Survey of Chemistry Exam Answer Key - 3115 Verified Questions

Survey of Chemistry

Exam Answer Key

Course Introduction

Survey of Chemistry provides a broad overview of the fundamental concepts and principles of chemistry, tailored for students seeking a foundational understanding of the subject. The course covers key topics such as atomic and molecular structure, chemical bonding, periodic trends, stoichiometry, states of matter, solutions, acids and bases, and basic chemical reactions. Emphasis is placed on the practical applications of chemistry in everyday life and across various scientific disciplines. Designed for non-science majors or those preparing for further studies, the course balances theoretical knowledge with hands-on laboratory experiences to reinforce core concepts and develop analytical skills.

Recommended Textbook

Chemistry 11th Edition by Raymond Chang

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25 Chapters

3115 Verified Questions

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Chapter 1: Chemistry: the Study of Change

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Sample Questions

Q1) The city of Los Angeles is now approximately 2400 miles south of Alaska.It is moving slowly northward as the San Andreas fault slides along.If Los Angeles is to arrive near Anchorage, Alaska, in 76 million years, at what average rate will it have to move in mm per month?

A)2.0 × 10^-10 mm/mo.

B)6.6 × 10^-6 mm/mo.

C)4.2 mm/mo.

D)9.5 mm/mo.

E)51 mm/mo.

Answer: C

Q2) Classify the following as an element, a compound, or a mixture: Table salt (non-iodized).

Answer: Compound

Q3) A unifying principle that explains a body of facts and relations is referred to as A)a hypothesis.

B)a law.

C)a theory.

D)none of the above.

Answer: C

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Page 3

Chapter 2: Atoms, Molecules, and Ions

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Sample Questions

Q1) Name the compound Co₂(SO₃)₃ .

A)cobalt sulfate

B)cobalt(II)sulfite

C)cobalt(II)sulfate

D)cobalt(III)sulfite

E)cobalt(III)sulfate

Answer: D

Q2) Which of the following scientists developed the nuclear model of the atom?

A)John Dalton

B)Robert Millikan

C)J.J.Thomson

D)Henry Moseley

E)Ernest Rutherford

Answer: E

Q3) Give the formula of carbonic acid.

Answer: H₂CO₃

Q4) What is the formula for the compound hydrogen peroxide

Answer: H₂O₂

Q5) Name the compound CH₃CH₂NH₂

Answer: Ethylamine

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 O₂ when the following equation is properly balanced? ___ CH₃OH + ___ O₂ \(\rarr\) ___ CO₂ + ___ H₂O

A)1

B)2

C)3

D)7

E)none of these

Answer: C

Q2) Balance the following and list the coefficients in order from left to right. ___ Cr + ___ H₂SO₄ \(\rarr\) ___ Cr₂(SO₄)₃+ ___ H₂

A)2, 3, 1, 2

B)2, 3, 1, 3

C)1, 3, 1, 3

D)4, 6, 2, 6

E)1, 3, 1, 2

Answer: B

Q3) How many ICl₃ molecules are present in 1.75 kg of ICl₃?

Answer: 4.52 × 10²⁴

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Chapter 4: Reactions in Aqueous Solutions

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Sample Questions

Q1) Which of the following ions is a weak acid?

A)SO₄^2-

B)H₂SO₄

C)HSO₄^-

D)HNO₃

E)NO₃^-

Q2) Hydrogen is oxidized in the following chemical reaction. H₂ + Cl₂\(\rarr\)2HCl

A)True

B)False

Q3) During a titration the following data were collected.A 50.0 mL portion of an HCl solution was titrated with 0.500 M NaOH; 200.mL of the base was required to neutralize the sample.How many grams of HCl are present in 500.mL of this acid solution?

Q4) 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.The solution turned a blue color characteristic of copper(II)ions.Write the net ionic equation for this reaction.

Q5) What is the molarity of a solution that contains 5.0 moles of solute in 2.00 liters of solution?

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Chapter 5: Gases

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Sample Questions

Q1) A gas-filled balloon with a volume of 12.5 L at 0.90 atm and 21°C is allowed to rise to the stratosphere where the temperature is -5°C and the pressure is 1.0 millibar.What is the final volume of the balloon in Liters? 1.000 atm = 1.013 bar.

Q2) Give five examples (list names)of compounds that exist as gases at room temperature and pressure.

