Basic Inorganic Chemistry
Final Exam
Course Introduction
Basic Inorganic Chemistry introduces students to the fundamental principles and concepts of inorganic chemistry, focusing on the properties, structures, reactions, and uses of elements and inorganic compounds. The course covers topics such as atomic structure, periodic table trends, chemical bonding, molecular geometry, acid-base theories, coordination chemistry, and the chemistry of main group and transition elements. Emphasis is placed on understanding the reactivity and applications of inorganic substances, as well as developing problem-solving skills through laboratory experiments and theoretical exercises.
Recommended Textbook
Principles of General Chemistry 2nd Edition by Martin Silberberg
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Page 2
Chapter 1: Keys to the Study of Chemistry
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Sample Questions
Q1) Which of the following is a chemical change?
A)boiling of water
B)melting wax
C)broiling a steak on a grill
D)condensing water vapor into rainfall
E)carving a piece of wood
Answer: C
Q2) Which of the following is not an S.I.base unit?
A)meter
B)ampere
C)second
D)gram
E)kelvin
Answer: D
Q3) An important aim in much chemical work is to use macroscopic measurements in order to gain an understanding of the microscopic world.
A)True
B)False
Answer: True
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Chapter 2: The Components of Matter
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Q1) What is the name of P<sub>4</sub>Se<sub>3</sub>?
A)phosphorus selenide
B)phosphorus triselenide
C)tetraphosphorus selenide
D)phosphoric selenide
E)tetraphosphorus triselenide
Answer: E
Q2) The molecular formula of a compound provides more information than its structural formula.
A)True
B)False
Answer: False
Q3) The substance,CaSe,is used in materials which are electron emitters.What is its name?
A)calcium monoselenide
B)calcium(II)selenide
C)calcium selenide
D)calcium(I)selenide
E)calcium(II)selenium
Answer: C
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Chapter 3: Stoichiometry of Formulas and Equations
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Sample Questions
Q1) Balance the following equation for partial oxidation of ammonia,an important reaction in the production of nitric acid: NH<sub>3</sub>(g)+ O<sub>2</sub>(g) \(\to\) NO(g)+ H<sub>2</sub>O(l)
Answer: 4NH<sub>3</sub>(g)+ 5O<sub>2</sub>(g) \(\to\) 4NO(g)+ 6H<sub>2</sub>O(l)
Q2) Sodium bromate is used in a mixture which dissolves gold from its ores.Calculate the mass in grams of 4.68 mol of sodium bromate.
A)706 g
B)482 g
C)383 g
D)32.2 g
E)0.0310 g
Answer: A
Q3) When a solution is diluted with water,the ratio of the initial to final volumes of solution is equal to the ratio of final to initial molarities.
A)True
B)False
Answer: True
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Chapter 4: Three Major Classes of Chemical Reactions
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Sample Questions
Q1) Which one of the following substances is a strong acid?
A)HNO<sub>3</sub>
B)H<sub>2</sub>CO<sub>3</sub>
C)NH<sub>3</sub>
D)CH<sub>3</sub>COOH
E)H<sub>3</sub>PO<sub>4</sub>
Q2) A 0.00100 mol sample of Ca(OH)<sub>2</sub> requires 25.00 mL of aqueous HCl for neutralization according to the reaction below.What is the concentration of the HCl? Equation: Ca(OH)<sub>2</sub>(s)+ 2HCl(aq) \(\to\) CaCl<sub>2</sub>(aq)+
H<sub>2</sub>O(l)
A)0.0200 M
B)0.0400 M
C)0.0800 M
D)4.00 \(\times\) 10<sup>-5</sup> M
E)None of these choices is correct.
Q3) An aqueous solution of lead nitrate,Pb(NO<sub>3</sub>)<sub>2</sub>,is mixed with one of sodium chromate,Na<sub>2</sub>CrO<sub>4</sub>,resulting in the formation of a precipitate of lead chromate.Write a balanced net ionic equation for this precipitation reaction,showing all phases.
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Page 6
Chapter 5: Gases and the Kinetic-Molecular Theory
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Sample Questions
Q1) State Avogadro's Law.
Q2) A 20.0-L container holds 15.3 mol of Cl<sub>2</sub> gas at 227°C.
a.Calculate the pressure in atmospheres,assuming ideal behavior.
b.Calculate the pressure in atmospheres,assuming van der Waals behavior.The van der Waals constants for Cl<sub>2</sub> are a = 6.49 atm·L<sup>2</sup>/mol<sup>2</sup> and b = 0.0562 L/mol.
