Chemistry for Physical Sciences
Final Exam
Course Introduction
This course provides a comprehensive introduction to the fundamental principles of chemistry with a focus on their applications within the physical sciences. Students will explore atomic and molecular structure, chemical bonding, stoichiometry, states of matter, thermodynamics, chemical kinetics, and equilibrium. Emphasis is placed on developing problem-solving skills and understanding the quantitative and qualitative aspects of chemical phenomena as they relate to physical processes. Laboratory components and practical exercises reinforce theoretical knowledge, preparing students for more advanced studies in physics, engineering, and related scientific disciplines.
Recommended Textbook Principles of Chemistry The Molecular Science 1st Edition by John W. Moore
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Page 2
Chapter 1: The Nature of Chemistry
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Sample Questions
Q1) Which characteristic below best fits the description of a liquid?
A) rapid molecular motion
B) highly disordered molecules
C) large distances between the molecules
D) molecules that are close together but are moving past one another
E) highly ordered molecules
Answer: D
Q2) Which of the following would not convert a heterogeneous mixture into a more homogeneous material?
A) filtering a salt solution
B) magnetically removing iron from an iron-sulfur mixture
C) filtering sand away from water
D) picking rice grains from a mixture of rice and peas
E) doing chromatography on an ink sample
Answer: A
Q3) The ability to conduct electricity is a(n) _____________ property.
Answer: physical
Q4) Burning of hydrogen fuel is a(n) _____________ change.
Answer: chemical
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Chapter 2: Atoms and Elements
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Sample Questions
Q1) The Millikan oil drop experiment was used to determine
A) whether the atom was radioactive.
B) the electron distribution in the atom.
C) the atomic number of an atom.
D) the charge of an electron.
E) the nuclear character of the atom.
Answer: D
Q2) How many hydrogen atoms are in a 27.0 g sample of water (H<sub>2</sub>O)?
A) 1.50
B) 2.99
C) 4.98 × 10<sup>-24</sup>
D) 1.81 × 10<sup>24</sup>
E) 2.93 × 10<sup>26</sup>
Answer: D
Q3) Write the atomic symbol for an atom that contains 24 protons, 24 electrons and 26 neutrons.
Answer: 11ea782d_fd70_540d_bbd5_b1588926a94a_TB5061_11
Q4) <sup>14</sup>O, <sup>16</sup>O, <sup>18</sup>O are called _____________.
Answer: isotopes
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Chapter 3: Chemical Compounds
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Sample Questions
Q1) Caffeine has the empirical formula C<sub>4</sub>H<sub>5</sub>N<sub>2</sub>O and an approximate molecular weight of 194.2 g/mol. What is its molecular formula?
A) C<sub>4</sub>H<sub>5</sub>N<sub>2</sub>O
B) C<sub>2</sub>H<sub>2</sub>NO
C) C<sub>8</sub>H<sub>10</sub>N<sub>4</sub>O<sub>2</sub>
D) C<sub>4</sub>H<sub>4</sub>N<sub>4</sub>O<sub>4</sub>
E) C<sub>2</sub>H<sub>2</sub>N<sub>2</sub>O<sub>2</sub>
Answer: C
Q2) Determine the number of molecules in 325 mg of aspirin, C<sub>9</sub>H<sub>8</sub>O<sub>4</sub>.
A) 1.81 × 10<sup>-3</sup>
B) 1.09 × 10<sup>21</sup>
C) 3.34 × 10<sup>23</sup>
D) 1.09 × 10<sup>24</sup>
E) 3.34 × 10<sup>26</sup>
Answer: B
Q3) Give the formula for an alkane containing five carbon atoms. Answer: C<sub>5</sub>H<sub>12</sub>
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Chapter 4: Quantities of Reactants and Products
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Sample Questions
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) A reaction has a theoretical yield of 48.23 g. It produces 28.81 g of product. Calculate the percent yield.
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>
Q4) A combination reaction is considered the opposite of a(n) _____________ reaction.
Q5) Stoichiometric coefficients found in a balanced equation can be used to derive _____________ ratios.
