Chemistry for Life Sciences
Test Preparation
Course Introduction
Chemistry for Life Sciences introduces students to the fundamental chemical principles that underpin biological systems and processes. The course covers essential topics such as atomic structure, chemical bonding, acids and bases, thermodynamics, and chemical reactions, with a focus on their relevance to living organisms. Emphasis is placed on biomolecules including proteins, carbohydrates, lipids, and nucleic acids and how their chemical properties influence cellular function and metabolism.
Designed for students pursuing careers in health, medicine, or biological sciences, this course bridges the gap between chemistry and biology, equipping learners with the foundational knowledge necessary to understand complex biochemical pathways and the molecular basis of life.
Recommended Textbook
Chemistry A Molecular Approach 4th Edition by Nivaldo J. Tro
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25 Chapters
3452 Verified Questions
3452 Flashcards
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Page 2
Chapter 1: Matter, measurement, and Problem Solving
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169 Verified Questions
169 Flashcards
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Sample Questions
Q1) Identify the crystalline solid.
A) plastic
B) cloth
C) diamond
D) water
E) coffee
Answer: C
Q2) The outside temperature is 35°C.What is the temperature in K?
A) -238 K
B) 308 K
C) 95 K
D) 31 K
E) 63 K
Answer: B
Q3) Define matter.
Answer: Matter is anything that occupies space and has mass.
Q4) Define the law of the conservation of energy.
Answer: Energy is neither created nor destroyed.
Q5) Define energy.
Answer: Energy is the capacity to do work.
Page 3
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Chapter 2: Atoms and Elements
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Sample Questions
Q1) Ions differ in the number of A) electrons.
B) neutrons.
C) protons.
D) neutrons and protons.
E) electrons and protons.
Answer: A
Q2) How many electrons are in a neutral atom of chlorine-37?
A) 1
B) 17
C) 18
D) 37
Answer: B
Q3) How many atoms are in 5.00 moles of O?
A) 4.52 × 10<sup>23</sup> atoms
B) 1.51 × 10<sup>24</sup>atoms
C) 5.02 × 10<sup>23</sup>atoms
D) 3.01 × 10<sup>24</sup> atoms
E) 3.01 × 10<sup>23</sup> atoms
Answer: D
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Chapter 3: Molecules,compounds,and Chemical Equations
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Sample Questions
Q1) Give the name for KHSO<sub>3</sub>.
A) monopotassium sulfite
B) potassium sulfate
C) potassium hydrogen sulfate
D) potassium hydrogen sulfite
E) potassium disulfite
Answer: D
Q2) Methane and oxygen react to form carbon dioxide and water.What mass of water is formed if 9.6 g of methane reacts with 38.4 g of oxygen to produce 26.4 g of carbon dioxide?
A) 21.6 g
B) 26.4 g
C) 44.4 g
D) 48.0 g
Answer: A
Q3) Why aren't prefixes used in naming ionic compounds?
Answer: The charges on the ions dictate how many must be present to form a neutral unit.Molecular compounds do not have such constraints and therefore must use prefixes to denote the number of atoms present.
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Chapter 4: Chemical Quantities and Aqueous Reactions
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Sample Questions
Q1) Which pair of compounds is soluble in water?
A) Hg<sub>2</sub>I<sub>2</sub> and AgCl
B) CoS and K<sub>2</sub>S
C) NaI <sub> </sub> and Cu(NO<sub>3</sub>)<sub>2</sub>
D) LiNO<sub>3</sub> and CaCO<sub>3</sub>
Q2) How many moles of KF are required to make 200 mL of a 3.00 M solution?
A) 0.200 moles
B) 0.600 moles
C) 3.00 moles
D) 0.450 moles
E) 0.800 moles
Q3) Determine the molarity of a solution formed by dissolving 3.00 moles of NaCl in enough water to yield 4.00 L of solution.
A) 1.33 M
B) 2.00 M
C) 0.750 M
D) 3.00 M
E) 12.00 M
Q4) Describe the difference between complete ionic and net ionic equations.
Q5) Define an electrolyte.
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Chapter 5: Gases
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Sample Questions
Q1) How many molecules of N<sub>2</sub> are contained in a 10.0 L tank at 7.53 atm and 485 K?
