Survey of Chemistry Test Preparation
Course Introduction
Survey of Chemistry offers a broad introduction to the fundamental concepts of chemistry, including atomic structure, chemical bonding, stoichiometry, states of matter, solutions, acids and bases, chemical reactions, and key principles of organic and biochemistry. Designed for students with diverse academic backgrounds, this course emphasizes practical applications of chemistry in everyday life, environmental issues, and health sciences. Through lectures, demonstrations, and laboratory experiments, students develop problem-solving skills and scientific literacy, laying the groundwork for further study in the physical and life sciences.
Recommended Textbook
Chemistry 10th Edition by Raymond Chang
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25 Chapters
2734 Verified Questions
2734 Flashcards
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Page 2
Chapter 1: Chemistry: The Study of Change
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153 Verified Questions
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Sample Questions
Q1) Identify the following as a physical or chemical change: Water is broken down into hydrogen and oxygen.
Answer: Chemical
Q2) Classify the following as either a physical or a chemical property: Ice melts at 0<sup>\(\circ\)</sup>C.
Answer: Physical
Q3) Suppose a house has a floor area of 2,250 square feet. What is this area in units of square centimeters?
A)2.42 cm<sup>2</sup>
B)2.09 * 10<sup>6</sup> cm<sup>2</sup>
C)5.02 * 10<sup>4</sup> cm<sup>2</sup>
D)6.86 * 10<sup>4</sup> cm<sup>2</sup>
E)101 cm<sup>2</sup>
Answer: B
Q4) An automobile engine has a piston displacement of 1,600 cm<sup>3</sup>. Express this volume in liters.
Answer: 1.6 L
Q5) What are the three states of matter?
Answer: Solid, liquid, and gas
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Chapter 2: Atoms, Molecules, and Ions
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Sample Questions
Q1) Give the number of protons (p), electrons (e), and neutrons (n)in one atom of nickel-62.
A)28 p, 28 e, 28 n
B)28 p, 28 e, 34 n
C)62 p, 28 e, 28 n
D)62 p, 62 e, 28 n
Answer: B
Q2) What is the formula for the binary compound formed by potassium and nitrogen? A)KN
B)K<sub>2</sub>N
C)NK<sub>2</sub>
D)K<sub>3</sub>N
E)NK<sub>3</sub>
Answer: D
Q3) What are the seven elements that naturally occur as diatomic molecules? Answer: Hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, iodine
Q4) Give the formula for potassium oxide. Answer: K<sub>2</sub>O
Q5) Write the formula of ammonium chlorate. Answer: NH<sub>4</sub>ClO<sub>3</sub>
Page 4
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Chapter 3: Mass Relationships in Chemical Reactions
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Sample Questions
Q1) What is the coefficient of H<sub>2</sub>O when the following equation is properly balanced with the smallest set of whole numbers? ___ Na + ___ H<sub>2</sub>O \(\rarr\) ___ NaOH + ___ H<sub>2</sub>
A)1
B)2
C)3
D)4
E)5
Answer: B
Q2) Formaldehyde has the formula CH<sub>2</sub>O. How many molecules are there in 0.11 g of formaldehyde?
A)6.1 * 10<sup>-27</sup>
B)3.7 * 10<sup>-3</sup>
C)4
D)2.2 * 10<sup>21</sup>
E)6.6 * 10<sup>22</sup>
Answer: D
Q3) How many Mg atoms are present in 170 g of Mg?
Answer: 4.2 * 10<sup>24</sup>
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Page 5
Chapter 4: Reactions in Aqueous Solution
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Sample Questions
Q1) One method of determining the concentration of hydrogen peroxide (H<sub>2</sub>O<sub>2</sub>)in a solution is through titration with iodide ion. The net ionic equation for this reaction is H<sub>2</sub>O<sub>2</sub> + 2I<sup>-</sup> +2H<sup>+</sup> \(\rarr\) I<sub>2</sub> + 2H<sub>2</sub>O
A 50.00 mL sample of a hydrogen peroxide solution is found to react completely with 37.12 mL of a 0.1500 M KI solution. What is the concentration of hydrogen peroxide in the sample?