Q3) Define Avogadro's Law

Q4) A 0.271 g sample of an unknown vapor occupies 294 mL at 140.°C and 847 mmHg.The empirical formula of the compound is CH₂.What is the molecular formula of the compound?

A)CH₂

B)C₂H₄

C)C₃H₆

D)C₄H₈

E)C₆H₁₂

Q5) Give five examples (list names)of elements that occur as gases at room temperature and pressure?

Q6) At STP, 1 mole of gas has a molar volume of 22.4 L.What is the density (g/L)of oxygen at STP?

Q7) What is the significance of the magnitude of the van der Waals "a" constant?

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Chapter 6: Thermochemistry

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Sample Questions

Q1) For which of these reactions will the difference between \(\Delta\)H° and \(\Delta\)E° be the smallest?

A)N₂(g)+ 3H₂(g)\(\rarr\)2NH₃(g)

B)4PH₃(g)\(\rarr\) P₄(g)+ 6H₂(g)

C)H₂(g)+ Cl₂(g)\(\rarr\) 2HCl(g)

D)CO₂(g)+ 2H₂O(l)\(\rarr\) CH₄(g)+ 2O₂(g)

E)P₄(s)+ 10Cl₂(g)\(\rarr\) 4PCl₅(s)

Q2) Calculate the amount of work done, in joules, when 2.5 mole of H₂O vaporizes at 1.0 atm and 25°C.Assume the volume of liquid H₂O is negligible compared to that of vapor.(1 L·atm = 101.3 J)

A)61.1 J

B)518 J

C)5.66 kJ

D)6.19 kJ

E)6,190 kJ

Q3) Chemical reactions in a bomb calorimeter occur at constant pressure.

A)True

B)False

Q4) Define specific heat.

Page 8

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Chapter 7: Quantum Theory and the Electronic Structure of

Atoms

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Sample Questions

Q1) If a hydrogen atom and a helium atom are traveling at the same speed,

A)the wavelength of the hydrogen atom will be about 4 times longer than the wavelength of the helium atom.

B)the wavelength of the hydrogen atom will be about 2 times longer than the wavelength of the helium.

C)the wavelength of the hydrogen atom will be roughly equal to the wavelength of the helium atom.

D)the wavelength of the helium atom will be about 2 times longer than the wavelength of the hydrogen atom.

E)the wavelength of the helium atom will be about 4 times longer than the wavelength of the hydrogen atom.

Q2) Calculate the frequency of visible light having a wavelength of 686 nm.

A)4.37 × 10¹⁴/s

B)4.37 × 10⁵/s

C)6.17 × 10¹⁴/s

D)2.29 × 10^-15/s

E)2.29 × 10^-6/s

Q3) Write the ground state electron configuration for a lead atom.

Q4) What is the total number of electrons possible in the 2p orbitals?

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Chapter 8: Periodic Relationships Among the Elements

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Sample Questions

Q1) As opposed to early periodic tables based on the law of octaves, modern periodic tables arrange the elements in order of increasing

A)nuclear binding energy.

B)number of neutrons.

C)atomic mass.

D)atomic number.

E)atomic size.

Q2) The sulfide ion, S^2-, is isoelectronic with which one of the following?

A)O^2-

B)F^-

C)Na⁺

D)Al³⁺

E)K⁺

Q3) How many 3d electrons does an Fe³⁺ ion have?

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Chapter 9: Chemical Bonding I: Basic Concepts

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Sample Questions

Q1) Carbonic acid, H₂CO₃, is a weak acid that contributes to the taste and produces the carbon dioxide bubbles in all carbonated beverages.Write a Lewis structure for H₂CO₃,

Q2) List all types of bonding present in the compound NH₄NO₃ I.ionic bond

II.polar covalent bond

III.nonpolar covalent bond

A)I only

B)II only

C)III only

D)I and II

E)II and III

Q3) Write a Lewis structure for the phosphate ion, PO₄^3-, that expands the octet to minimize formal charge and if necessary places negative formal charges on the most electronegative atom(s).

Q4) Write the chemical equation for which the enthalpy of reaction is the lattice energy of KCl(s).

Q5) Write the Lewis structure of boron trifluoride.

Q6) Write the Lewis dot symbol for the sulfide ion.

Page 11

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Chapter 10: Chemical Bonding Ii: Molecular Geometry and Hybridization

of Atomic Orbitals

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Sample Questions

Q1) An sp² hybridized central carbon atom with no lone pairs of electrons has what type of bonding?

