Course Introduction
Introduction to Chemistry Study Guide
Questions
Introduction to Chemistry provides a foundational understanding of the principles and concepts that form the basis of chemical science. This course covers essential topics such as atomic structure, chemical bonding, the periodic table, stoichiometry, chemical reactions, states of matter, solutions, and basic thermodynamics. Students will gain practical laboratory skills and develop problem-solving abilities through experiments and exercises designed to illustrate key chemical concepts. By the end of the course, learners will be equipped with the knowledge needed to pursue more advanced studies in chemistry and related scientific fields.
Recommended Textbook
General Chemistry The Essential Concepts 7th Edition by Raymond Chang
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Page 2
Chapter 1: Introduction
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Sample Questions
Q1) A pure yellow crystalline substance, when heated in a vacuum, releases a greenish gas and a red powder. Is the original yellow crystalline substance a compound or element?
Answer: Compound
Q2) Complete the following sentence. A hypothesis is
A)a tentative explanation for a set of observations that can be tested by further experimentation.
B)a statement describing a relationship between phenomena that is always the same under the same conditions.
C)a unifying principle that explains a body of facts and relations.
D)a model used to visualize the invisible.
Answer: A
Q3) Which of the following is an example of a physical property?
A)corrosiveness of sulfuric acid
B)toxicity of cyanide
C)flammability of gasoline
D)neutralization of stomach acid with an antacid
E)lead becomes a liquid when heated to 601<sup>\(\circ\)</sup>C
Answer: E
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Chapter 2: Atoms, Molecules, and Ions
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Sample Questions
Q1) The Stock system name for Co<sub>2</sub>(SO<sub>3</sub>)<sub>3</sub> is:
A)cobalt sulfate
B)cobalt(II)sulfite
C)cobalt(II)sulfate
D)cobalt(III)sulfite
E)cobalt(III)sulfate
Answer: D
Q2) The chemical formula for iron(II)nitrate is:
A)Fe<sub>2</sub>(NO<sub>3</sub>)<sub>3</sub>
B)Ir(NO<sub>2</sub>)<sub>2</sub>
C)Fe<sub>2</sub>N<sub>3</sub>
D)Fe(NO<sub>3</sub>)<sub>2</sub>
E)Fe(NO<sub>2</sub>)<sub>2</sub>
Answer: D
Q3) Using a cathode ray tube, J. J. Thomson determined the magnitude of the electric charge on the electron.
A)True
B)False
Answer: False
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Page 4
Chapter 3: Stoichiometry
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Sample Questions
Q1) What is the mass of 0.20 mole of C<sub>2</sub>H<sub>6</sub>O (ethanol)?
A)230 g
B)46 g
C)23 g
D)4.6 g
E)None of these.
Answer: E
Q2) An average atom of uranium (U)is approximately how many times heavier than an atom of potassium?
A)6.1 times
B)4.8 times
C)2.4 times
D)12.5 times
E)7.7 times
Answer: A
Q3) What percent by mass of oxygen is present in carbon monoxide, CO?
Answer: 57%
Q4) Calculate the molecular mass, in g/mol, of P<sub>4</sub>O<sub>10</sub>.
Answer: 283.9 g/mol
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Chapter 4: Reactions in Aqueous Solution
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Sample Questions
Q1) Give an example of a monoprotic acid.
Q2) Which of the following compounds is a nonelectrolyte?
A)NaF
B)HNO<sub>3</sub>
C)CH<sub>3</sub>COOH (acetic acid)
D)NaOH
E)C<sub>6</sub>H<sub>12</sub>O<sub>6 </sub>(glucose)
Q3) A piece of copper metal was added to an aqueous solution of silver nitrate, and within a few minutes it was observed that a grey crystalline solid formed on surface of the copper and the solution turned a blue color characteristic of copper(II)ions. Write the balanced chemical equation for this reaction.