A)5.568 * 10<sup>-2</sup> M
B)0.2227 M
C)0.1010 M
D)0.4041 M
E)0.1114 M
Q2) Identify the reducing agent in the following reaction. 2KBr + F<sub>2</sub> \(\rarr\)Br<sub>2</sub> + 2KF
Q3) Give an example of a triprotic acid.
Q4) Write balanced molecular and net ionic equations for the reaction that would occur between Al(s)and Co(NO<sub>3</sub>)<sub>2</sub>(aq).Be sure to include the correct states in your final equations.If no reaction is expected, write "no reaction."
Q5) Name and give the formulas for six strong acids.
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Chapter 5: Gases
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Sample Questions
Q1) The pressure of a gas sample was measured to be 489 mmHg.Which of the following is not an equivalent statement of that pressure? (1 atm = 1.01325 * 10<sup>5</sup> Pa)
A)65.2 kPa
B)6.52 * 10<sup>4</sup> Pa
C)489 torr
D)0.811 atm
Q2) A small bubble rises from the bottom of a lake, where the temperature and pressure are 4°C and 3.0 atm, to the water's surface, where the temperature is 25°C and the pressure is 0.95 atm.Calculate the final volume of the bubble if its initial volume was 2.1 mL.
A)0.72 mL
B)6.2 mL
C)41.4 mL
D)22.4 mL
E)7.1 mL
Q3) A particular coal sample contains 2.32% S.When the coal is burned, the sulfur is converted to sulfur dioxide gas. What volume of SO<sub>2</sub>(g), measured at 25°C and 749 mmHg, is produced by burning 2.0 * 10<sup>6</sup> lb of this coal? (1 lb = 454 g)
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Page 7
Chapter 6: Thermo-Chemistry
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Sample Questions
Q1) A piece of copper with a mass of 218 g has a heat capacity of 83.9 J/°C. What is the specific heat of copper?
A)0.385 J/g·°C
B)1.83 * 10<sup>4</sup> J/g·°C
C)2.60 J/g·°C
D)1.32 J/g·°C
E)24.5 J/g·°C
Q2) Aluminum oxide can be reduced to aluminum metal using carbon, the other reaction product being carbon monoxide. Determine the enthalpy change when 12.5 g of aluminum is produced by this method.[\(\Delta\)H°<sub>f</sub>(carbon monoxide)= -110.5 kJ/mol; \(\Delta\)H°<sub>f</sub>(aluminum oxide)= -1669.8 kJ/mol]
A)725 kJ
B)697 kJ
C)310 kJ
D)361 kJ
E)1504 kJ
Q3) A 26.2 g piece of copper metal is heated from 21.5°C to 201.6°C. Calculate the amount of heat absorbed by the metal. The specific heat of Cu is 0.385 J/g·°C.
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Chapter 7: Quantum Theory and the Electronic Structure of Atoms
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Sample Questions
Q1) A ground-state atom of iron has ___ unpaired electrons and is _____.
A)0, diamagnetic
B)6, diamagnetic
C)3, paramagnetic
D)5, paramagnetic
E)4, paramagnetic
Q2) When the electron in a hydrogen atom falls from its first excited energy level to the ground state energy level, a photon with wavelength \(\lambda\) is emitted. A proton having this same wavelength would have a velocity of
A)3.87 m/s
B)5990 m/s
C)1.21 * 10<sup>-7</sup> m/s
D)3.26 m/s
E)5.99 m/s
Q3) A ground-state atom of arsenic has A)no unpaired electrons.
B)one unpaired electron.
C)two unpaired electrons. D)three unpaired electrons.
E)four unpaired electrons.
Page 9
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Chapter 8: Periodic Relationships Among the Elements
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Sample Questions
Q1) Which ion is isoelectronic with Ar?
