Chemistry for Life Sciences
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Course Introduction
Chemistry for Life Sciences introduces students to fundamental chemical principles with an emphasis on their application in biological systems. The course covers atomic and molecular structure, chemical bonding, acids and bases, thermodynamics, and reaction kinetics, all contextualized to processes relevant in living organisms. Students explore the chemistry of water, organic molecules, pH regulation, enzyme action, energy production, and metabolic pathways. The curriculum is designed to build a solid foundation for further studies in biology, biochemistry, medicine, and related health sciences, providing practical laboratory experience to reinforce theoretical knowledge.
Recommended Textbook Principles of Chemistry The Molecular Science 1st Edition by John W. Moore
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Chapter 1: The Nature of Chemistry
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Sample Questions
Q1) Which characteristic below best fits the description of a solid?
A) large distances between the molecules
B) molecules that are close together but are moving past one another
C) highly disordered molecules
D) rapid molecular motion
E) highly ordered molecules
Answer: E
Q2) The freezing point and boiling point of water are often used to calibrate thermometers. Give those temperatures in degrees Celsius.
A) 0 and 100
B) 273 and 373
C) 32 and 212
D) 0 and 373
E) 100 and 273
Answer: A
Q3) A mixture that is nonuniform in composition is a(n) _____________ mixture. Answer: heterogeneous
Q4) How many atoms are in a diatomic molecule?.
Answer: two
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Chapter 2: Atoms and Elements
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Q1) Which statement about electrons is true?
A) Electrons are found in the nucleus of the atom.
B) Electrons are attracted to negatively charged electrodes.
C) All atoms have electrons as part of their structure.
D) Electrons have much more mass than any atom.
E) Electrons are positively charged.
Answer: C
Q2) <sup>14</sup>O, <sup>16</sup>O, <sup>18</sup>O are called _____________.
Answer: isotopes
Q3) Which metric prefix means 1000?
A) pico
B) micro
C) milli
D) kilo
E) nano
Answer: D
Q4) There are _____________ centiliters in a liter.
Answer: 100
Q5) Give an example of an Alkaline Earth Metal.
Answer: Be, Mg, Ca, Sr, Ba, Ra
4
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Chapter 3: Chemical Compounds
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Sample Questions
Q1) Determine the percent potassium present in K<sub>2</sub>SO<sub>4</sub>.
A) 22.4%
B) 28.9%
C) 44.9%
D) 57.9%
E) None of these
Answer: C
Q2) An organic compound has an empirical formula of C<sub>2</sub>H<sub>3</sub>O and had an approximate molecular weight of 86. What is its molecular formula?
A) C<sub>2</sub>H<sub>3</sub>O
B) C<sub>4</sub>H<sub>6</sub>O<sub>2</sub>
C) C<sub>6</sub>H<sub>8</sub>O<sub>3</sub>
D) C<sub>8</sub>H<sub>12</sub>O<sub>4</sub>
E) C<sub>10</sub>H<sub>14</sub>O<sub>5</sub>
Answer: B
Q3) Name the compound with the formula VCl<sub>3</sub>. Answer: vanadium(III) chloride
Q4) Write the formula for chromium(III) sulfate. Answer: Cr<sub>2</sub>(SO<sub>4</sub>)<sub>3</sub>
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Chapter 4: Quantities of Reactants and Products
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Q1) The efficiency of a particular synthesis method is evaluated by determining the: A) molecular weight of the product.
B) stoichiometric coefficients.
C) limiting reactant.
D) theoretical yield.
E) percent yield.
Q2) Which of the following cannot be determined from a balanced chemical equation?
A) The relative mass of each reactant and product.
B) The number of moles of reactants and products.
C) The number of molecules of reactants and products.
D) The number of atoms of each element reacting.
E) Whether the reaction will proceed as written.
Q3) The complete combustion of a hydrocarbon produces 90.36 g of CO<sub>2</sub> and 46.25 g of H<sub>2</sub>O. What is the empirical formula of the hydrocarbon?
A) CH
B) CH<sub>2</sub>
C) C<sub>2</sub>H<sub>5</sub>
D) C<sub>3</sub>H<sub>8</sub>
E) C<sub>3</sub>H<sub>4</sub>
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Chapter 5: Chemical Reactions
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Q1) How many moles of NaOH are present in 25.0 mL of a 0.1000 M NaOH solution?
A) 100 mol
B) 2.50 × 10<sup>-3</sup> mol
C) 0.100 mol
D) 2.50 mol
E) 25.0 mol
Q2) What is the oxidation number of O in Fe<sub>2</sub>O<sub>3</sub>?
A) +2
B) -2
C) -3
D) -5
E) +5
Q3) Which statement about bases is true?
