCH<sub>3</sub>
D) CH<sub>3</sub>OH
E) Pb
Q3) Which one of the following substances does not exist in the indicated solid type?
A) graphite - network
B) Na - metallic
C) SiO<sub>2</sub> - molecular
D) NaCl - ionic
E) diamond - network
Q4) How do the electrical properties of semiconductors differ from those of metals?
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Chapter 13: The Properties of Mixtures: Solutions and Colloids
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Sample Questions
Q1) Select the weakest electrolyte from the following set.
A) Na<sub>2</sub>SO<sub>4</sub>
B) KCl
C) CH<sub>3</sub>CH<sub>2</sub>COOH, propionic acid
D) CaCl<sub>2</sub>
E) LiOH
Q2) Children under the age of six with more than 0.10 ppm of lead in their blood can suffer a reduction in I.Q. or have behavior problems. What is the molality of a solution which contains 0.10 ppm of lead?
A) 4.8 × 10<sup>-10</sup> m
B) 4.8 × 10<sup>-7</sup> m
C) 4.8 × 10<sup>-4</sup> m
D) 4.8 × 10<sup>-1</sup> m
E) none of the above
Q3) Which of the following aqueous solutions will have the lowest osmotic pressure?
A) 0.10 m KOH
B) 0.10 m RbCl
C) 0.05 m CaSO<sub>4</sub>
D) 0.05 m BaCl<sub>2</sub>
E) 0.10 m K<sub>2</sub>SO<sub>4</sub>
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Chapter 14: Periodic Patterns in the Main Group Elements:
Bonding, Structure, and Reactivity
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Sample Questions
Q1) Elements of group 1A have higher melting and boiling points than those of group 2A.
A)True
B)False
Q2) Bromine will form compounds with each of the other elements in Period 4 of the periodic table. How does the type of bonding in the compounds change as one moves from potassium bromide to selenium bromide?
A) ionic to polar covalent
B) polar covalent to ionic
C) polar covalent to non-polar covalent
D) coordinate covalent to polar covalent
E) none of the above
Q3) Predict the products for the following set of reactants. H<sub>2</sub>O(l) + ClF<sub>5</sub>(l) \(\to\)
A) HClO<sub>2</sub>(aq) + HF(aq)
B) HClO<sub>3</sub>(aq) + HF(aq)
C) HClO(aq) + HF(aq)
D) HCl(aq) + HF(aq)
E) HClO(aq) + HFO(aq)
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Chapter 15: Organic Compounds and the Atomic Properties of
Carbon
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Q1) The backbone of protein molecules consists of repeating N-C-C-O units.
A)True
B)False
Q2) Explain how a disulfide bridge can arise a protein molecule.
Q3) Esters can be formed by the dehydration-condensation of a carboxylic acid and an alcohol.
A)True
B)False
Q4) A characteristic reaction of alkenes is addition.
A)True
B)False
Q5) Each amino acid has two functional groups in common and one of 20 other groups attached to the \(\alpha\)-carbon. The two functional groups are
A) carboxyl and amine.
B) ester and amine.
C) carboxyl and amide.
D) alcohol and amine.
E) carboxyl and peptide.
Q7) Explain how a salt bridge can arise a protein molecule. Page 17
Q6) Draw and name four alkenes with the formula C<sub>4</sub>H<sub>8</sub>.
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Chapter 16: Kinetics: Rates and Mechanisms of Chemical Reactions
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Q1) The elementary reaction HBr(g) + Br(g) \(\to\) H(g) + Br<sub>2</sub>(g) is endothermic.
a. Would you expect the rate constant for the back reaction to be smaller or larger than that for the forward reaction? Explain, briefly.
b. Draw a fully-labeled reaction energy diagram for this reaction, showing the locations of the reactants, products, and transition state.
Q2) In a reversible reaction, a catalyst will speed up the forward reaction but not affect the reverse reaction.
A)True
B)False
Q3) A reaction has an activation energy of 195.0 kJ/mol. When the temperature is increased from 200.°C to 220.°C, the rate constant will increase by a factor of A) 1.1.
B) 4.3 × 10<sup>4</sup>.
C) 3.2.
D) 7.5.
E) none of the above.
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Chapter 17: Equilibrium: the Extent of Chemical Reactions
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Sample Questions
Q1) Hydrogen iodide, HI, is formed in an equilibrium reaction when gaseous hydrogen and iodine gas are heated together. If 20.0 g of hydrogen and 20.0 g of iodine are heated, forming 10.0 g of hydrogen iodide, what mass of hydrogen remains unreacted?
A) 10.0 g hydrogen remains
B) 10.9 g hydrogen remains
C) 15.0 g hydrogen remains
D) 19.9 g hydrogen remains
E) Need to know the equilibrium constant in order to calculate the answer.
Q2) Which of the following has an effect on the magnitude of the equilibrium constant?
A) removing products as they are formed
B) adding more of a reactant
C) adding a catalyst
D) increasing the pressure, in a gas-phase reaction
E) change in temperature
Q3) When a reaction system reaches equilibrium, the forward and reverse reactions stop.
A)True
B)False
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Chapter 18: Acid-Base Equilibria
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Q1) Butyric acid is responsible for the odor in rancid butter. A solution of 0.25 M butyric acid has a pH of 2.71. What is the K<sub>a</sub> for the acid?
A) 0.36
B) 2.4 × 10<sup>-2</sup>
C) 7.8 × 10<sup>-3</sup>
D) 1.5 × 10<sup>-5</sup>
E) none of the above
Q2) Which of the following aqueous liquids will have the highest pH?