Q4) During a titration the following data were collected. A 10. mL portion of an unknown monoprotic acid solution was titrated with 1.0 M NaOH; 40. mL of the base were required to neutralize the sample. What is the molarity of the acid solution?
Q5) Identify the reducing agent in the following reaction. 4Al + 3O<sub>2 </sub>\(\rarr\) 2Al<sub>2</sub>O<sub>3</sub>
Q6) Identify the following as either a good or poor conductor of electricity: a crystal of Na<sub>2</sub>SO<sub>4</sub>.
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Chapter 5: Gases
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Sample Questions
Q1) Calculate the volume of H<sub>2</sub>(g)at 273 K and 2.00 atm that will be formed when 275 mL of 0.725 M HCl solution reacts with excess Mg to give hydrogen gas and aqueous magnesium chloride.
A)0.56 L
B)1.12 L
C)2.23 L
D)4.47 L
E)3.54 L
Q2) Ammonium nitrite undergoes decomposition to produce only gases as shown below.
NH<sub>4</sub>NO<sub>2</sub>(s)\(\rarr\) N<sub>2</sub>(g)+ 2H<sub>2</sub>O(g) How many liters of gas will be produced by the decomposition of 32.0 g of NH<sub>4</sub>NO<sub>2</sub> at 525°C and 1.5 atm?
Q3) What volume of H<sub>2</sub> is formed at STP when 6.0 g of Al is treated with excess NaOH?
2NaOH + 2Al + 6H<sub>2</sub>O \(\rarr\) 2NaAl(OH)<sub>4</sub> + 3H<sub>2</sub>(g)
Q4) At standard temperature and pressure, a given sample of water vapor occupies a volume of 2.80 L. How many moles of water vapor are present?
Q5) What is the definition of a "gas"?
Page 7
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Chapter 6: Energy Relationships in Chemical Reactions
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Sample Questions
Q1) An average home in Colorado requires 20. GJ of heat per month. How many grams of natural gas (methane)must be burned to supply this energy? CH<sub>4</sub>(g)+ 2O<sub>2</sub>(g) \(\rarr\)CO<sub>2</sub>(g)+ 2H<sub>2</sub>O(l)\(\Delta\)H°<sub>rxn</sub>= -890.4 kJ/mol
A)1.4 * 10<sup>3</sup> g
B)3.6 * 10<sup>5</sup> g
C)7.1 * 10<sup>-4</sup> g
D)2.2 * 10<sup>4</sup> g
E)1.4 * 10<sup>4</sup> g
Q2) The combustion of octane produces heat according to the equation 2C<sub>8</sub>H<sub>18</sub>(l)+
25O<sub>2</sub>(g)\(\rarr\)16CO<sub>2</sub>(g)+
18H<sub>2</sub>O(l)\(\Delta\)H°<sub>rxn</sub>= -11,020 kJ/mol
What is the heat of combustion per gram of octane?
A)-5,510 kJ/g
B)-96.5 kJ/g
C)-48.2 kJ/g
D)-193 kJ/g
E)-6.292 * 10<sup>5</sup> kJ/g
Q3) What is the standard enthalpy of formation of H<sub>2</sub>(g)at 25°C?
Q4) Define specific heat.
Page 8
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Chapter 7: The Electronic Structure of Atoms
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Sample Questions
Q1) Write the ground state electron configuration for a lead atom.
Q2) Which of the following is the electron configuration of an excited state of a copper atom?
A)[Ar]4s<sup>2</sup>3d<sup>9</sup>
B)[Ar]4s<sup>1</sup>3d<sup>10</sup>
C)[Ar]4s<sup>1</sup>3d<sup>8</sup>
D)[Ar]4s<sup>2</sup>3d<sup>8</sup>
E)[Ar]4s<sup>0</sup>3d<sup>10</sup>
Q3) The ground-state electron configuration of Cr, Mo, and Ag are exceptions to the Aufbau principle. Which of the following is the electron configuration for Mo?