A) Bases increase the hydroxide ion concentration of water when dissolved in it.
B) Bases increase the hydronium ion concentration of water when dissolved in it.
C) Bases turn phenolphthlalein colorless.
D) Bases react with limestone to produce gas bubbles.
E) Bases react with many metals to produce a flammable gas.
Q4) How many moles of ions are in a 250.0 mL sample of 0.150 M NaCl?
Q5) When an element is _____________, it loses electrons.
Page 7
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Chapter 6: Energy and Chemical Reactions
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Q1) The temperature of 3.50 kg of water is raised by 1.17°C when 1.00 g of hydrazine N<sub>2</sub>H<sub>4</sub> is burned in a bomb calorimeter. The calorimeter has a heat capacity of 883 J/°C. How much heat is given off by the sample? The specific heat of water = 4.184 J g<sup>-1</sup> °C<sup>-1</sup>.
A) 0.944 kJ
B) 16.3 kJ
C) 17.1 kJ
D) 18.2 kJ
E) 21.5 kJ
Q2) Determine the incorrect relationship given below.
A) 1 µJ = 1 × 10<sup>-6</sup> J
B) 1000 cal = 1 kcal
C) 44.0 kJ = 1.05 × 10<sup>4</sup> cal
D) 1000 J = 1 kJ
E) 80.0 cal/g = 312 J/g
Q3) How much energy in kilojoules is required to vaporize a 25.0 g sample of water at 100.0°C? The heat of vaporization of water = 2260 J g<sup>-1</sup> at 100°C.
Q4) Enthalpy change is equal to heat transfer at constant _______________.
Q5) Energy of motion is called _____________.
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Chapter 7: Electron Configurations and the Periodic Table
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Q1) Which statement about the formation of solid sodium chloride from sodium metal atoms and chlorine atoms is false?
A) Sodium atoms transfer electrons to chlorine atoms.
B) The sodium cation has an octet of outer electrons.
C) Sodium chloride contains alternating atoms in a crystal lattice.
D) The chloride anion has an octet of outer electrons.
E) The formation of sodium chloride from the elements releases energy to the surroundings.
Q2) Which word or phrase least applies to the quantum number represented by the symbol l?
A) subshell
B) s, p, d, f
C) orbital type
D) orientation
E) shape
Q3) As one moves closer to the nucleus, what happens to the value of the principal quantum number?
Q4) Atomic radii _____________ going from left to right across a row of the periodic table.
Q5) Each electron in an atom or ion must have a unique set of four _____________.
Page 9
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Chapter 8: Covalent Bonding
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Q1) Benzene is an example of a resonance hybrid because it contains _____________ electrons.
Q2) What is the formal charge on carbon in CO?
A) -2
B) -1
C) 0
D) +1
E) +2
Q3) Which represents a non-polar covalent bond?
A) H-O
B) C-N
C) C-C
D) Li-F
E) S-O
Q4) Hydrocarbons with only single bonds are called _____________.
Q5) Hydrocarbons that contain double bonds are called _____________.
Q6) Oxygen atoms contribute _____________ electrons to a Lewis Dot structure.
Q7) The correctly drawn Lewis dot structure for linear H<sub>2</sub>CO<sub>2</sub> contains _____________ lone electron pairs, _____________ single bonds and _____________ double bonds.
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Chapter 9: Molecular Structure
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Q1) Which accurately describes a molecule with all polar bonds?
A) It is always a polar molecule.
B) It is never a polar molecule.
C) It may be a polar molecule.
D) It may be an ionic molecule.
E) It cannot be a nonpolar molecule.
Q2) Use the VSEPR model to predict the electron-pair geometry of O<sub>3</sub>.
A) linear
B) tetrahedral
C) triangular bipyramidal
D) bent
E) triangular planar
Q3) For which molecule will the electron-pair geometry be different than the molecular geometry?
A) NH<sub>4</sub><sup>+</sup>
B) CO
C) CH<sub>4</sub>
D) BH<sub>3</sub>
E) H<sub>2</sub>O
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Chapter 10: Gases and the Atmosphere
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Q1) If a 2.0-liter sample of gas experiences a decrease in pressure from 1.74 atm to 0.555 atm at 25°C, what is the resulting volume (in L) at 25°C?
A) 0.48 L
B) 0.64 L
C) 1.9 L
D) 6.3 L
E) 20 L
Q2) Which statement is false?
A) Atmospheric gases are an example of gases mixing completely.
B) An exhaled breath of air will disperse throughout the entire atmosphere.
C) A gas inside a container will remain inside if the container is opened.
D) Atmospheric pressure is a result of gas molecules exerting pressure on their surroundings.