A)Fe<sup>2+</sup>
B)F<sup>-</sup>
C)Br<sup>-</sup>
D)Ga<sup>3+</sup>
E)Ca<sup>2+</sup>
Q2) The law of octaves states that every eighth element has similar properties when arranged in order of increasing A)nuclear binding energy.
B)number of electrons.
C)number of neutrons.
D)atomic mass.
E)atomic number.
Q3) The electron affinity of fluorine is essentially equal to
A)the negative of the ionization energy F.
B)the ionization energy F<sup>-</sup>.
C)the negative of the ionization energy F<sup>-</sup>.
D)the ionization energy Ne.
E)the negative of the ionization energy Ne.
Q4) Write the ground-state electron configuration for Ca<sup>2+</sup>.
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Chapter 9: Chemical Bonding I: Basic Concepts
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Sample Questions
Q1) Which one of the following molecules has an atom with an incomplete octet?
A)NF<sub>3</sub>
B)H<sub>2</sub>O
C)AsCl<sub>3</sub>
D)GeH<sub>4</sub>
E)BF<sub>3</sub>
Q2) Which one of the following is most likely to be an ionic compound?
A)NCl<sub>3</sub>
B)BaCl<sub>2</sub>
C)CO
D)SO<sub>2</sub>
E)SF<sub>4</sub>
Q3) The formal charge on the bromine atom in BrO<sub>3</sub><sup>- </sup> drawn with three single bonds is A)-2 B)-1 C)0
Q4) Write the Lewis structure of boron trifluoride.
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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization
of Atomic Orbitals
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Sample Questions
Q1) Use VSEPR theory to predict the molecular geometry of SF<sub>4</sub> (sulfur tetrafluoride).
Q2) Which of the following molecules have the same geometries?
A)SF<sub>4</sub> and CH<sub>4</sub>
B)CO<sub>2</sub> and H<sub>2</sub>O
C)CO<sub>2</sub> and BeH<sub>2</sub>
D)N<sub>2</sub>O and NO<sub>2</sub>
Q3) Which should have the longer bond, B<sub>2</sub> or B<sub>2</sub><sup>-</sup>?
Q4) How does the geometrical structure of PF<sub>5</sub> differ from that of IF<sub>5</sub>?
Q5) According to VSEPR theory, which one of the following molecules has tetrahedral geometry?
A)NH<sub>3</sub>
B)CCl<sub>4</sub>
C)CO<sub>2</sub>
D)SF<sub>4</sub>
E)PCl<sub>5</sub>
Q7) Indicate the number of \(\pi\)-bonds in N<sub>2</sub>H<sub>2</sub>. Page 12
Q6) Draw a Lewis structure for PF<sub>5</sub> that shows the correct atom arrangement predicted by VSEPR theory.
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Chapter 11: Intermolecular Forces and Liquids and Solids
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Sample Questions
Q1) The shape of the water-to-glass meniscus results from the strong adhesive forces between glass and water.
A)True
B)False
Q2) Which one of the following substances should exhibit hydrogen bonding in the liquid state?
A)PH<sub>3</sub>
B)He
C)H<sub>2</sub>S
D)CH<sub>4</sub>
E)CH<sub>3</sub>OH
Q3) Which one of the following substances is expected to have the lowest melting point?
A)BrI
B)CsI
C)LiI
D)NaI
E)RbI
Q4) Indicate all the types of intermolecular forces of attraction in HCl(g).
Q5) Indicate all the types of intermolecular forces of attraction in HF(l).
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Chapter 12: Physical Properties of Solutions
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Sample Questions
Q1) Calculate the approximate freezing point of a solution made from 21.0 g NaCl and 1.00 * 10<sup>2</sup> g of H<sub>2</sub>O. [K<sub>f</sub> of water is 1.86°C/m.]
