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D∆¥ 1 PH¥$!¢∆l & ¢h€mi¢∆l Pr!n¢!pl€$ > Chemistry by Raymond Chang (based on the summary of end chapters only) > Flashcards

Chemistry by Raymond Chang (based on the summary of end chapters only) Flashcards

1
Q

The study of chemistry involves three basic steps. Which of the following is NOT one of those steps?

A) Observation
B) Interpretation
C) Analysis
D) Representation

A

C) Analysis

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2
Q

What does the “observation” step in chemistry refer to?

A) Understanding molecular structures
B) Measurements in the macroscopic world
C) The use of symbols for communication
D) Developing scientific theories

A

B) Measurements in the macroscopic world

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3
Q

What does the “representation” step in chemistry involve?

A) Creating visual diagrams of atoms and molecules
B) The use of shorthand notation symbols and equations for communication
C) Testing hypotheses
D) Gathering information through experiments

A

B) The use of shorthand notation symbols and equations for communication

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4
Q

The scientific method is a systematic approach to research that begins with which of the following?

A) Testing hypotheses
B) Developing theories
C) Gathering information through observation and measurements
D) Using scientific notation

A

C) Gathering information through observation and measurements

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5
Q

Which substances in chemistry are formed by the chemical combination of atoms in fixed proportions?

A) Elements
B) Mixtures
C) Compounds
D) Solutions

A

C) Compounds

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6
Q

All substances can exist in which three states?

A) Liquid, gas, and plasma
B) Solid, liquid, and gas
C) Solid, plasma, and gas
D) Liquid, gas, and plasma

A

B) Solid, liquid, and gas

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7
Q

What can cause the interconversion between different states of matter?

A) Change in density
B) Change in temperature
C) Change in molecular structure
D) Change in weight

A

B) Change in temperature

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8
Q

Which units are used to express physical quantities in chemistry?

A) Metric units
B) Imperial units
C) SI units
D) Standard units

A

C) SI units

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9
Q

Scientific notation helps in handling which types of quantities?

A) Only very large quantities
B) Only very small quantities
C) Both very large and very small quantities
D) Only exact quantities

A

C) Both very large and very small quantities

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10
Q

In scientific notation, what is the range for the value of N in the form N×10^n ?

A) 0 to 1
B) 1 to 10
C) 10 to 100
D) Any positive number

A

B) 1 to 10

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11
Q

Which of the following statements best describes the scientific method?

A) A fixed set of steps followed exactly the same way each time
B) A systematic approach that begins with hypotheses
C) A method of testing and refining laws and theories
D) A systematic approach to research that starts with gathering information through observation and measurements

A

A) A fixed set of steps followed exactly the same way each time

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12
Q

Which of the following properties of a substance can be observed without changing its identity?

A) Chemical properties
B) Physical properties
C) Reactive properties
D) Ionic properties

A

B) Physical properties

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13
Q

Which of the following best defines chemical properties?

A) Properties that can be observed without changing the substance
B) Properties that describe the color and texture of a substance
C) Properties that, when demonstrated, change the identity of the substance
D) Properties that are unique to gases

A

C) Properties that, when demonstrated, change the identity of the substance

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14
Q

What type of mixture can be separated into pure components by physical means?

A) Homogeneous mixtures only
B) Heterogeneous mixtures only
C) Both homogeneous and heterogeneous mixtures
D) Only solutions

A

C) Both homogeneous and heterogeneous mixtures

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15
Q

What are the simplest substances in chemistry called?

A) Mixtures
B) Compounds
C) Elements
D) Solutions

A

C) Elements

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16
Q

Which of the following describes the combination of atoms of different elements in fixed proportions?

A) Mixture
B) Element
C) Solution
D) Compound

A

D) Compound

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17
Q

Which unit system is commonly used to express physical quantities in all sciences, including chemistry?

A) Metric system
B) SI units
C) Customary units
D) British Imperial system

A

B) SI units

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18
Q

Which of the following is NOT a step in the study of chemistry as described in the text?

A) Observation
B) Representation
C) Hypothesis
D) Interpretation

A

C) Hypothesis

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19
Q

What is required to separate the components of a mixture?

A) Chemical reactions
B) Physical means
C) Electrical processes
D) Magnetic processes

A

B) Physical means

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20
Q

According to Dalton’s atomic theory, what is true about atoms of the same element?

A) They are always positively charged.
B) They are identical.
C) They cannot combine with other elements.
D) They are different from each other.

A

B) They are identical.

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21
Q

What does the law of conservation of mass state?

A) Atoms can be created and destroyed in chemical reactions.
B) Atoms are only combined in certain ratios.
C) Atoms are neither created nor destroyed in chemical reactions.
D) Atoms are always in a gaseous state.