A) 1.89 × 10<sup>24</sup> molecules
B) 1.14 × 10<sup>24</sup> molecules
C) 8.32 × 10<sup>24</sup> molecules
D) 4.89 × 10<sup>24</sup> molecules
E) 3.63 × 10<sup>24</sup> molecules
Q2) Determine the total volume of all gases (at STP)formed when 50.0 mL of TNT (C<sub>3</sub>H<sub>5</sub>(NO<sub>3</sub>)<sub>3</sub>,d = 1.60 g/mL,molar mass = 227.10 g/mol)reacts according to the following reaction. 4 C<sub>3</sub>H<sub>5</sub>(NO<sub>3</sub>)<sub>3</sub>(l) 6 N<sub>2</sub>(g)+ O<sub>2</sub>(g)+ 12 CO<sub>2</sub>(g)+ 10 H<sub>2</sub>O(g)
A) 4.93 L
B) 57.2 L
C) 29.6 L
D) 448 L
E) 175 L
Q3) Define effusion.
Q4) Explain what the term "mean free path" describes.How does it change with decreasing pressure?
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Chapter 6: Thermochemistry
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Sample Questions
Q1) Define specific heat capacity.
A) the quantity of heat required to raise the temperature of 1 mole of a substance by 1°C
B) the quantity of heat required to change a system's temperature by 1°C
C) the quantity of heat required to raise the temperature of 1 gram of a substance by 1°C
D) the quantity of heat required to raise the temperature of 1 kilogram of a substance by 1°F
E) the quantity of heat required to raise the temperature of 1 liter of a substance by 1°F
Q2) How much heat is absorbed when 40.00 g of C(s)reacts in the presence of excess SO<sub>2</sub>(g)to produce CS<sub>2</sub>(l)and CO(g)according to the following chemical equation? 5 C(s)+ 2 SO<sub>2</sub>(g) CS<sub>2</sub>(l)+ 4 CO(g) H° = 239.9 kJ
A) 159.8 kJ
B) 239.9 kJ
C) 798.7 kJ
D) 1918 kJ
Q3) Describe the energy changes that occur when a book is held 6 ft off the floor and then dropped.
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Chapter 7: The Quantum-Mechanical Model of the Atom
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Sample Questions
Q1) A sunburn is caused by overexposure to ________ radiation.
A) ultraviolet
B) gamma
C) microwave
D) x-ray
E) radio
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) Why don't we observe the wavelength of everyday macroscopic objects?
Q5) Define diffraction.
Q6) Describe how a neon light works.
9
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Chapter 8: Periodic Properties of the Elements
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Sample Questions
Q1) Identify the s-block element.
A) Cm
B) Ta
C) Na
D) F
E) None of the above
Q2) Which reaction below represents the electron affinity of S?
A) S(g) + e S (g)
B) S (g) + e S<sup>2</sup> (g)
C) S(g) S (g) + e
D) S (g) S(g) + e
E) S<sup>2</sup> (g) S (g) + e
Q3) Which atom in each group (I and II)has the smallest atomic radius? (I)Ca,Ti,Br (II)O,S,Se
A) Ca; O
B) Ca; Se
C) Br; O
D) Br; Se
Q4) Define paramagnetic.
Q5) Define ionization energy.
Page 10
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Chapter 9: Chemical Bonding I: Lewis Theory
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Sample Questions
Q1) Which element can expand its valence shell to accommodate more than eight electrons?
A) B
B) F
C) S
D) He
E) H
Q2) Choose the bond below that is the strongest.
A) C-Cl
B) C=O
C) C-Br
D) I-I
E) C C
Q3) How are electron affinity and electronegativity different?
Q4) Give the number of valence electrons for SO<sub>4</sub><sup>2-</sup>.
A) 32
B) 30
C) 34
D) 28
E) 36
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Chapter 10: Chemical Bonding Ii: Molecular Shapes,valence
Bond Theory,
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Sample Questions
Q1) Draw the molecular orbital diagram shown to determine which of the following is most stable.
A) C<sub>2</sub><sup>2</sup>
B) N<sub>2</sub><sup>2</sup>
C) B<sub>2</sub>
D) C<sub>2</sub><sup>2</sup>
E) B<sub>2</sub><sup>2</sup>
Q2) Place the following in order of decreasing dipole moment. I.cis-CHCl=CHCl II.trans-CHCl=CHCl III.cis-CHF=CHF
A) III > I > II
B) II > I > III
C) I > III > II
D) II > III > I
E) I = III > II
Q3) What is the molecular geometry of BrF<sub>4</sub> <sup>-</sup>?