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Chapter 5: Chemical Reactions
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Sample Questions
Q1) A 25.00 mL sample of HCl solution is neutralized by exactly 31.22 mL of 0.08152 M Ca(OH)<sub>2</sub>. What is the molarity of the HCl solution?
A) 0.08152 M
B) 0.1018 M
C) 0.2036 M
D) 0.1021 M
E) 0.09453 M
Q2) Which compound will not dissolve in water in large amounts?
A) KNO<sub>3</sub>
B) NH<sub>4</sub>Cl
C) Ca(OH)<sub>2</sub>
D) AgCl
E) Ag<sub>2</sub>SO<sub>4</sub>
Q3) 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
Q4) When an element is _____________, it loses electrons.
Page 7
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Chapter 6: Energy and Chemical Reactions
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Sample Questions
Q1) In an endothermic reaction, heat is transferred from the _____________ to the
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) According to the First Law of Thermodynamics, the total _____________ of the universe is constant.
Q4) Enthalpy change is equal to heat transfer at constant _______________.
Q5) What is the molar heat capacity of table salt, NaCl (specific heat = 0.88 J g<sup>-1</sup> °C<sup>-1</sup>)?
A) 5.30 × 10<sup>22</sup> J mol<sup>-1</sup> °C<sup>-1</sup>
B) 24.6 J mol<sup>-1</sup> °C<sup>-1</sup>
C) 51.4 J mol<sup>-1</sup> °C<sup>-1</sup>
D) 117 J mol<sup>-1</sup> °C<sup>-1</sup>
E) 245 J mol<sup>-1</sup> °C<sup>-1</sup>
Q6) Energy of motion is called _____________.
Page 8
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Chapter 7: Electron Configurations and the Periodic Table
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Sample Questions
Q1) Each electron in an atom or ion must have a unique set of four _____________.
Q2) The first ionization energy is the amount of energy required to remove one ____________ from a neutral atom.
Q3) As one moves closer to the nucleus, what happens to the value of the principal quantum number?
Q4) The d orbitals occur in groups of ____ and hold up to ____ electrons.
A) 3; 6
B) 4; 8
C) 5; 10
D) 6; 12
E) 7; 14
Q5) Light has a frequency of 7.21 × 10<sup>13</sup> hertz. What is its wavelength? The speed of light is 3.00 × 10<sup>8</sup> m/s.
A) 2.16 × 10<sup>9</sup> nm
B) 2.40 × 10<sup>5</sup> nm
C) 4.16 × 10<sup>3</sup> nm
D) 2.40 × 10<sup>-4</sup> nm
E) 4.16 × 10<sup>-6</sup> nm
Q6) Which halogen atom has the smallest atomic radius?
Page 9
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Chapter 8: Covalent Bonding
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Sample Questions
Q1) Which bond is longest?
A) C-O
B) C-I
C) C-Br
D) C-C
E) C-N
Q2) Which molecule contains a triple bond?
A) C<sub>2</sub>H<sub>4</sub>
B) CCl<sub>4</sub>
C) H<sub>2</sub>O
D) N<sub>2</sub>
E) O<sub>2</sub>
Q3) Which bond is shortest?
A) carbon-oxygen single bond
B) carbon-hydrogen single bond
C) hydrogen-hydrogen single bond
D) carbon-carbon double bond
E) carbon-oxygen triple bond
Q4) The ____________ of an element is its ability to pull electrons towards itself when participating in a covalent bond.
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Chapter 9: Molecular Structure
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Q1) The carbon-carbon p bond in ethylene, CH<sub>2</sub>CH<sub>2</sub>, results from the overlap of which of the following?
A) s atomic orbitals
B) p atomic orbitals
C) sp hybrid orbitals
D) sp<sup>2</sup> hybrid orbitals
E) sp<sup>3</sup> hybrid orbitals
Q2) Use the VSEPR model to predict the electron-pair geometry and the molecular geometry of CO<sub>2</sub>.
A) tetrahedral, linear
B) tetrahedral, tetrahedral
C) linear, bent
D) bent, linear
E) linear, linear
Q3) Which molecule is polar?