A) 0.1 M CH<sub>3</sub>COOH, pK<sub>a</sub> = 4.7
B) 0.1 M CuCl<sub>2</sub>, pK<sub>a</sub> = 7.5
C) 0.1 M H<sub>3</sub>C<sub>6</sub>H<sub>5</sub>O<sub>7</sub>, pK<sub>a</sub> = 3.1
D) 0.1 M ZnCl<sub>2</sub>, pK<sub>a</sub> = 9.0
E) pure water
Q3) Which of the following is a Lewis base?
A) BCl<sub>3</sub>
B) Cu<sup>2+</sup>
C) Cl<sup>-</sup>
D) Mn<sup>2+</sup>
E) NH<sub>4</sub><sup>+</sup>
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Chapter 19: Ionic Equilibria in Aqueous Systems
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Q1) Calculate the solubility of zinc hydroxide, Zn(OH)<sub>2</sub>, in 1.00 M NaOH. K<sub>sp</sub> = 3.0 × 10<sup>-16</sup> for Zn(OH)<sub>2</sub>, K<sub>f</sub> = 3.0 × 10<sup>15</sup> for Zn(OH)<sub>4</sub><sup>2-</sup>
A) 0.60 M
B) 0.52 M
C) 0.37 M
D) 0.32 M
E) 0.24 M
Q2) A buffer is prepared by adding 150 mL of 1.0 M NaOH to 250 mL of 1.0 M NaH<sub>2</sub>PO<sub>4</sub>. How many moles of HCl must be added to this buffer solution to change the pH by 0.18 units? If necessary, assume the total volume remains unchanged at 400 mL.
A) 0.025 mol HCl
B) 0.063 mol HCl
C) 0.082 mol HCl
D) 0.50 mol HCl
E) 1.0 mol HCl
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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Q1) A reaction has \(\Delta\)G = 10.0 kJ and \(\Delta\)G° = 15.0 kJ at a temperature of 50 °C. Calculate the value of the reaction quotient Q under these conditions.
A) 0.16
B) 9.1 × 10<sup>-5</sup>
C) 1.1 × 10<sup>4</sup>
D) 6.4
E) 6.0 × 10<sup>-6</sup>
Q2) In order for a process to be spontaneous,
A) the entropy of the system must increase.
B) the entropy of the surroundings must increase.
C) the entropy of the universe must decrease.
D) the entropy of the surroundings must decrease.
E) the entropy change of the surroundings plus the entropy change of the system must be positive.
Q3) The term microstate refers to the energy state of a single molecule in a system of many molecules.
A)True
B)False
Q4) State the second and third laws of thermodynamics.
Page 23
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Chapter 21: Electrochemistry: Chemical Change and Electrical Work
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Q1) Which one of the following statements relating to the glass electrode is correct?
A) The glass electrode detects hydrogen gas.
B) The glass of a glass electrode serves to conduct electrons.
C) When pH is measured, only a single electrode, the glass electrode, need be used.
D) The potential of the glass electrode varies linearly with the pH of the solution.
E) None of the above statements is correct.
Q2) Which, if any, of the following metals would not be capable of acting as a sacrificial anode when used with iron E°<sub>Fe</sub> = -0.44 V; all E° values refer to the M<sup>2+</sup>/M half-cell reactions.
A) manganese, Mn, E° = -1.18 V
B) cadmium, Cd, E° = -0.40 V
C) magnesium, Mg, E° = -2.37 V
D) zinc, Zn, E° = -0.76 V
E) All of these metals are capable of acting as sacrificial anodes with iron.
Q3) Write down equations representing the anode half-reaction, the cathode half-reaction, and the overall cell reaction for the lead-acid battery.
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Chapter 22: The Elements in Nature and Industry
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Q1) Electrometallurgy uses _________________ to separate a metal from its ore.
A) solid phase chemical properties
B) electrical processes
C) thermal processes
D) aqueous chemical processes
E) molten salt processes
Q2) What gas is produced at the anode in the Downs cell, in which molten NaCl is electrolyzed?
Q3) Carbon atoms in the carbon cycle spend most of their time in the oceans.
A)True
B)False
Q4) The most abundant element in the Earth's crust is
A) H.
B) Fe.
C) Al.
D) Si.
E) O.
Q5) What gas is produced at the anode in the Hall-Heroult process for the electrolytic production of aluminum?
Page 25
Q6) Briefly describe the three main pathways for nitrogen fixation.
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Chapter 23: The Transition Elements and Their Coordination Compounds
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Q1) Which one of the following has the ground state electron configuration [Ar]3d<sup>10</sup>4s<sup>1</sup>?
A) In<sup>+</sup>
B) Cd<sup>2+</sup>
C) Ag<sup>+</sup>
D) Ag
E) Cu
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) Octahedral complexes can exhibit geometric, optical, and linkage isomerism.
A)True
B)False
Q4) Why is the +2 oxidation state so common among transition elements?
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Q5) What is the difference between a coordination compound and a complex ion?
Chapter 24: Nuclear Reactions and Their Applications
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Q1) Most foodstuffs contain natural, radioactive isotopes.
A)True
B)False
Q2) Gamma rays are high energy electrons.
A)True
B)False
Q3) What is the mechanism by which control rods slow down the fission rate in a nuclear reactor?
Q4) After 4 half-lives, the fraction of a radioactive isotope which still remains is approximately one-eighth.
A)True
B)False
Q5) Write a complete, balanced equation to represent the alpha decay of radon-210.
Q6) An isotope with a low value of N/Z will generally decay through
A) .\(\alpha\) decay.
B) .\(\beta\) decay.
C) .\(\gamma\) decay.
D) electron capture.
E) spontaneous fission.
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Q7) Write a complete, balanced equation to represent the beta decay of thallium-207.