A)1 \(\pi\) and 2 \(\sigma\) bonds

B)1 \(\pi\) and 3 \(\sigma\) bonds

C)2 \(\pi\) and 2 \(\sigma\) bonds

D)3 \(\pi\) and 2 \(\sigma\) bonds

E)0 \(\pi\) and 4 \(\sigma\) bonds

Q2) Two p_y orbitals from two different atoms can interact to form a pi bonding molecular orbital.

A)True

B)False

Q3) Indicate the type of hybrid orbitals used by the central atom in BrF₃.

A)sp

B)sp²

C)sp³

D)sp³d

E)sp³d²

Q4) According to the VSEPR theory, the geometrical structure of PF₅ is

Q5) Indicate the number of \(\pi\)-bonds in N₂H₂.

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Chapter 11: Intermolecular Forces and Liquids and Solids

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Sample Questions

Q1) Which of the following would be expected to have the lowest vapor pressure at room temperature?

A)ethanol, bp = 78°C

B)methanol, bp = 65°C

C)water, bp = 100°C

D)acetone, bp = 56°C

Q2) Which is expected to have a higher vapor pressure, C₈H₁₈ or C₄H₁₀?

Q3) Boron nitride, BN₃, melts at approximately at 3,000°C under high pressure.This material is almost as hard as diamond.What kind of crystal is this?

Q4) The normal boiling point of bromine is 58.8°C.Given that the vapor pressure of bromine is 75.0 torr at 2.5°C, calculate the molar enthalpy of vaporization of bromine.

A)2.90 kJ/mol

B)3.57 kJ/mol

C)3.76 kJ/mol

D)29.7 kJ/mol

E)31.3 kJ/mol

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) What is the mole fraction of NaOH in a 32.0 % by mass NaOH aqueous solution?

Q2) Calculate the boiling point of a 4.5 m solution of Na₂SO₄ in water (assume 100% dissociation, K_b (H₂O)= 0.52 °C/m)

A)93.0 °C

B)97.7 °C

C)102.3°C

D)104.7°C

E)107.0°C

Q3) A solution made of pentane and hexane has a mole fraction, X = 0.250 of pentane.Which substance is the solute?

Q4) Using your knowledge of osmosis or osmotic pressure explain the following statement: Drinking salt water actually dehydrates our tissues.

Q5) List the following solutions in order of increasing boiling point: 0.15m NaI, 0.10m K₃PO₄, 0.20 ethylene glycol

Q6) A 100.mL sample of water is taken from the Great Salt Lake, and the water is allowed to evaporate.The salts that remain (mostly NaCl)have a mass of 31.9 g Calculate the original concentration of NaCl, in g per liter, in each water sample.

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Page 14

Chapter 13: Chemical Kinetics

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Sample Questions

Q1) The reaction A + 2B \(\rarr\) products has been found to have the rate law, rate = k[A] [B]².While holding the concentration of A constant, the concentration of B is increased from x to 3x.Predict by what factor the rate of reaction increases.

A)3

B)6

C)9

D)27

E)30

Q2) For the reaction X₂ + Y + Z \(\rarr\) XY + XZ, it is found that the rate equation is rate = k [X₂][Y].Why does the concentration of Z have no effect on the rate?

A)The concentration of Z is very small and the others are very large.

B)Z must react in a step after the rate determining step.

C)Z is an intermediate.

D)The fraction of molecules of Z that have very high energies is zero.

E)The activation energy for Z to react is very high.

Q3) Given the rate law for a reaction, rate = k[A][B]², where rate is measured in units of M s^-1, what are the units for the rate constant k?

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Chapter 14: Chemical Equilibrium

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Q1) A reaction with an equilibrium constant K_c = 1.5 x 10^-25 would consist of which of the following at equilibrium:

A)approximately equal reactants and products

B)some reactants and products with reactants slightly favored

C)some reactants and products with products slightly favored

D)essentially all reactants

E)essentially all products

Q2) A reaction with an equilibrium constant K_c = 1.5 x 10²¹ would consist of which of the following at equilibrium:

A)approximately equal reactants and products

B)some reactants and products with reactants slightly favored

C)some reactants and products with products slightly favored

D)essentially all reactants

E)essentially all products

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.

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Chapter 15: Acids and Bases

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Sample Questions

Q1) What mass of sodium cyanide must be added to 250.mL of water in order to obtain a solution having a pH of 10.50? [K_a(HCN)= 4.9 × 10^-10]

A)240 g

B)0.032 g

C)0.059 g

D)0.94 g

E)0.24 g

Q2) Calculate the hydrogen ion concentration in a solution of fruit juice having a pH of 4.25.