A) BF<sub>3</sub>
B) H<sub>2</sub>Se
C) N<sub>2</sub>
D) GeF<sub>4</sub>
E) CO<sub>2</sub>
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Chapter 10: Gases and the Atmosphere
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Q1) Aluminum and oxygen react to form aluminum oxide. How many L of oxygen at 296 K and 1.4 atm of pressure are necessary to form 3.25 g of aluminum oxide? R = 0.0821 L atm mol<sup>-1</sup> K<sup>-1</sup>.
A) 6.5 × 10<sup>-2</sup> L
B) 0.55 L
C) 0.83 L
D) 1.3 L
E) 56 L
Q2) A 5.00 L flask contains 3.50 g of sulfur trioxide, 2.45 g of carbon monoxide, and 3.99 g of argon all at 35°C. What is the pressure (in atm) in the flask?
A) 0.221 atm
B) 0.234 atm
C) 50.5 atm
D) 5.70 atm
E) 1.17 atm
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Chapter 11: Liquids, Solids, and Materials
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Q1) The volatility of a liquid is increased by
A) decreasing the temperature and increasing the atmospheric pressure.
B) decreasing the temperature and decreasing the atmospheric pressure.
C) increasing the temperature and increasing the atmospheric pressure.
D) increasing the temperature and decreasing the atmospheric pressure.
E) decreasing the temperature and maintaining the atmospheric pressure.
Q2) The density of water has its maximum value at 4<sup>\(\circ\)</sup>C because
A) the thermal motion of the molecules overcomes the hydrogen bonds and allows the molecules to approach each other more closely.
B) the thermal motion of the molecules is much greater than the energy of the hydrogen bonds and allows the molecules a great deal of freedom of motion.
C) covalent bonds begin breaking at this temperature.
D) the molecules are trapped in place by very strong hydrogen bonds.
E) the maximum number of hydrogen bonds exists between the molecules at this temperature.
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13
Chapter 12: Chemical Kinetics: Rates of Reactions
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Q1) The rate constant for the first order decomposition of SO<sub>2</sub>Cl<sub>2</sub> is 1.37 × 10<sup>-3</sup> min<sup>-1</sup> at a certain temperature. If the concentration is initially 0.500 M, predict it after 300 min.
A) 6.85 × 10<sup>-4</sup> M
B) 0.331 M
C) 0.500 M
D) 0.754 M
E) 1.33 M
Q2) Which statement, comparing catalyzed and uncatalyzed versions of a reaction, is incorrect?
A) The initial and final energy states are the same.
B) The reaction mechanisms are different.
C) The reaction progress diagrams are different.
D) The balanced equation for the overall reaction is the same.
E) The activation energy is lower in the uncatalyzed reaction.
Q3) When studying kinetics, it is more accurate to measure instantaneous reaction rates at particular times rather than the average rates of reaction over time intervals. Explain why.
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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) If a reaction is product-favored on the basis of enthalpy and reactant-favored on the basis of entropy, then it most likely is ____ and ____ net moles of gas.
A) exothermic; produces
B) exothermic; consumes
C) exothermic; neither produces nor consumes
D) endothermic; produces
E) endothermic; consumes
Q3) A chemical equilibrium may involve
A) reactants and products exhibiting different states of matter
B) molecular rearrangements
C) a chemical change of one compound into another
D) reaction of elements to produce one or more compounds
E) any of these
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Chapter 14: The Chemistry of Solutes and Solutions
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Q1) In order to calculate the molality of a solution from its molarity, what other information would be needed?
A) molar mass of solute and molar mass of solvent
B) molar mass of solute and density of solvent
C) molar mass of solute and density of solution
D) molar mass of solvent and density of solution
E) molar mass of solvent and density of solute
Q2) If two liquids are miscible, they ____ because ____.
A) mix together; their intermolecular interactions are similar
B) do not mix together; their intermolecular interactions are similar
C) do not mix together; their intermolecular interactions are dissimilar
D) mix together; their intermolecular interactions are dissimilar
E) do not mix together; their densities are dissimilar
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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Sample Questions
Q1) Which reaction illustrates water acting as a base?