Q3) Nitrogen will behave most like an ideal gas
A)at high temperature and high pressure.
B)at high temperature and low pressure.
C)at low temperature and high pressure.
D)at low temperature and low pressure.
E)at intermediate (moderate)temperature and pressure.
Q4) A 255-mL gas sample weighing 0.292 g is at 52,810 Pa and 127°C.
a.How many moles of gas are present?
b.What is the molar mass of the gas?
Q5) For real gases,PV > nRT,always.
A)True
B)False
Q6) Calculate the density in g/L of gaseous SF<sub>6</sub> at 50.0°C and 650.torr.
Page 7
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Chapter 6: Thermochemistry: Energy Flow and Chemical Change
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Q1) The standard heat (enthalpy)of formation of graphite is zero.
A)True
B)False
Q2) A system initially has an internal energy E of 501 J.It undergoes a process during which it releases 111 J of heat energy to the surroundings,and does work of 222 J.What is the final energy of the system,in J?
A)168 J
B)390 J
C)612 J
D)834 J
E)It cannot be calculated without more information.
Q3) Standard heats (enthalpies)of formation of compounds,\(\Delta\)H°<sub>f</sub> ,may be positive or negative.
A)True
B)False
Q4) The standard state of a substance in aqueous solution is a 1 M solution.
A)True
B)False
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Chapter 7: Quantum Theory and Atomic Structure
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Sample Questions
Q1) Consider the following adjectives used to describe types of spectrum: Continuous line atomic emission absorption
How many of them are appropriate to describe the spectrum of radiation absorbed by a sample of mercury vapor?
A)one
B)two
C)three
D)four
E)five
Q2) A photon has an energy of 5.53 \(\times\) 10<sup>-17</sup> J.What is its frequency in s<sup>-1</sup>?
A)3.66 \(\times\) 10<sup>-50</sup> s<sup>-1</sup>
B)1.20 \(\times\) 10<sup>-17</sup> s<sup>-1</sup>
C)3.59 \(\times\) 10<sup>-9</sup> s<sup>-1</sup>
D)2.78 \(\times\) 10<sup>8</sup> s<sup>-1</sup>
E)8.35 \(\times\) 10<sup>16</sup> s<sup>-1</sup>
Q3) The energy of a photon is directly proportional to the wavelength of the radiation. A)True
B)False
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Chapter 8: Electron Configuration and Chemical Periodicity
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Q1) Write down the full electron configuration for ground state atoms of a.S (element 16).
b.Pd (element 46).
Q2) In Mendeleev's version of the periodic table,the elements were arranged in order of increasing atomic number.
A)True
B)False
Q3) Which of the following elements will form a cation with a +2 charge?
A)Si
B)Sr
C)Ga
D)Cs
E)S
Q4) Which of the following electron configurations represents the ground state for an element?
A)[Ne]3s<sup>1</sup>3p<sup>1</sup>
B)[He]2s<sup>1</sup>2p<sup>3</sup>
C)[Ne]3s<sup>2</sup>3p<sup>2</sup>3d<sup>1</sup>
D)[Ne]3s<sup>2</sup>3p<sup>3</sup>3d<sup>1</sup>
E)[Ne]3s<sup>2</sup>3p<sup>3</sup>
Page 10
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Chapter 9: Models of Chemical Bonding
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Q1) Arrange aluminum,nitrogen,phosphorus and indium in order of increasing electronegativity.
A)Al < In < N < P
B)Al < In < P < N
C)In < Al < P < N
D)In < P < Al < N
E)None of these orders is correct.
Q2) Arrange calcium,rubidium,sulfur,and arsenic in order of decreasing electronegativity.
A)S > As > Rb > Ca
B)S > As > Ca > Rb
C)As > S > Rb > Ca
D)As > S > Ca > Rb
E)None of these orders is correct.
Q3) Give a clear and concise definition of the term "electronegativity";i.e. ,what does it measure?
Q4) Bond energy increases as bond order increases,for bonding between a given pair of atoms.
A)True B)False
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Chapter 10: The Shapes of Molecules
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Q1) In which one of the following is the best Lewis structure a resonance structure?