E) An air pump demonstrates that gases can be compressed.
Q3) Which is a mathematical representation of Charles's Law?
A) P/T = constant
B) PV = constant
C) V × T = constant
D) n × T = constant
E) V/T = constant
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Chapter 11: Liquids, Solids, and Materials
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Q1) A positive slope for a region of a heating curve indicates that ____ in that region.
A) no energy is being absorbed by the system
B) energy is being given off by the system, but it cannot be measured
C) energy is being absorbed by the system and is being used for a phase change
D) additional data is needed to explain this observation
E) energy is being absorbed by the system and is being used to increase the temperature
Q2) Which choice is an example of a metallic solid?
A) NaCl
B) crystal
C) iron
D) quartz
E) glass
Q3) What is the correct order for the surface tension of the following substances?
A) CHCl<sub>3</sub> < H<sub>2</sub>O < C<sub>8</sub>H<sub>18</sub>
B) H<sub>2</sub>O < CHCl<sub>3</sub> < C<sub>8</sub>H<sub>18</sub>
C) CHCl<sub>3</sub> < C<sub>8</sub>H<sub>18</sub> < H<sub>2</sub>O
D) C<sub>8</sub>H<sub>18</sub> < CHCl<sub>3</sub> < H<sub>2</sub>O
E) C<sub>8</sub>H<sub>18</sub> < H<sub>2</sub>O < CHCl<sub>3</sub>
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Chapter 12: Chemical Kinetics: Rates of Reactions
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Q1) The function of an automobile catalytic converter is to convert
A) CO to CO<sub>2</sub> and NO to N<sub>2</sub>.
B) C to CO and NO to N<sub>2</sub>.
C) CO to CO<sub>2</sub> and O<sub>2</sub> to O<sub>3</sub>.
D) CO to CO<sub>2</sub> and NO to N<sub>2</sub>O.
E) CO to CO<sub>2</sub> and N<sub>2</sub> to NO.
Q2) A reaction displays zero-order kinetics for its single reactant. It therefore follows that a plot of ____ versus time is linear, and that the slope of this plot = ____.
A) [reactant]; -k
B) [reactant]; k
C) 1/[reactant]; -k
D) 1/[reactant]; k
E) ln[reactant]; -k
Q3) The substrate in an enzyme-controlled biological reaction is a
A) molecule serving as a rigid support for the enzyme.
B) molecule that is acted on by the catalyst.
C) species necessary for the proper functioning of the enzyme.
D) biological molecule used for storage of energy.
E) catalyst for that chemical reaction.
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Chapter 13: Chemical Equilibrium
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Q1) For an exothermic reaction, an increase in temperature
A) always means an increase in K<sub>c</sub>
B) always means a decrease in K<sub>c</sub>
C) the reaction will become more product-favored at higher temperatures
D) the reaction will become more reactant-favored at higher temperatures
E) both b and d
Q2) The equilibrium constant for a particular chemical reaction is
A) dependent on the total pressure of the reaction
B) dependent on the temperature of the reaction
C) a numerical value between -¥ and +¥
D) unreliably determined if variable amounts of either solid or liquid reactants or products are present
E) both b and c
Q3) A chemical reaction reaches equilibrium when
A) both the forward and reverse reactions stop
B) the forward and reverse reactions both occur at the same rate
C) all of the limiting reactant has been used up
D) all of the limiting reactant has been used up and all of the limiting product has been created
E) the actual yield of the reaction equals the theoretical yield
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Chapter 14: The Chemistry of Solutes and Solutions
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Q1) Drinkable water can be produced from seawater by the process of ____, using a membrane which is permeable to ____ but not to ____.
A) osmosis; water; ions
B) reverse osmosis; water; ions
C) reverse osmosis; sodium ions; chloride ions
D) osmosis; ions; water
E) reverse osmosis; ions; water
Q2) A concentrated antifreeze solution contains 580. g ethylene glycol mixed with 540. g of water. The weight percent of water in this solution is
A) 27.0 %
B) 48.2 %
C) 51.8 %
D) 93.1 %
E) 107 %
Q3) Explain what is meant by a. osmosis.
b. osmotic pressure.
c. reverse osmosis.
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Chapter 15: Acids and Bases
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Q1) Which statement is not correct?
A) a Lewis acid is a substance that can accept a pair of electrons to form a new bond
B) a Lewis base is a substance that can donate a pair of electrons to form a new bond
C) Al(OH)<sub>3</sub> is an amphoteric substance that can form either a positive or negative ion
D) Neutral molecules cannot act as Lewis acids
E) Ag(NH<sub>3</sub>)<sub>2</sub><sup>+</sup> is a complex ion, formed from a Ag<sup>+</sup> ion (Lewis acid) and two NH<sub>3</sub> molecules (Lewis bases)
Q2) A solution is not neutral. Which one of these statements is true?