A)[Kr]5s<sup>1</sup>4d<sup>5</sup>
B)[Kr]5s<sup>2</sup>4d<sup>4</sup>
C)[Xe]6s<sup>2</sup>5d<sup>4</sup>
D)[Ar]4s<sup>2</sup>4d<sup>4</sup>
E)[Kr]5s<sup>2</sup>4d<sup>6</sup>
Q4) The colors of the visible spectrum are blue, green, orange, red, violet, and yellow. Of these colors, _______ has the longest wavelength.
Q5) Write the ground state electron configuration for Ni.
Q6) Write the ground state electron configuration for Cr.
Page 9
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Chapter 8: The Periodic Table
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Sample Questions
Q1) Which of the following is a basic oxide?
A)CO<sub>2</sub>
B)MgO
C)As<sub>2</sub>O<sub>3</sub>
D)SO<sub>2</sub>
E)Cl<sub>2</sub>O<sub>7</sub>
Q2) What is the charge on the stable ion formed by selenium?
A)+2
B)+1
C)-1
D)-2
E)-3
Q3) Consider the following reaction: 3Li + Z \(\rarr\) Li<sub>3</sub>Z. What is the formula for the compound if we substitute magnesium for lithium?
A)MgZ
B)Mg<sub>2</sub>Z
C)MgZ<sub>2</sub>
D)Mg<sub>3</sub>Z
E)Mg<sub>3</sub>Z<sub>2</sub>
Q4) Write the ground-state electron configuration for Mg<sup>2+</sup>.
Page 10
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Chapter 9: Chemical Bonding I: the Covalent Bond
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Sample Questions
Q1) Which one of the following is most likely to be a covalent compound?
A)Rb<sub>2</sub>O
B)BaO
C)SrO
D)SeO<sub>2</sub>
E)MnO<sub>2</sub>
Q2) The bond in which of the following pairs of atoms would have the greatest percent ionic character (i.e., most polar)?
A)C - O
B)S - O
C)Na - I
D)Na - Br
E)F - F
Q3) The Si - Cl bond has less ionic character than the C - Cl bond.
A)True
B)False
Q4) Write a Lewis structure for the chlorate ion, ClO<sub>3</sub><sup>-</sup>, that obeys the octet rule, showing all non-zero formal charges, and give the total number of resonance structures for ClO<sub>3</sub><sup>-</sup> that obey the octet rule.
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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization
of Atomic Orbitals
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Sample Questions
Q1) Give the number of lone pairs around the central atom and the molecular geometry of XeF<sub>2</sub>.
A)0 lone pairs, linear
B)1 lone pair, bent
C)2 lone pairs, bent
D)3 lone pairs, bent
E)3 lone pairs, linear
Q2) Which one of the following molecules has tetrahedral geometry?
A)XeF<sub>4</sub>
B)BF<sub>3</sub>
C)AsF<sub>5</sub>
D)CF<sub>4</sub>
E)NH<sub>3</sub>
Q3) The bond angle in ICl<sub>2</sub><sup>-</sup> is expected to be
A)a little less than 109.5°.
B)109.5°.
C)a little more than 109.5°.
D)120°.
E)180°.
Q4) Indicate the number of \(\pi\)-bonds in N<sub>2</sub>H<sub>2</sub>.
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Chapter 11: Introduction to Organic Chemistry
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Sample Questions
Q1) Oxidation of the 2-propanol will produce a/an
A)aldehyde.
B)amine.
C)alkene.
D)ketone.
E)carboxylic acid.
Q2) Unsaturated hydrocarbons
A)contain at least one double or triple carbon-carbon bond.
B)contain at least one element other than hydrogen and carbon.
C)contain the maximum number of hydrogens that can bond with the carbon atoms present.
D)cannot form structural isomers.
E)cannot undergo addition reactions.