A)SO<sub>3</sub>
B)BF<sub>3</sub>
C)I<sub>3</sub><sup>-</sup>
D)SCO (C = central atom)
E)SO<sub>3</sub><sup>2-</sup>
Q2) Using SO<sub>2</sub> as an example,describe the sort of experimental data which might suggest that no single Lewis structure is an accurate representation of its bonding.
Q3) According to VSEPR theory,a molecule with the general formula AX<sub>4</sub>E will have a _____ molecular shape.
A)bent
B)see-saw
C)trigonal planar
D)T-shaped
E)square planar
Q4) A molecule which contains polar bonds will always have a dipole moment.
A)True
B)False
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Chapter 11: Theories of Covalent Bonding
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Q1) According to molecular orbital theory,all diatomic molecules with an even number of electrons will be diamagnetic.
A)True
B)False
Q2) In the context of molecular orbital (MO)theory,explain how atomic p orbitals can give rise to:
a.a \(\sigma\) MO
b.a \(\pi\) MO
Q3) What type of hybridization is needed to explain why ethyne (acetylene,C<sub>2</sub>H<sub>2</sub>),is linear?
Q4) Valence bond theory predicts that xenon will use _____ hybrid orbitals in XeOF<sub>4</sub>.
A)sp
B)sp<sup>2</sup>
C)sp<sup>3</sup>
D)sp<sup>3</sup>d
E)sp<sup>3</sup>d<sup>2</sup>
Q5) In one sentence state the basic principle of valence bond theory.
Q6) In one sentence state how molecular orbitals are usually obtained.
Page 13
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Chapter 12: Intermolecular Forces: Liquids,solids,and Phase Changes
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Q1) In cubic closest packing,the unit cell is body-centered cubic.
A)True
B)False
Q2) A sample of octane in equilibrium with its vapor in a closed 1.0-L container has a vapor pressure of 50.0 torr at 45°C.The container's volume is increased to 2.0 L at constant temperature and the liquid/vapor equilibrium is reestablished.What is the vapor pressure?
A)> 50.0 torr
B)50.0 torr
C)25.0 torr
D)The mass of the octane vapor is needed to calculate the vapor pressure.
E)The external pressure is needed to calculate the vapor pressure.
Q3) In hydrogen iodide __________________ are the most important intermolecular forces.
A)dipole-dipole forces
B)London dispersion forces
C)hydrogen bonding
D)covalent bonds
E)polar covalent bonds
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Q4) How do the electrical properties of semiconductors differ from those of metals?
Chapter 13: The Properties of Solutions
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Q1) A solution of sucrose (sugar)in water is in equilibrium with solid sucrose.If more solid sucrose is now added,with stirring,
A)the concentration of the solution will increase.
B)the concentration of the solution will decrease.
C)the concentration of the solution will remain the same.
D)the volume of solution will increase.
E)a supersaturated solution will be produced.
Q2) Which of the following aqueous liquids will have the lowest freezing point?
A)0.5 m C<sub>12</sub>H<sub>22</sub>O<sub>11</sub> (sucrose)
B)0.5 m Ca(NO<sub>3</sub>)<sub>2</sub>
C)0.5 m NiSO<sub>4</sub>
D)0.5 m Li<sub>3</sub>PO<sub>4</sub>
E)pure water
Q3) How many moles of sulfate ions are present in 1.0 L of 0.5 M Li<sub>2</sub>SO<sub>4</sub>?
A)0.5 mol
B)1.0 mol
C)1.5 mol
D)2.0 mol
E)3.0 mol
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Chapter 14: The Main-Group Elements: Applying Principles of Bonding and Structure
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Q1) In which of the following ways is lithium different from the other alkali metals?
A)Its salts are much more soluble in water than those of the other alkali metals.
B)It has an unusually high density.
C)It forms molecular compounds with the hydrocarbon groups of organic halides.
D)Its ionization energy is lower than expected.
E)It does not react with water at room temperature.
Q2) Phosphoric acid (H<sub>3</sub>PO<sub>4</sub>)is a strong acid.
A)True
B)False
Q3) Name two different classes (types)of hydride.
Q4) Which,if any,of the following generalized formulas does not exist for interhalogen compounds? (X and Y represent different halogens. )
A)XY
B)XY<sub>2</sub>
C)XY<sub>3</sub>
D)XY<sub>5</sub>
E)All the above can represent stable interhalogen compounds.
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Chapter 15: Organic Compounds and the Atomic Properties of Carbon
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Q1) Which of the following polymers is a condensation polymer?