A)1.0 × 10^-14 M

B)5.6 × 10^-5 M

C)4.0 × 10^-25 M

D)2.5 × 10^-4 M

E)5.6 × 10^-4 M

Q3) Write the formula for the conjugate acid of H₂PO₄^-.

Q4) When 2.0 × 10^-2 mole of nicotinic acid (a monoprotic acid)is dissolved in 350.mL of water, the pH is 3.05.What is the K_a of nicotinic acid?

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Page 17

Chapter 16: Acid-Base Equilibria and Solubility Equilibria

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Sample Questions

Q1) Calculate the pH of a solution that is 0.15 M CH₃COOH and 0.75 M CH₃COONa.

Q2) Calculate the molar solubility of BaCO₃ in a 0.10 M solution of Na₂CO₃(aq).(K_sp (BaCO₃)= 8.1 x 10^-9)

A)8.1 x 10^-9 M

B)9.0 x 10^-5 M

C)8.1 x 10^-8 M

D)2.8 x 10^-4 M

E)0.10 M

Q3) The K_sp for silver(I)phosphate is 1.8 × 10^-18.Calculate the molar solubility of silver(I)phosphate.

A)1.6 × 10^-5 M

B)2.1 × 10^-5 M

C)3.7 × 10^-5 M

D)7.2 × 10^-1 M

E)1.8 × 10^-1 M

Q4) Write a net ionic equation for the reaction that occurs when a small amount of nitric acid is added to a NaNO₂/HNO₂ buffer.

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Chapter 17: Entropy Free Energy and Equilibrium

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Q1) Under what conditions (always, never, high temperature only, low temperature only)is the reaction O₂(g)\(\rarr\) 2O(g)expected to be spontaneous?

Q2) Under what conditions (always, never, high temperature only, low temperature only)is the process: H₂O(l)\(\rarr\)H₂O(s)expected to be spontaneous?

Q3) The entropy of a perfectly ordered crystalline substance at 0 K is 0 J/mol.

A)True

B)False

Q4) Choose the substance with the higher entropy per mole at a given temperature: Br₂(l)or Br₂(g).

Q5) Which of the following is consistent with a reaction that proceeds spontaneously in the reverse direction (assume all variables are in terms of the forward direction only)?

A)(\(\Delta\)G > 0, Q < K)

B)(\(\Delta\)G° = 0, Q = K)

C)(\(\Delta\)G < 0, Q > K)

D)(\(\Delta\)G° > 0, Q = K)

E)(\(\Delta\)G > 0, Q > K)

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Chapter 18: Electrochemistry

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Sample Questions

Q1) The measured voltage of a cell in which the following reaction occurs is 0.96 V: H₂(g, 1.0 atm)+ 2Ag⁺(aq, 1.0 M)\(\rarr\) 2H⁺(aq, pH = ?)+ 2Ag(s)

Calculate the pH of the H⁺(aq)solution.

A)1.4

B)2.7

C)5.4

D)7.1

E)14.9

Q2) Complete and balance the following redox reaction under acidic conditions: ClO₂^-(aq)\(\rarr\) ClO₂(g)+ Cl^-(aq)

Q3) Consider the following reaction: 2Fe²⁺(aq)+ Cu²⁺ \(\rarr\) 2Fe³⁺(aq)+ Cu. When the reaction comes to equilibrium, what is the cell voltage?

A)0.43 V

B)1.11 V

C)0.78 V

D)-0.43 V

E)0 V

Q4) Will H₂(g)form when Sn is placed in 1.0 M HCl?

Page 20

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Chapter 19: Nuclear Chemistry

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Sample Questions

Q1) Decay of lutetium-167 by electron capture yields

A)ytterbium-167.

B)lutetium-166.

C)thulium-163.

D)tantalum-171.

E)hafnium-167.

Q2) A typical radius of an atomic nucleus is about

A)100 µm

B)5000 mm

C)100 nm

D)5 × 10^-3 pm

E)500 pm

Q3) Decay of silicon-27 by positron emission yields

A)magnesium-23.

B)sulfur-31.

C)phosphorus-27.

D)silicon-26.

E)aluminum-27.

Q4) Strontium-90 has a half-life of 28.8 years.How much strontium-90 was present initially, if after 144 years 10.0 g remain?