Q3) Which of these reactions leads to a change in the hybridization of one or more carbon atoms?
A)free radical halogenation of an alkane
B)hydrolysis of an ester to yield an acid and an alcohol
C)substitution of an aromatic ring using a halogen
D)oxidation of an alcohol to yield a carboxylic acid
E)neutralization of an amine using a strong mineral acid
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Chapter 12: Intermolecular Forces and Liquids and Solids
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Sample Questions
Q1) Vanadium crystallizes in a body-centered cubic lattice, and the length of the edge of a unit cell is 305 pm. What is the density of V?
A)5.96 g/cm<sup>3</sup>
B)2.98 g/cm<sup>3</sup>
C)2.98 * 10<sup>-6</sup> g/cm<sup>3</sup>
D)5.96 * 10<sup>-30</sup> g/cm<sup>3</sup>
E)11.9 g/cm<sup>3</sup>
Q2) Which liquid is expected to have the larger surface tension at a given temperature, CCl<sub>4</sub> or H<sub>2</sub>O? Briefly explain.
Q3) Platinum has a face-centered cubic crystal structure and a density of 21.5 g/cm<sup>3</sup>. What is the radius of the platinum atom?
A)69 pm
B)98 pm
C)139 pm
D)196 pm
E)277 pm
Q4) Iron crystallizes in a body-centered cubic unit. The edge of this cell is 287 pm. Calculate the density of iron.
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Chapter 13: Physical Properties of Solutions
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Sample Questions
Q1) At 10°C one volume of water dissolves 3.10 volumes of chlorine gas at 1.00 atm pressure. What is the Henry's Law constant in mol/L·atm?
A)3.8
B)0.043
C)36.
D)3.1
E)0.13
Q2) What is the percent by mass of sodium phosphate in a 0.142 M Na<sub>3</sub>PO<sub>4</sub>(aq)solution that has a density of 1.015 g/mL?
Q3) Calculate the mass of solute in the following solution: 50.0 mL of 0.0300 M C<sub>12</sub>H<sub>22</sub>O<sub>11</sub>.
Q4) A solution is prepared by adding 40.3 g of Mg(NO<sub>3</sub>)<sub>2</sub> to 127 g of water. Calculate the mole fraction and molality of magnesium nitrate in this solution.
Q5) 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?
Q6) Define solvation.
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Chapter 14: Chemical Kinetics
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Sample Questions
Q1) An increase in the temperature of the reactants causes an increase in the rate of reaction. The best explanation for this behavior is that as the temperature increases,
A)the concentration of reactants increases.
B)the activation energy decreases.
C)the collision frequency increases.
D)the fraction of collisions with total kinetic energy greater than E<sub>a</sub> increases.
E)the activation energy increases.
Q2) For the following reaction, \(\Delta\)P(C<sub>6</sub>H<sub>14</sub>)/\(\Delta\)t was found to be -6.2 * 10<sup>-3</sup> atm/s. C<sub>6</sub>H<sub>14</sub>(g)\(\rarr\) C<sub>6</sub>H<sub>6</sub>(g)+ 4H<sub>2</sub>(g)
Determine \(\Delta\)P(H<sub>2</sub>)/\(\Delta\)t for this reaction at the same time.
A)6.2 * 10<sup>-3</sup> atm/s
B)1.6 * 10<sup>-3</sup> atm/s
C)2.5 * 10<sup>-2</sup> atm/s
D)-1.6 * 10<sup>-3</sup> atm/s
E)-2.5 * 10<sup>-2</sup> atm/s
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Chapter 15: Chemical Equilibrium
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Q1) Equilibrium is established for the reaction 2X(s)+ Y(g)\(\rarr\)<sub> </sub>2Z(g)at 500K, K<sub>c</sub> = 100. Determine the concentration of Z in equilibrium with 0.2 mol X and 0.50 M Y at 500K.