A)polystyrene
B)Teflon®
C)polyvinylchloride
D)polypropylene
E)Dacron®
Q2) Describe and contrast fibrous and globular proteins in terms of their amino acid composition,their structure and their function.
Q3) The carbon atoms in a molecule of cyclohexane lie in the same plane.
A)True
B)False
Q4) Explain what is meant by "complementary" in the context of DNA strands.
Q5) A characteristic reaction of alkenes is addition.
A)True
B)False
Q6) Draw and name all stable molecules with the formula C<sub>5</sub>H<sub>12</sub>.
Q8) All ketones are capable of hydrogen bonding. A)True Page 17
Q7) Draw and name two non-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 decomposition of hydrogen peroxide is a first-order process with a rate constant of 1.06 \(\times\) 10<sup>-3</sup> min<sup>-1</sup>.How long will it take for the concentration of H<sub>2</sub>O<sub>2</sub> to drop from 0.0200 M to 0.0120 M?
A)< 1 min
B)7.55 min
C)481 min
D)4550 min
E)31,400 min
Q2) The rate of a reaction is determined by the rate of the fastest step in the mechanism.
A)True
B)False
Q3) The units of the rate of reaction depend on the order of the reaction.
A)True
B)False
Q4) All bimolecular reactions are second-order reactions.
A)True
B)False
19
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Chapter 17: Equilibrium: the Extent of Chemical Reactions
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Q1) In order to write the correct mass-action expression for a reaction one must
A)know the kinetic rate law for the reaction.
B)know the mechanism for the reaction.
C)have a properly balanced chemical equation.
D)have values for the concentrations of the reactants.
E)know the limiting reactant.
Q2) For a gas-phase equilibrium,a change in the pressure of any single reactant or product will change K<sub>p</sub>.
A)True
B)False
Q3) For a solution equilibrium,a change in concentration of a reactant or product does not change K<sub>c</sub>.
A)True
B)False
Q4) Although a system may be at equilibrium,the rate constants of the forward and reverse reactions will in general be different.
A)True
B)False
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Chapter 18: Acid-Base Equilibria
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Q1) Arsenic acid,H<sub>3</sub>AsO<sub>4</sub>,is used industrially to manufacture insecticides.Arsenic acid is a polyprotic acid with K<sub>1</sub> = 2.5 \(\times\) 10<sup>-4</sup>,K<sub>2</sub> = 5.6 \(\times\) 10<sup>-8</sup>,and K<sub>3</sub> = 3 \(\times\) 10<sup>-13</sup>.What is the concentration of the HAsO<sub>4</sub><sup>2-</sup> in a solution whose initial arsenic acid concentration was 0.35 M?
A)9.4 \(\times\) 10<sup>-3</sup> M
B)2.5 \(\times\) 10<sup>-4</sup> M
C)8.8 \(\times\) 10<sup>-5</sup> M
D)5.6 \(\times\) 10<sup>-8</sup> M
E)None of these choices is correct.
Q2) Which one of the following is a Lewis acid but not a Brønsted-Lowry acid?
A)Fe<sup>3+</sup>
B)H<sub>3</sub>O<sup>+</sup>
C)HSO<sub>4</sub><sup>-</sup>
D)NH<sub>3</sub>
E)H<sub>2</sub>O
Q3) Formic acid,HCOOH,is a weak acid with a pK<sub>a</sub> of 3.74.Draw up a reaction table for the reaction of 0.300 M formic acid with water,and calculate the pH of this solution.
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Chapter 19: Ionic Equilibria in Aqueous Systems
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Q1) A 50.0-mL sample of 0.50 M HCl is titrated with 0.50 M NaOH.What is the pH of the solution after 28.0 mL of NaOH have been added to the acid?
A)0.85
B)0.75
C)0.66
D)0.49
E)3.8
Q2) Calculate the solubility of lead(II)iodide,PbI<sub>2</sub>,in 0.025 M KI.K<sub>sp</sub> = 7.9 \(\times\) 10<sup>-9</sup>
A)4.5 \(\times\) 10<sup>-2</sup> M
B)2.8 \(\times\) 10<sup>-2</sup> M
C)8.9 \(\times\) 10<sup>-5</sup> M
D)5.0 \(\times\) 10<sup>-5</sup> M
E)1.3 \(\times\) 10<sup>-5</sup> M
Q3) What is the pH of 375 mL of solution containing 0.150 mol of propenoic acid (HA)and 0.250 mol of sodium propenoate (NaA)? (K<sub>a</sub> for propenoic acid is 5.52 \(\times\) 10<sup>-5</sup>. )
Q4) Make a clear distinction between buffer range and buffer capacity.