A

C) Atoms are neither created nor destroyed in chemical reactions.

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22
Q

Which law states that elements in a compound are always combined in the same proportions by mass?

A) Law of definite proportions
B) Law of conservation of mass
C) Law of multiple proportions
D) Law of atomic theory

A

A) Law of definite proportions

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23
Q

When two elements form more than one compound, what law states that the mass ratios of one element to a fixed mass of the other are in small whole numbers?

A) Law of definite proportions
B) Law of multiple proportions
C) Dalton’s atomic theory
D) Law of conservation of mass

A

B) Law of multiple proportions

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24
Q

Which part of an atom contains protons and neutrons?

A) The outer shell
B) The nucleus
C) The electron cloud
D) The atomic shell

A

B) The nucleus

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25
Q

What is the relative charge of protons, neutrons, and electrons?

A) Protons are positive, neutrons are neutral, electrons are negative
B) Protons are negative, neutrons are neutral, electrons are positive
C) Protons are neutral, neutrons are positive, electrons are negative
D) Protons are positive, neutrons are negative, electrons are neutral

A

A) Protons are positive, neutrons are neutral, electrons are negative

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26
Q

What determines the atomic number of an element?

A) The total number of protons and neutrons
B) The number of neutrons in the nucleus
C) The number of protons in the nucleus
D) The mass of the element

A

C) The number of protons in the nucleus

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27
Q

What is an isotope?

A) An atom with more protons than neutrons
B) An atom with the same number of protons but a different number of neutrons
C) An atom with a different number of protons and neutrons
D) An atom with fewer protons than neutrons

A

B) An atom with the same number of protons but a different number of neutrons

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28
Q

What is the purpose of a chemical formula?

A) To show only the elements present in a compound
B) To show the type and number of atoms in the smallest unit of a compound
C) To describe the shape of a molecule
D) To display the physical properties of a substance

A

B) To show the type and number of atoms in the smallest unit of a compound

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29
Q

What does a molecular formula convey?

A) Only the elements present in a molecule
B) The specific number and type of atoms in each molecule of a compound
C) The energy levels of each atom
D) The physical state of each element

A

B) The specific number and type of atoms in each molecule of a compound

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30
Q

What is an empirical formula?

A) A formula showing the exact number of atoms in a molecule
B) A formula showing the simplest ratio of the atoms in a compound
C) A formula that depicts the physical state of a compound
D) A formula used only for ionic compounds

A

B) A formula showing the simplest ratio of the atoms in a compound

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31
Q

Which type of compound consists of discrete, individual molecules?

A) Ionic compounds
B) Molecular compounds
C) Inorganic compounds
D) Hydrocarbon compounds

A

B) Molecular compounds

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32
Q

What type of compounds are made of cations and anions?

A) Organic compounds
B) Molecular compounds
C) Ionic compounds
D) Hydrocarbon compounds

A

C) Ionic compounds

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33
Q

What elements do organic compounds primarily contain?

A) Carbon and oxygen
B) Carbon and hydrogen
C) Hydrogen and nitrogen
D) Oxygen and nitrogen

A

B) Carbon and hydrogen

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34
Q

Which is the simplest type of organic compound?

A) Hydrocarbon
B) Salt
C) Acid
D) Base

A

A) Hydrocarbon

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35
Q

According to the law of multiple proportions, when two elements form multiple compounds, the mass ratios of one element to a fixed mass of the other are:

A) Random numbers
B) Small whole numbers
C) Decimals
D) Fractions

A

B) Small whole numbers

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36
Q

Which particles are found in the nucleus of an atom?

A) Protons and electrons
B) Electrons and neutrons
C) Protons and neutrons
D) Only protons

A

C) Protons and neutrons

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37
Q

Modern chemistry began with which theory?

A. Bohr’s model
B. Quantum theory
C. Dalton’s atomic theory
D. Rutherford’s atomic model

A

C. Dalton’s atomic theory

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38
Q

According to Dalton’s atomic theory, which of the following statements is not true?

A. Atoms are indivisible particles.
B. Atoms of different elements combine in whole-number ratios.
C. Atoms of the same element are identical in all respects.
D. Atoms are created in chemical reactions.

A

D. Atoms are created in chemical reactions.

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39
Q

The law that states that elements in a compound are always combined in the same proportions by mass is called:

A. Law of Conservation of Mass
B. Law of Definite Proportions
C. Law of Multiple Proportions
D. Avogadro’s Law

A

B. Law of Definite Proportions

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40
Q

When two elements can form more than one type of compound, their masses combine in small whole-number ratios. This is known as:

A. Law of Conservation of Energy
B. Law of Constant Composition
C. Law of Multiple Proportions
D. Law of Definite Proportions

A

C. Law of Multiple Proportions

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41
Q

An atom consists of a dense central nucleus surrounded by which type of particles?