A) seesaw
B) square planar
C) square pyramidal
D) pyramidal
E) trigonal bipyramidal
Page 12
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Chapter 11: Liquids,solids,and Intermolecular Forces
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Sample Questions
Q1) The normal boiling point of water is ________ at sea level.
A) 0°F
B) 32°C
C) 212<sup> </sup>K
D) 212°F
E) 273°F
Q2) In liquid methanol, CH<sub>3</sub>OH<sub> </sub> <sub> </sub> <sub> </sub> <sub> </sub>
Which intermolecular forces are present?
A) Dispersion, hydrogen bonding and dipole-dipole forces are present.
B) Only dipole-dipole and hydrogen bonding forces are present.
C) Only dispersion and ion-dipole forces are present.
D) Only hydrogen bonding forces are present.
Q3) Sketch the phase diagram of benzene.Make sure to label the axes and the different phases of benzene.Use the physical data provided below.
melting point = 279 K
boiling point = 353 K
T<sub>c</sub> = 562 K
P<sub>c</sub> = 48.4 atm
Triple Point = 0.05 atm,279 K
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Chapter 12: Solids and Modern Materials
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Sample Questions
Q1) Identify the element with the smallest band gap.
A) lead
B) silicon
C) germanium
D) carbon
E) tin
Q2) Identify the type of solid for AgCl.
A) metallic atomic solid
B) ionic solid
C) nonbonding atomic solid
D) molecular solid
E) networking atomic solid
Q3) Identify the compound that has fluorite 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.
Page 14
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Chapter 13: Solutions
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Sample Questions
Q1) Fluids used for an intravenous transfusion must be ________ with bodily fluids.
A) isosmotic
B) hyperosmotic
C) hyposmotic
D) neosmotic
E) magnosmotic
Q2) A solution is prepared by adding 30.00 g of lactose (milk sugar)to 110.0 g of water at 55° The partial pressure of water above the solution is ________ torr.The vapor pressure of pure water at 55°C is 118.0 torr.The molar mass of lactose is 342.3 g/mol.
A) 1.670
B) 94.1
C) 169.4
D) 116.3
E) 92.7
Q3) Which of the following solutions will have the lowest freezing point?
A) 0.010 m NH<sub>4</sub>Cl
B) 0.010 m Li<sub>2</sub>CO<sub>3</sub>
C) 0.035 m CH<sub>3</sub>CH<sub>2</sub>CH<sub>2</sub>OH
D) 0.015 m CaI<sub>2</sub>
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Chapter 14: Chemical Kinetics
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Sample Questions
Q1) Given the following rate law,how does the rate of reaction change if the concentration of Y is doubled? Rate = k[X]<sup>4</sup>[Y]<sup>3</sup>
A) The rate of reaction will increase by a factor of 6.
B) The rate of reaction will decrease by a factor of 4.
C) The rate of reaction will increase by a factor of 8.
D) The rate of reaction will increase by a factor of 4.
E) The rate of reaction will remain unchanged.
Q2) The first-order decay of radon has a half-life of 3.823 days.How many grams of radon decompose after 5.55 days if the sample initially weighs 100.0 grams?
A) 83.4 g
B) 16.6 g
C) 50.0 g
D) 36.6 g
E) 63.4 g
Q3) Define half-life.
Q4) Explain how the order of a reaction can be determined.
Q5) Explain what the exponential factor in the Arrhenius equation represents.
Q6) What function do enzymes serve?
Q7) Define the frequency factor.
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Chapter 15: Chemical Equilibrium
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Sample Questions
Q1) Consider the following reaction at equilibrium.What will happen if the volume increased? 4 FeS<sub>2</sub>(s)+ 11 O<sub>2</sub>(g) 2 Fe<sub>2</sub>O<sub>3</sub>(s)+ 8 SO<sub>2</sub>(g)
A) The equilibrium constant will increase.
B) The equilibrium will change in the direction of the reactants.
C) The equilibrium will change in the direction of the products.
D) No change in equilibrium is observed.
E) The equilibrium constant will decrease.
Q2) Consider the following reaction,equilibrium concentrations,and equilibrium constant at a particular temperature.Determine the equilibrium concentration of SO<sub>2</sub>(g). 2 SO<sub>2</sub>(g)+ O<sub>2</sub>(g) 2 SO<sub>3</sub>(g)K<sub>c</sub> = 1.7 × 10<sup>8</sup> [SO<sub>3</sub>]<sub>eq</sub> = 0.0034 M [O<sub>2</sub>]<sub>eq</sub> = 0.0018 M
A) 2.8 × 10<sup>13</sup> M
B) 1.88 M
C) 6.1 × 10<sup>-6</sup> M
D) 1.0 × 10<sup>3</sup> M
E) 1.4 M
Q3) Why is an equilibrium constant unitless?