A)3.59°C
B)6.68°C
C)-13.4°C
D)-6.68°C
E)-3.59°C
Q2) How many liters of ethylene glycol antifreeze (C<sub>2</sub>H<sub>6</sub>O<sub>2</sub>)would you add to your car radiator containing 15.0 L of water if you needed to protect your engine to -17.8°C? (The density of ethylene glycol is 1.1 g/mL. For water, K<sub>f</sub> = 1.86°C/m.)
Q3) Which of the following aqueous solutions has the lowest freezing point?
A)0.18 m KCl
B)0.15 m Na<sub>2</sub>SO<sub>4</sub>
C)0.12 m Ca(NO<sub>3</sub>)<sub>2</sub>
D)pure water
E)0.20 m C<sub>2</sub>H<sub>6</sub>O<sub>2</sub> (ethylene glycol)
Q4) What is the molality of a 0.142 M Na<sub>3</sub>PO<sub>4</sub>(aq)solution that has a density of 1.015 g/mL?
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Chapter 13: Chemical Kinetics
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Sample Questions
Q1) The first-order decomposition of phosphene to phosphorus and hydrogen 4PH<sub>3</sub>(g)F1F1F1S1 F1F1F10 P<sub>4</sub>(g)+ 6H<sub>2</sub>(g)
Has a half-life of 35.0 s at 680°C. Starting with 520 mmHg of pure phosphene in an 8.00-L flask at 680°C, how long will it take for the total pressure in the flask to rise to 1.000 atm?
A)628 s
B)33.4 s
C)51.2 s
D)111 s
E)48.3 s
Q2) For the first-order reaction, A F1F1F1S1 F1F1F10 products, if half of the initial concentration of A reacts in 20 min, then the remaining half will completely react in the next 20 min.
A)True
B)False
Q3) An experimental drug, D, is known to decompose in the blood stream.Tripling the concentration of the drug increases the decomposition rate by a factor of nine.Write the rate law for decomposition of D and give the units of the rate constant.
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Chapter 14: Chemical Equilibrium
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Sample Questions
Q1) Describe why addition of a catalyst does not affect the equilibrium constant for a reaction.
Q2) What conditions are used in the Haber process to enhance the yield of ammonia? Explain why each condition affects the yield in terms of the Le Châtelier principle.
Q3) Which of these statements is true about chemical equilibria in general?
A)At equilibrium the total concentration of products equals the total concentration of reactants, that is, [products] = [reactants].
B)Equilibrium is the result of the cessation of all chemical change.
C)There is only one set of equilibrium concentrations that equals the K<sub>c</sub> value.
D)At equilibrium, the rate constant of the forward reaction is equal to the rate constant for the reverse reaction.
E)At equilibrium, the rate of the forward reaction is equal to as the rate of the reverse reaction.
Q4) The dissociation of solid silver chloride in water to produce silver ions and chloride ions has an equilibrium constant of 1.8 * 10<sup>-18</sup>.Based on the magnitude of the equilibrium constant, is silver chloride very soluble in water? Why?
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Page 17
Chapter 15: Acids and Bases
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Sample Questions
Q1) Al(OH)<sub>3</sub> is an amphoteric hydroxide. Write a balanced ionic equation to show its reaction with KOH.
Q2) Calculate the concentration of malonate ion (C<sub>3</sub>H<sub>2</sub>O<sub>4</sub><sup>2-</sup>)in a 0.200 M solution of malonic acid (C<sub>3</sub>H<sub>4</sub>O<sub>4</sub>).[For malonic acid, K<sub>a1</sub> = 1.4 * 10<sup>-3</sup>, K<sub>a2</sub> = 2.0 * 10<sup>-6</sup>.]
A)2.8 * 10<sup>-4</sup> M
B)0.016 M
C)1.8 * 10<sup>-4</sup> M
D)1.4 * 10<sup>-3</sup> M
E)2.0 * 10<sup>-6</sup> M
Q3) Arrange the acids H<sub>2</sub>Se, H<sub>2</sub>Te, and H<sub>2</sub>S in order of increasing acid strength.