A)3.2 M
B)3.5 M
C)4.5 M
D)7.1 M
E)None of these.
Q2) 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?
Q3) 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.
Q4) Describe why addition of a catalyst does not affect the equilibrium constant for a reaction.
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Chapter 16: Acids and Bases
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Q1) The pOH of a solution is 10.40. Calculate the hydrogen ion concentration in the solution.
A)4.0* 10<sup>-11</sup> M
B)3.6 M
C)4.0 * 10<sup>-10</sup> M
D)2.5 * 10<sup>-4</sup> M
E)1.8 * 10<sup>-4</sup> M
Q2) The equilibrium constant for the reaction C<sub>7</sub>H<sub>15</sub>COOH(aq)+ HCOO<sup>-</sup>(aq)\(\rarr\)
C<sub>7</sub>H<sub>15</sub>COO<sup>-</sup>(aq)+ HCOOH(aq)
Is 7.23 * 10<sup>-2</sup> at 25°C. If K<sub>a</sub> for formic acid (HCOOH)is 1.77 * 10<sup>-4</sup>, what is the acid dissociation constant for C<sub>7</sub>H<sub>15</sub>COOH?
A)2.45 * 10<sup>-3</sup>
B)4.08 * 10<sup>-2</sup>
C)7.81 * 10<sup>-4</sup>
D)1.00 * 10<sup>-4</sup>
E)1.28 * 10<sup>-5</sup>
Q3) Al(OH)<sub>3</sub> is an amphoteric hydroxide. Write a balanced ionic equation to show its reaction with HNO<sub>3</sub>.
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Chapter 17: Acid-Base Equilibria and Solubility Equilibria
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Q1) The molar solubility of tin(II)iodide is 1.28 * 10<sup>-2</sup> mol/L. What is K<sub>sp</sub> for this compound?
A)8.4 * 10<sup>-6</sup>
B)1.28 * 10<sup>-2</sup>
C)4.2 * 10<sup>-6</sup>
D)1.6 * 10<sup>-4</sup>
E)2.1 * 10<sup>-6</sup>
Q2) An environmental chemist obtained a 200. mL sample of lake water believed to be contaminated with a single monoprotic strong acid. Titrating this sample with a 0.0050 M NaOH(aq)required 7.3 mL of the NaOH solution to reach the endpoint. What is the concentration of H<sup>+</sup> in the lake?
Q3) The solubility of lead(II)iodide is 0.064 g/100 mL at 20<sup>o</sup>C. What is the solubility product for lead(II)iodide?
A)1.1 * 10<sup>-8</sup>
B)3.9 * 10<sup>-6</sup>
C)1.1 * 10<sup>-11</sup>
D)2.7 * 10<sup>-12</sup>
E)1.4 * 10<sup>-3</sup>
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Chapter 18: Thermodynamics
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Q1) For the reaction 3H<sub>2</sub>(g)+ N<sub>2</sub>(g) \(\rarr\) 2NH<sub>3</sub>(g), K<sub>c</sub> = 9.0 at 350°C. In what direction does this reaction proceed at 350°C under standard state conditions?
Q2) Dissolving an ionic solid in water always results in an increase in entropy.
A)True
B)False
Q3) In the gas phase, formic acid forms a dimmer, 2HCOOH(g) \(\rarr\) (HCOOH)<sub>2</sub>(g). For this reaction, \(\Delta\)H° = -60.1 kJ/mol and \(\Delta\)G° = -13.9 kJ/mol at 25°C. Find the equilibrium constant (K<sub>p</sub>)for this reaction at 75 °C.
A)8960
B)273
C)0.120
D)8.33
E)1.12 * 10<sup>-4</sup>
Q4) For the reaction H<sub>2</sub>O<sub>2</sub>(g) \(\rarr\) H<sub>2</sub>O(g)+ <sup>1</sup>/<sub>2</sub>O<sub>2</sub>(g), \(\Delta\)H° = -106 kJ/mol and \(\Delta\)S° = 58 J/K·mol at 25°C. Calculate \(\Delta\)G° for this reaction at this temperature.