Q4) Why aren't solids or liquids included in an equilibrium expression?
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Chapter 16: Acids and Bases
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Sample Questions
Q1) Which of the following is an Arrhenius acid?
A) H<sub>2</sub>SO<sub>3</sub>
B) CsOH
C) NH<sub>2</sub>CH<sub>3</sub>
D) CH<sub>3</sub>CI<sub>3</sub>
E) LiCl
Q2) What is the pH of a 0.100 M NH<sub>3</sub> solution that has K<sub>b</sub> = 1.8 × 10<sup>-5</sup>? The equation for the dissociation of NH<sub>3</sub> is: NH<sub>3</sub>(aq)+ H<sub>2</sub>O(l) NH<sub>4</sub><sup>+</sup>(aq)+ OH<sup>-</sup>(aq)
A) 1.87
B) 2.87
C) 11.13
D) 10.13
Q3) Which of the following is a weak acid?
A) HClO<sub>4</sub>
B) H<sub>2</sub>SO<sub>4</sub>
C) HBr
D) HCN
E) HNO<sub>3</sub>
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Chapter 17: Aqueous Ionic Equilibrium
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Sample Questions
Q1) Stalactites and stalagmites form as ________ precipitates out of the water evaporating in underground caves.
A) hydrochloric acid
B) potassium hydroxide
C) lithium bromide
D) calcium carbonate
E) sodium bicarbonate
Q2) Calculate the pH of a solution formed by mixing 150.0 mL of 0.10 M HC<sub>7</sub>H<sub>5</sub>O<sub>2</sub> with 100.0 mL of 0.30 M NaC<sub>7</sub>H<sub>5</sub>O<sub>2</sub>.The K<sub>a</sub> for HC<sub>7</sub>H<sub>5</sub>O<sub>2</sub> is 6.5 × 10<sup>-5</sup>.
A) 4.19
B) 9.69
C) 4.49
D) 4.31
E) 10.51
Q3) Give the name of the compound that is in antifreeze and is toxic to pets.
Q4) Explain the common ion effect with respect to molar solubility.
Q5) Describe the information necessary (and why)to choose a good indicator for a weak acid/strong base titration.
Page 19
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Chapter 18: Free Energy and Thermodynamics
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Sample Questions
Q1) Identify the element with the highest standard free energy of formation.
A) Ba(s)
B) Ca(s)
C) Li(s)
D) K(s)
E) All elements have a value of zero.
Q2) For a given reaction, H = -19.9 kJ/mol and S = -55.5 J/Kmol.The reaction will have G = 0 at ________ K.Assume that H and S do not vary with temperature.
A) 359
B) 2789
C) 298
D) 2.79
E) 0.359
Q3) Which of the following processes has a S < 0?
A) ethanol freezes
B) isopropyl alcohol condenses
C) 2-propanol (g, at 555 K) 2-propanol (g, at 400 K)
D) carbon dioxide(g) carbon dioxide(s)
E) All of the above processes have a S < 0.
Q4) Define allotrope.
Page 20
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Chapter 19: Electrochemistry
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Sample Questions
Q1) Identify oxidation.
A) increase in oxidation number
B) loss of electrons
C) gain of protons
D) loss of neutrons
E) both A and B
Q2) Identify a characteristic that does not describe a standard hydrogen electrode.
A) 1 M HCl solution with hydrogen gas
B) 1 atmosphere
C) zinc electrode
D) inert platinum electrode
E) potential of zero
Q3) What element is being oxidized in the following redox reaction?
Cr(OH)<sub>4</sub><sup>-</sup>(aq)+
CrO<sub>4</sub><sup>2-</sup>(aq)+ Cl<sup>-</sup>(aq)
A) Cr
B) O
C) H
D) Cl
Q4) What is electrolysis?
ClO<sup>-</sup>(aq)
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Chapter 20: Radioactivity and Nuclear Chemistry
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Sample Questions
Q1) Which of the following statements is TRUE?
A) Positrons are similar in ionizing power and penetrating power to beta particles.
B) A positron is the antiparticle of the electron.
C) Beta decay occurs when a neutron changes into a proton while emitting an electron.