A) Cu(H<sub>2</sub>O)<sub>4</sub><sup>2+</sup> + 4 NH<sub>3</sub> Cu(NH<sub>3</sub>)<sub>4</sub><sup>2+</sup> + 4 H<sub>2</sub>O
B) H<sub>2</sub>CO<sub>3</sub> H<sub>2</sub>O + CO<sub>2</sub>
C) NH<sub>3</sub> + H<sub>2</sub>O NH<sub>4</sub><sup>+</sup> + OH<sup>-</sup>
D) H<sub>2</sub>PO<sub>4</sub><sup>-</sup> + H<sub>2</sub>O OH<sup>-</sup> + H<sub>3</sub>PO<sub>4</sub>
E) HSO<sub>4</sub><sup>-</sup> + H<sub>2</sub>O H<sub>3</sub>O<sup>+</sup> + SO<sub>4</sub><sup>2-</sup>
Q2) Carboxylic acids have the formula
A) R-CH<sub>3</sub>
B) R-CHO
C) R-COOH
D) R-NH<sub>2</sub>
E) R-OH
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) What is the minimum volume of 0.55 M NaOH required to completely convert 1.00 L of 0.45 M acetic acid (pK<sub>a</sub> = 4.74) to its conjugate base?
A) 0.82 L
B) 1.0 L
C) 1.2 L
D) 1.8 L
E) 2.2 L
Q2) Silver phosphate is less soluble in Na<sub>3</sub>PO<sub>4</sub> than in water because
A) a common ion displaces the solubility equilibrium towards the undissolved solute.
B) some insoluble compounds are amphoteric.
C) the solubility of most salts increases as temperature increases.
D) the formation of complex ions displaces the solubility equilibrium to the right.
E) the solubility of many salts is affected by the pH of the solution.
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18
Chapter 17: Thermodynamics: Directionality of Chemical Reactions
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Q1) The enthalpy change for a reaction can be calculated using the equation DH°<sub>rxn</sub> It would be possible to use DS<sup>\(\circ\)</sup><sub>f</sub> values to calculate DS<sup>\(\circ\)</sup><sub>rxn</sub> in a similar manner; however, this is not done in practice.
a. Describe the procedure which is used instead.
b. Explain why this mathematical process is possible for entropy calculations but not for enthalpy calculations.
Q2) What is the value of the equilibrium constant at 25<sup>\(\circ\)</sup>C for a reaction, if the value of DG<sup>\(\circ\)</sup><sub>rxn</sub> is -47.8 kJ at 25<sup>\(\circ\)</sup>C?
A) 1.70
B) 6.88
C) 2.30 × 10<sup>2</sup>
D) 2.74 × 10<sup>5</sup>
E) 2.37 × 10<sup>8</sup>
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Chapter 18: Electrochemistry and Its Applications
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Q1) A mass of 0.839 g of a divalent metal is plated out of a solution of the divalent metal ion. This takes 67.2 min at a current of 0.63 A. What is the metal? (Hint: find its atomic mass.)
A) Cd
B) Cu
C) Hg
D) Fe
E) Mg
Q2) In the salt bridge in an electrochemical cell,
A) negatively charged ions migrate towards the anode compartment
B) positively charged ions migrate towards the anode compartment
C) negatively charged ions migrate towards the cathode compartment
D) both b and c
E) neither a nor c
Q3) The unit used to measure electrical current is the
A) ampere
B) coulomb
C) faraday
D) joule
E) volt
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Chapter 19: Nuclear Chemistry
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Q1) If polonium-210 emits an alpha particle, the other product will be
A) lead-206.
B) mercury-204.
C) mercury-206.
D) polonium-208.
E) radon-206.
Q2) 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
Q3) A radioactive isotope that decomposes by positron emission loses
A) an electron
B) a proton
C) a helium nucleus
D) a particle similar to an electron, but positively charged
E) a positively-charged neutron
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