A)H<sub>2</sub>S < H<sub>2</sub>Se < H<sub>2</sub>Te
B)H<sub>2</sub>S < H<sub>2</sub>Te < H<sub>2</sub>Se
C)H<sub>2</sub>Te < H<sub>2</sub>S < H<sub>2</sub>Se
D)H<sub>2</sub>Se<sub> </sub>< H<sub>2</sub>S < H<sub>2</sub>Te
E)H<sub>2</sub>Se<sub> </sub>< H<sub>2</sub>Te < H<sub>2</sub>S
Q4) Write the chemical formula for sulfuric acid.
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Chapter 16: Acid-Base Equilibria and Solubility Equilibria
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Sample Questions
Q1) Describe how to prepare 500.mL of a cyanic acid (HCNO)/sodium cyanate (NaCNO)buffer having a pH of 4.80. [K<sub>a</sub>(HCNO)= 2.0 * 10<sup>-4</sup>]
Q2) What is the effective pH range for a sodium acetate/acetic acid buffer? (For CH<sub>3</sub>COOH, K<sub>a</sub> = 1.8 * 10<sup>-5</sup>)
Q3) Assuming equal concentrations of conjugate base and acid, which one of the following mixtures is suitable for making a buffer solution with an optimum pH of 4.6-4.8?
A)CH<sub>3</sub>COO<sub>2</sub>Na / CH<sub>3</sub>COOH (K<sub>a</sub> = 1.8 * 10<sup>-5</sup>)
B)NH<sub>3</sub> / NH<sub>4</sub>Cl (K<sub>a</sub> = 5.6 * 10<sup>-10</sup>)
C)NaOCl / HOCl (K<sub>a</sub> = 3.2 * 10<sup>-8</sup>)
D)NaNO<sub>2</sub> / HNO<sub>2</sub> (K<sub>a</sub> = 4.5 * 10<sup>-4</sup>)
E)NaCl / HCl
Q4) The concentration of Mg<sup>2+</sup> in seawater is 5.0 * 10<sup>-2 </sup>M. What hydroxide concentration is needed to remove 90% of the Mg<sup>2+</sup> by precipitation? (For Mg(OH)<sub>2</sub>, K<sub>sp</sub> = 1.2 * 10<sup>-11</sup>.)
Q5) Describe how to make a sodium formate (HCOONa)/formic acid (HCOOH)buffer that has a pH of 4.77.
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Page 19
Chapter 17: Chemistry in the Atmosphere
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Sample Questions
Q1) One of the steps in the fixation of nitrogen by lightening is given below. 2NO<sub>2</sub>(g)+ H<sub>2</sub>O(l)\(\to\) HNO<sub>2</sub>(aq)+ HNO<sub>3</sub>(aq)
In this reaction, is the element nitrogen oxidized or reduced?
A)Nitrogen is oxidized in this reaction.
B)Nitrogen is reduced in this reaction.
C)Nitrogen is oxidized and reduced in this reaction.
D)Nitrogen is neither oxidized nor reduced in this reaction.
Q2) The most prevalent noble gas in our atmosphere is
A)helium.
B)neon.
C)argon.
D)krypton.
E)xenon.
Q3) The compound CFCl<sub>3</sub> is used as a/an
A)enzyme
B)anesthetic
C)gaseous fuel
D)coolant
E)CFC replacement
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Chapter 18: Entropy, Free Energy, and Equilibrium
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Sample Questions
Q1) K<sub>w</sub> for the auto-ionization of water, H<sub>2</sub>O(l)\(\to\) H<sup>+</sup>(aq)+ OH<sup>-</sup> (aq), is 1.0 \(\times\) 10<sup>-14</sup>. What are the signs (+/-)of \(\Delta\)S° and \(\Delta\)H° for the reaction at 25°C?