21

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Chapter 20: Chemistry in the Atmosphere

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Q1) Acid rain is precipitation having a pH that is

A)above 7.0

B)below 5.5

C)below 7.0

D)above 8.6

E)above 10.0

Q2) Excluding water vapor, list the four most prevalent gases in our atmosphere from most abundant to least abundant.

Q3) Heat is radiated from Earth to space predominately in the form of:

A)infrared radiation.

B)radioactivity.

C)visible light.

D)solar flares.

E)UV radiation.

Q4) The ozone layer in the atmosphere protects animals and plants from A)infrared radiation.

B)acid rain.

C)visible sunlight.

D)radioactivity.

E)ultraviolet rays.

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Chapter 21: Metallurgy and the Chemistry of Metals

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Q1) Aluminum hydroxide, Al(OH)₃, is

A)an acid.

B)an amphoteric hydroxide.

C)a base.

D)an explosive hydroxide.

E)used to make amalgams.

Q2) Write a balanced chemical equation illustrating roasting.

Q3) According to the band theory, a band is

A)an elastic force that holds electrons close to atoms in an insulator.

B)the energy gap associated with semiconductors.

C)a large number of molecular orbitals that are close together in energy.

D)a large number of atoms in a crystal.

Q4) Which one of these metals would normally be obtained by chemical reduction?

A)aluminum

B)calcium

C)lithium

D)sodium

E)vanadium

Q5) Write two balanced chemical equations for the burning of magnesium in air.

Q6) Write the chemical formula of magnetite.

Page 23

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Chapter 22: Nonmetallic Elements and Their Compounds

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Q1) When liquid phosphorus trichloride reacts with water, the products are

A)PCl₅ and H₃PO₄.

B)H₃PO₃ and Cl₂.

C)H₃PO₄ and HCl.

D)H₃PO₄ and Cl₂.

E)H₃PO₃ and HCl.

Q2) Which of these aqueous solutions has the lowest pH?

A)0.100 M NaOH

B)0.100 M Na₂O

C)0.100 M Na₃N

D)all of these solutions have the same pH due to the leveling effect

E)these all are solutions of weak bases, so K_b values are needed in order to decide

Q3) Which of the following substances is a peroxide?

A)KO₂

B)Na₂O₂

C)CO₂

D)CaO

E)Al₂O₃

Q4) Write the chemical formula of the oxide ion.

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Chapter 23: Transition Metal Chemistry and Coordination Compounds

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Q1) The electron configuration of a Co³⁺ion is A)[Ar]3d⁶.

B)[Ar]4s¹3d⁵.

C)[Ar] 4s²3d⁴.

D)[Ar]3d⁵.

E)[Ar]3d⁴.

Q2) Which of these ligands produces the weakest crystal field?

A)CN^-

B)I^-

C)OH^-

D)H₂O

E)NH₃

Q3) How many geometric isomers are possible for the complex [CrF₃Br₃]^3-? Draw these isomers.

Q4) Write the chemical formula of diamminedichloroplatinum(II).

Q5) Predict the number of unpaired electrons in the [Cr(en)₃]²⁺ ion.

Q7) Name the complex ion Page 25

Q6) What is the oxidation number of Co in the complex Na[Co(H₂O)₂Cl₄]?

[Co(H₂O)₄Cl₂]⁺.

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Chapter 24: Organic Chemistry

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Q1) Amines are

A)organic bases that react with water to produce ammonia.

B)organic acids that react with water to produce ammonia.

C)organic bases that react with acids to form ammonium salts.

D)organic acids that react with bases to form ammonium salts.

E)None of the above.

Q2) Which functional group, when present in a compound that is allowed to stand in air, poses a danger of slowly yielding explosive peroxides?

A)ether

B)alcohol

C)carboxylic acid

D)ketone

E)unsaturated hydrocarbon

Q3) Alkynes have the general formula

A)C_nH_2n-4

B)C_nH_2n-2

C)C_nH_2n

D)C_nH_2n+2

E)C_nH_2n+4

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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 choice contains all three molecular units found in nucleotides?

A)phosphate, sugar, amino acid

B)amino acid, nitrogen-containing base, sugar

C)carboxylic acid, sugar, protein

D)phosphate, nitrogen-containing base, sugar

E)sugar, amino acid, protein

Q3) Which of the following polymers is formed by a condensation process?

A)PVC

B)nylon

C)Teflon

D)Plexiglas

E)neoprene

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