D) An alpha particle is a helium 2+ ion.
E) All of the above are true.
Q2) The half-life of cobalt-60 is 5.20 yr.How many milligrams of a 2.000-mg sample remain after 6.55 years?
A) 0.837
B) 3.23 × 10<sup>-15</sup>
C) 4.779
D) 1.588
E) 1.163
Q3) List two problems associated with nuclear power.
Q4) Describe what is meant by the term "valley of stability"?
Q5) Define chain reaction in terms of the fission of uranium nucleus.
Q6) Why is an alpha emitter much more harmful if is ingested than when applied to the skin?
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Chapter 21: Organic Chemistry
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Sample Questions
Q1) Which of the following names is correct?
A) 2-methyl-1-butene
B) 3,3-diethyl-4-propene
C) 2-ethyl-3-pentene
D) 1-propyl-2-pentene
E) None of the above is correct.
Q2) Identify the formula for an alkyne.
A) C<sub>n</sub>H<sub>2n+2</sub>
B) C<sub>n</sub>H<sub>2n-2</sub>
C) C<sub>n</sub>H<sub>2n</sub>
D) C<sub>n</sub>H<sub>2n-3</sub>
E) C<sub>n</sub>H<sub>2n+3</sub>
Q3) Which of the following compounds exhibits optical isomerism?
A) sec-butylbenzene
B) CH<sub>2</sub>I<sub>2</sub>
C) CF<sub>2</sub>BrI
D) CHF<sub>2</sub>I
E) isobutylbenzene
Q4) What is meant by the term "structural isomer"? Draw a structural isomer of CH<sub>3</sub>-CH<sub>2</sub>-CH<sub>2</sub>-CH<sub>3</sub>.
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Chapter 22: Biochemistry
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Sample Questions
Q1) Which one of the following amino acids contains a polar side chain?
A) valine
B) glycine
C) tyrosine
D) leucine
E) phenylalanine
Q2) Which of the following is a type of nucleic acid?
A) DNA
B) amino acid
C) tripeptide
D) carbohydrate
E) lipid
Q3) Identify a base that is NOT in RNA.
A) adenine
B) cytosine
C) thymine
D) guanine
E) uracil
Q4) What is a codon?
Q5) Why are lipids well-suited as structural components of cell membranes?
Page 24
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Chapter 23: Chemistry of the Nonmetals
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Sample Questions
Q1) Determine the oxidation state of nitrogen in LiNO<sub>2</sub><sup> </sup>.
A) +1
B) -2
C) +5
D) 0
E) +3
Q2) Describe the Frasch process for obtaining sulfur.
Q3) What is the oxidation number of oxygen in H<sub>2</sub>O<sub>2</sub>?
A) -4
B) +2
C) -3
D) -1
E) +1
Q4) Identify the form of carbon that is used for gems.
A) coal
B) charcoal
C) gasoline
D) diamond
E) fullerene
Q5) Why are phosphate compounds added to detergents?
Page 25
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Chapter 24: Metals and Metallurgy
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Sample Questions
Q1) Pyrometallurgy is the
A) refining of metal ores using reactions with aqueous solutions.
B) refining of metal ores using heat.
C) refining of metal ores using reduction reactions.
D) forming metal parts using heat and small crystals of metal.
E) forming metal parts using electricity.
Q2) Why is zinc used to coat steel objects?
Q3) An interstitial alloy contains carbon in one-fourth of the octahedral holes of a closest packed metal,M.What is the formula of this alloy?
A) M<sub>4</sub>C
B) MC<sub>4</sub>
C) MC
D) M<sub>2</sub>C
E) MC<sub>2</sub>
Q4) Why does a "two-phase" structure occur in a substitutional alloy?
Q5) Which substance is the best conductor of electricity?
A) astatine
B) diamond
C) tungsten
D) xenon
Page 26
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Chapter 25: Transition Metals and Coordination Compounds
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Sample Questions
Q1) How many unpaired electrons would you expect for the complex ion: [FeCl<sub>6</sub>]<sup>4-</sup>?
A) 2
B) 6
C) 4
D) 3
E) 0
Q2) Identify the transition metal that is used in hemoglobin synthesis in the human body.
A) potassium
B) zinc
C) copper
D) manganese
E) chromium
Q3) Choose the polydentate ligand from the substances below.
A) oxalate ion
B) EDTA
C) nitrite ion
D) hydroxide ion
E) carbon monoxide
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