A) [H<sub>3</sub>O<sup>+</sup>] = 1.0 × 10<sup>-7</sup> M
B) [H<sub>3</sub>O<sup>+</sup>] = [OH<sup>-</sup>]
C) [H<sub>3</sub>O<sup>+</sup>][OH<sup>-</sup>] = 1.0 × 10<sup>-7</sup>
D) [OH<sup>-</sup>] = 1.0 × 10<sup>-7</sup> M
E) [H<sub>3</sub>O<sup>+</sup>][OH<sup>-</sup>] = 1.0 × 10<sup>-14</sup>
Q3) A solution of acetic acid (K<sub>a</sub> = 1.8 × 10<sup>-5</sup>) has a pH exactly 2 higher than a solution of HCl(aq) of the same concentration. What is that concentration?
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Chapter 16: Additional Aqueous Equilibria
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Q1) An alkaline buffer would best be made using
A) a weak acid with a pK<sub>a</sub> > 7.00 or K<sub>a</sub> < 1.00 × 10<sup>-7</sup> and its conjugate base
B) a weak acid with a pK<sub>a</sub> < 7.00 or K<sub>a</sub> > 1.00 × 10<sup>-7</sup> and its conjugate base
C) a strong acid and a strong base
D) a weak acid with a K<sub>a</sub> > 1.00 × 10<sup>14</sup>
E) a strong base and its conjugate acid
Q2) A student attempts to dissolve a small amount of CuCO<sub>3</sub> in 0.0100 M NaOH. Despite vigorous stirring, the material does not apparently dissolve. However, after a while the student notices that the solid has changed appearance. When the pH of the solution is measured, it is found to have gone down. Explain these observations. Was the original goal of the experiment achieved?
Q3) In which liquid or solution would CuI(s) have the greatest solubility?
A) 0.1 M aqueous ammonia
B) distilled water
C) 0.1 M aqueous copper(I) nitrate
D) saturated aqueous calcium hydroxide
E) 0.1 M aqueous sodium iodide
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Chapter 17: Thermodynamics: Directionality of Chemical Reactions
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Q1) Compression of a gas is
A) reactant-favored, since energy in the compressed gas is more dispersed.
B) reactant-favored, since matter in the compressed gas is more dispersed.
C) reactant-favored, since energy in the compressed gas is less dispersed.
D) product-favored, since energy in the compressed gas is less dispersed.
E) product-favored, since energy in the compressed gas is more dispersed.
Q2) Nature favors exothermic reactions because after such a reaction
A) the energy previously concentrated in a few particles is now dispersed over more particles in the system.
B) the energy previously concentrated in a few particles is now dispersed over more particles in the surroundings.
C) the energy previously concentrated in a few particles is now dispersed over more particles in both the system and the surroundings.
D) the energy previously held in many particles is now concentrated in a few, resulting in a temperature rise.
E) the energy previously held in many particles is now concentrated in a few, resulting in a temperature fall.
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Chapter 18: Electrochemistry and Its Applications
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Q1) In the anode compartment of a simple electrochemical cell the electrode is being ____, and ____ are flowing in from the salt bridge.
A) oxidized; anions
B) oxidized; cations
C) oxidized; electrons
D) reduced; cations
E) reduced; anions
Q2) Is the part of a flashlight battery that is marked "+" the anode or the cathode? Explain why it is labeled in that way.
Q3) A secondary cell is one which ____; this is possible because ____.
A) is driven by a primary cell; the cell reactions are coupled
B) is disposable without harming the environment; no mercury is used
C) is rechargeable; the cell reaction is reversible
D) has double the lifetime of a primary cell; two primary cells are placed in parallel
E) has double the voltage of a primary cell; two primary cells are placed in series
Q4) Explain why galvanizing steel chain-link fencing is a more effective strategy to prevent corrosion than simply painting it.
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Chapter 19: Nuclear Chemistry
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Q1) The most penetrating type of radioactivity is ____, and the least penetrating is
A) alpha particles; gamma rays
B) beta particles; alpha particles
C) beta particles; gamma rays
D) gamma rays; alpha particles
E) gamma rays; beta particles
Q2) All of these are common modes of radioactive decay except
A) positron capture
B) positron emission
C) electron capture
D) beta emission
E) alpha emission
Q3) Which of these is the smallest contributor to background radiation exposure in the U.S.?
A) cosmic radiation
B) radon
C) x-rays
D) uranium
E) nuclear wastes
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