A)(\(\Delta\)S° = (+)and \(\Delta\)H° = (+))
B)(\(\Delta\)S° = (+)and \(\Delta\)H° = (-))
C)(\(\Delta\)S° = (-)and \(\Delta\)H° = (+))
D)(\(\Delta\)S° = (-)and \(\Delta\)H° = (-))
Q2) Predict the sign of \(\Delta\)S for the process Ni(s, 50°C)\(\to\) Ni(s, 25°C).
Q3) Arrange these reactions according to increasing \(\Delta\)S.
1)H<sub>2</sub>O(g)\(\to\) H<sub>2</sub>O(l)
2)2NO(g)\(\to\) N<sub>2</sub>(g)+ O<sub>2</sub>(g)
3)MgCO<sub>3</sub>(s)\(\to\)MgO(s)+ CO<sub>2</sub>(g)
A)1 < 2 < 3
B)2 < 3 < 1
C)3 < 2 < 1
D)2 < 1 < 3
E)1 < 3 < 2
Q4) Rubidium has a heat of vaporization of 69.0 kJ/mol at its boiling point (686°C).Calculate \(\Delta\)S for this process, Rb(l)\(\to\) Rb(g), at 1 atm and 686°C.
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Chapter 19: Electrochemistry
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Sample Questions
Q1) Suppose the reaction Pb(s)+ 2H<sup>+</sup>(aq)\(\to\)Pb<sup>2+</sup>(aq)+ H<sub>2</sub>(g)is carried out at pH = 4.00 and at a hydrogen gas pressure of 1.00 atm. The concentration of lead(II)ions that causes this reaction to be at equilibrium is
A)2.5 M
B)1.6 \(\times\) 10<sup>-2</sup> M
C)2.5 \(\times\) 10<sup>-4</sup> M
D)1.6 \(\times\) 10<sup>-6</sup> M
E)0.40 M
Q2) For the reaction Ni<sup>2+</sup>(aq)+ 2Fe<sup>2+</sup>(aq)\(\to\) Ni(s)+ 2Fe<sup>3+</sup>(aq), the standard cell potential E°<sub>cell</sub> is
A)+2.81 V.
B)+1.02 V.
C)+0.52 V.
D)-1.02 V.
E)-2.81 V.
Q3) Consider the reaction Fe + Sn<sup>2+</sup>(1 \(\times\) 10<sup>-3</sup> M)\(\to\) Fe<sup>2+</sup>(1.0 M)+ Sn. Calculate the voltage theoretically generated by the cell.
Q4) Will H<sub>2</sub>(g)form when Sn is placed in 1.0 M HCl?
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Chapter 20: Metallurgy and the Chemistry of Metals
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Sample Questions
Q1) Write the chemical formula of fluorite.
Q2) Which of these reactions represents the removal of silica from iron ore in a blast furnace?
A)SiO<sub>2</sub>(s)\(\to\)SiO<sub>2</sub>(g)
B)SiO<sub>2</sub>(s)+ CaO(s)\(\to\) CaSiO<sub>3</sub>(l)
C)SiO<sub>2</sub>(s)+ 4HF(g)\(\to\) SiF<sub>4</sub>(g)+ 2H<sub>2</sub>O(g)
D)SiO<sub>2</sub>(s)+ C(s)\(\to\)Si(l)+ CO<sub>2</sub>(g)
E)SiO<sub>2</sub>(s)+ CO(g)\(\to\)SiCO<sub>3</sub>(l)
Q3) In the Hall process, ____________ is reduced ____________.
A)nickel; electrolytically
B)aluminum; electrolytically
C)nickel; by reaction with metallic sodium
D)aluminum; by reaction with metallic sodium
E)copper; electrolytically
Q4) Which one of these metals would normally be obtained by electrolytic reduction?
A)aluminum
B)chromium
C)copper
D)iron
E)zinc
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Chapter 21: Nonmetallic Elements and Their Compounds
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Sample Questions
Q1) When liquid phosphorus trichloride reacts with water, the products are
A)PCl<sub>5</sub> and H<sub>3</sub>PO<sub>4</sub>.