Q5) How does the entropy change when a molecular solid is dissolved in water?
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Chapter 19: Redox Reactions and Electrochemistry
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Q1) How many faradays are transferred in an electrolytic cell when a current of 2.0 amperes flows for 12 hours?
A)24 F
B)8.6 * 10<sup>4</sup> F
C)0.90 F
D)6.2 * 10<sup> -3</sup> F
E)1.1 F
Q2) For the reaction Ni<sup>2+</sup>(aq)+ 2Fe<sup>2+</sup>(aq)\(\rarr\) 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) How many grams of copper are deposited on the cathode of an electrolytic cell if an electric current of 2.00 A is passed through a solution of CuSO<sub>4</sub> for a period of 19.0 min?
Q4) Complete and balance the following redox reaction under acidic conditions: ClO<sub>2</sub><sup>-</sup>(aq)\(\rarr\) ClO<sub>2</sub>(g)+ Cl<sup>-</sup>(aq)
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Chapter 20: The Chemistry of Coordination Compounds
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Q1) Write the chemical formula of the dibromobis(oxalato)cobaltate(III)ion.
Q2) 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
Q3) What terms describe the geometric isomers that are possible for the complex [CrF<sub>2</sub>Cl<sub>4</sub>]<sup>3-</sup>?
Q4) In the complex ion [Fe(CN)<sub>6</sub>]<sup>4-</sup>, the oxidation number of Fe is A)+1. B)+2. C)+3. D)-4. E)+6.
Q5) A molecule or atom that accepts an electron pair to form a coordinate covalent bond is called a Lewis _______.
Page 22
Q6) Write the chemical formula of triammineaquodichlorochromium(III)chloride.
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Chapter 21: Nuclear Chemistry
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Q1) How many <sup>14</sup>C atoms are in a charcoal sample that has a decay rate of 3,500 disintegrations per min? (For <sup>14</sup>C, t<sub>1/2</sub> = 5,730 yr.)
A)2.9 * 10<sup>7 </sup>atoms
B)8.0 * 10<sup>-7 </sup>atoms
C)1.4 * 10<sup>14 </sup>atoms
D)1.5 * 10<sup>13 </sup>atoms
E)6.02 * 10<sup>20 </sup>atoms
Q2) What role does cadmium metal (Cd)play in a nuclear reactor?
A)slows down the fission neutrons (moderator)
B)transfers heat from the reactor to the heat exchanger (primary coolant)
C)controls chain reaction (control rods)
D)transfers heat from the condenser to the environment (cooling tower)
E)undergoes fission (fuel rods)
Q3) Tritium is a radioisotope of hydrogen having a half-life of 12.3 years. If you initially had 1.0 * 10<sup>-7</sup> mole of tritium, how many moles of tritium would be left after 78 years?
Q4) Protactinium-234 has a half-life of 1 minute. How much of a 400. g sample protactinium would remain after 4 minutes?
Q5) What nuclear fuel is produced in a breeder reactor?
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Chapter 22: Organic Polymerssynthetic and Natural
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Source URL: https://quizplus.com/quiz/63763
Sample Questions
Q1) Which of the following polymers is formed by a condensation process?
A)PVC
B)nylon
C)Teflon
D)Plexiglas
E)neoprene
Q2) Hydrogen bonding, dispersion forces, ionic forces, and dipole-dipole forces all affect the structure of a protein?
A)True
B)False
Q3) Both DNA and RNA have double-helical structures.
A)True
B)False
Q4) Which choice lists both the sugar and the nitrogen base that are a part of RNA but are not a part of DNA?
A)deoxyribose and thymine
B)ribose and deoxyribose
C)ribose and uracil
D)uracil and thymine
E)ribose and thymine
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