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Chapter
Direction of Chemical Reactions
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Q1) Under a given set of conditions,all microstates of a system are equally probable.
A)True
B)False
Q2) Which relationship or statement best describes \(\Delta\)S° for the following reaction?
O<sub>3</sub>(g)+ NO(g) \(\to\) O<sub>2</sub>(g)+ NO<sub>2</sub>(g)
A)(\(\Delta\)S° \(\approx\) 0)
B)(\(\Delta\)S° < 0)
C)(\(\Delta\)S° > 0)
D)(\(\Delta\)S° = \(\Delta\)H°/T)
E)More information is needed to make a reasonable prediction.
Q3) Which of the following is always true for an exothermic process?
A)q<sub>sys</sub> > 0,\(\Delta\)S<sub>surr</sub> < 0
B)q<sub>sys</sub> < 0,\(\Delta\)S<sub>surr</sub> > 0
C)q<sub>sys</sub> < 0,\(\Delta\)S<sub>surr</sub> < 0
D)q<sub>sys</sub> > 0,\(\Delta\)S<sub>surr</sub> > 0
E)w < 0
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) What mass of copper will be deposited when 18.2 A are passed through a CuSO<sub>4</sub> solution for 45.0 minutes?
A)16.2 g
B)33.4 g
C)40.6 g
D)81.3 g
E)163 g
Q2) A voltaic cell is prepared using copper and silver.Its cell notation is shown below. Cu(s)| Cu<sup>2+</sup>(aq)|| Ag<sup>+</sup>(aq)| Ag(s) Which of the following processes occurs at the cathode?
A)Cu(s) \(\to\) Cu<sup>2+</sup>(aq)+ 2e<sup>-</sup>
B)Cu<sup>2+</sup>(aq)+ 2e<sup>-</sup> \(\to\) Cu(s) C)Ag(s) \(\to\) Ag<sup>+</sup>(aq)+ e<sup>-</sup>
D)Ag<sup>+</sup>(aq)+ e<sup>-</sup> \(\to\) Ag(s)
E)Cu(s)+ 2Ag<sup>+</sup>(aq) \(\to\) Cu<sup>2+</sup>(aq)+ 2Ag(s)
Q3) A secondary cell (battery)can operate either as a galvanic or an electrolytic cell.
A)True B)False
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Chapter 22: The Transition Elements and Their Coordination Compounds
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Q1) The most common oxidation state for ions of the transition elements is
A)+2
B)+3
C)+4
D)+5
E)+6
Q2) Which of the following octahedral complexes should have the largest crystal field splitting energy,\(\Delta\)?
A)[Cr(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup>
B)[Cr(SCN)<sub>6</sub>]<sup>3-</sup>
C)[Cr(NH<sub>3</sub>)<sub>6</sub>]<sup>3+</sup>
D)[Cr(CN)<sub>6</sub>]<sup>3-</sup>
E)[Cr(en)<sub>3</sub>]<sup>3+</sup> (en = ethylenediamine)
Q3) What is the highest possible oxidation state for palladium,Pd?
A)+1
B)+2
C)+3 D)+4
E)+6
Q4) Why is the +2 oxidation state so common among transition elements?
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Chapter 23: Nuclear Reactions and Their Applications
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Q1) Write a complete,balanced equation to represent the alpha decay of radon-210.
Q2) Calculate to four significant figures
a.the mass defect in kg and b.the energy released in kJ/mol,when a neutron decays to produce a proton and an electron.The neutron,proton and electron masses are 1.67493 \(\times\) 10<sup>-27</sup> kg,1.67262 \(\times\) 10<sup>-27</sup> kg and 9.10939 \(\times\) 10<sup>-31</sup> kg,respectively.
Q3) Positron decay and electron capture have the same net effect on the Z and N values of a nucleus.
A)True
B)False
Q4) Which of the following materials is put into a nuclear reactor to slow the chain reaction?
A)heavy water
B)moderators
C)control rods
D)reflectors
E)chlorine
Q5) Write a complete,balanced equation to represent the beta decay of thallium-207.
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