B)H<sub>3</sub>PO<sub>3</sub> and Cl<sub>2</sub>.
C)H<sub>3</sub>PO<sub>4</sub> and HCl.
D)H<sub>3</sub>PO<sub>4</sub> and Cl<sub>2</sub>.
E)H<sub>3</sub>PO<sub>3</sub> and HCl.
Q2) Which one of these species exists as a diatomic element at room temperature?
A)Na
B)N
C)B
D)Ar
E)P
Q3) The main commercial source of bromine is A)silver bromide ore. B)seawater.
C)the Frasch process.
D)phosphate rock.
E)the Ostwald process.
Q4) Write an equation for a laboratory preparation of nitrogen gas.
Q5) Write the formula for the binary hydride you would expect calcium to form.
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Chapter 22: Transition Metal Chemistry and Coordination Compounds
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Sample Questions
Q1) In K<sub>4</sub>[Fe(CN)<sub>6</sub>], how many 3d electrons does the iron atom have?
A)3
B)4
C)5
D)6
E)7
Q2) What is the coordination number of cobalt in [Co(NH<sub>3</sub>)<sub>6</sub>]Cl<sub>3</sub>?
Q3) Write the chemical formula of triammineaquodichlorochromium(III)chloride.
Q4) How many geometric isomers are possible for the complex [CrF<sub>3</sub>Br<sub>3</sub>]<sup>3-</sup>? Draw these isomers.
Q5) Predict the number of unpaired electrons in the [Fe(CN)<sub>6</sub>]<sup>3-</sup> ion.
Q6) The total number of electrons in the 3d orbitals of Cr<sup>3+</sup> is A)1 B)2 C)3 D)4 E)5
Page 25
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Chapter 23: Nuclear Chemistry
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Sample Questions
Q1) What element is the stable end-product of the uranium radioactive decay series?
A)Th
B)Pu
C)Ra
D)Au
E)Pb
Q2) Protactinium-234 has a half-life of 1 minute.How much of a 400.g sample protactinium would remain after 2 minutes?
Q3) The blood volume of a patient can be measured by using chromium-51, a positron emitter, administered as a solution of sodium chromate.Write the nuclear equation for the decay of chromium-51.
Q4) Gamma-rays cause radiation damage when they interact with matter by producing A)ions and free radicals.
B)isotopes.
C)daughter products.
D)oxidation.
E)reduction.
Q5) Radioactive nitrogen-13 has a half-life of 10 minutes.After 40 minutes, how much of this isotope would remain in a sample that originally contained 96 mg?
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Chapter 24: Organic Chemistry
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Sample Questions
Q1) "Wood alcohol" is the common name for A)methanol.
B)ethanol.
C)propyl alcohol.
D)ethylene.
E)acetylene.
Q2) Which of the following compounds are isomers of each other?
I.pentane
II.2-methylbutane
III.2,3-dimethylbutane
IV.2,2-dimethylpropane
V.1-hexene
Q3) The oxidation product of 1-propanol when using Cr<sub>2</sub>O<sub>7</sub><sup>2-</sup> as the oxidizing agent is acetone.
A)True
B)False
Q4) The reaction of hydrogen chloride gas with propene will yield 1-chloropropane as the main product.
A)True
B)False
Page 28
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Chapter 25: Synthetic and Natural Organic Polymers
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Q1) Amides are synthesized from two classes of organic compounds. Those two types of compounds are
A)carboxylic acids and alkenes.
B)amines and alcohols.
C)alcohols and carboxylic acids.
D)amines and carboxylic acids.
E)alkenes and amines.
Q2) Which one of these materials is a copolymer?
A)Styrene-butadiene
B)polyvinyl chloride
C)polypropylene
D)poly-cis-isoprene
E)polyethylene
Q3) Both DNA and RNA have double-helical structures.
A)True
B)False
Q4) The monomer used to prepare polyvinyl chloride (PVC)is CHCl=CHCl.
A)True
B)False
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