A. Protons only
B. Neutrons only
C. Electrons only
D. Protons, neutrons, and electrons

A

D. Protons, neutrons, and electrons

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42
Q

Which particles in the atom are positively charged?

A. Neutrons
B. Electrons
C. Protons
D. Nucleus

A

C. Protons

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43
Q

The atomic number of an element is determined by the number of:

A. Neutrons in the nucleus
B. Protons in the nucleus
C. Electrons in the outer shell
D. Protons and neutrons combined

A

B. Protons in the nucleus

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44
Q

Isotopes are atoms of the same element that have:

A. Different numbers of electrons
B. Different numbers of protons
C. Different numbers of neutrons
D. Different atomic numbers

A

C. Different numbers of neutrons

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45
Q

Which of the following represents the simplest ratio of atoms in a molecule?

A. Structural formula
B. Empirical formula
C. Molecular formula
D. Ionic formula

A

B. Empirical formula

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46
Q

Chemical compounds can be classified as either:

A. Molecular or Ionic
B. Covalent or Metallic
C. Anions or Cations
D. Organic or Inorganic

A

A. Molecular or Ionic

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47
Q

What type of compounds contain carbon and elements like hydrogen, oxygen, and nitrogen?

A. Inorganic compounds
B. Metallic compounds
C. Organic compounds
D. Ionic compounds

A

C. Organic compounds

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48
Q

Hydrocarbons are the simplest type of which compound?

A. Inorganic
B. Molecular
C. Organic
D. Ionic

A

C. Organic

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49
Q

The formula of a compound can be deduced from:

A. The molecular weight of the compound
B. The atomic weights of its elements
C. A set of simple rules based on compound names
D. The color of the compound

A

C. A set of simple rules based on compound names

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50
Q

Atomic masses are measured in:

A. Molar units
B. Atomic mass units (amu)
C. Molecular units
D. Mass per volume

A

B. Atomic mass units (amu)

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51
Q

The atomic mass of an element is the average of:

A. Mass of a single atom
B. Mass of its isotopes
C. Number of protons and neutrons
D. Mass of the nucleus only

A

B. Mass of its isotopes

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52
Q

The molar mass of an element or compound is equal to:

A. The atomic mass in amu multiplied by Avogadro’s number
B. The atomic mass in grams per mole
C. The number of atoms per molecule
D. Half the atomic number

A

B. The atomic mass in grams per mole

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53
Q

Percent composition by mass of a compound refers to:

A. The percentage of each isotope in an element
B. The mass percentage of each element in the compound
C. The total mass of the compound in grams
D. The average molecular mass

A

B. The mass percentage of each element in the compound

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54
Q

If the percent composition by mass of a compound is known, what can be deduced?

A. The isotopic abundance of elements
B. The empirical and molecular formulas
C. The atomic number of elements
D. The color of the compound

A

B. The empirical and molecular formulas

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55
Q

Chemical reactions are represented by:

A. Physical equations
B. Chemical equations
C. Mathematical equations
D. Nuclear equations

A

B. Chemical equations

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56
Q

Chemical equations must be balanced to obey which law?

A. Law of Multiple Proportions
B. Law of Conservation of Mass
C. Law of Definite Proportions
D. Law of Atomic Theory

A

B. Law of Conservation of Mass

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57
Q

In a chemical equation, the substances on the left side of the arrow are called:

A. Products
B. Molecules
C. Reactants
D. Compounds

A

C. Reactants

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58
Q

Stoichiometry is primarily concerned with:

A. The mass of compounds only
B. The study of the ratio of products and reactants in chemical reactions
C. The color changes in a reaction
D. The energy changes during a reaction

A

B. The study of the ratio of products and reactants in chemical reactions

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59
Q

What is the limiting reagent in a chemical reaction?

A. The reactant that is present in the smallest stoichiometric amount
B. The product that forms first in a reaction
C. The reactant that is present in excess
D. The catalyst used in the reaction

A

A. The reactant that is present in the smallest stoichiometric amount

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60
Q

The amount of product actually obtained in a chemical reaction is called the:

A. Theoretical yield
B. Actual yield
C. Percent composition
D. Empirical yield

A

B. Actual yield

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61
Q

Percent yield is calculated by:

A. Dividing theoretical yield by actual yield
B. Dividing actual yield by theoretical yield and multiplying by 100%
C. Dividing actual yield by molar mass
D. Multiplying the limiting reagent by 100%

A

B. Dividing actual yield by theoretical yield and multiplying by 100%

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62
Q

Which of the following statements is true regarding the conductivity of aqueous solutions?

A) Solutions only conduct electricity if the solutes are nonelectrolytes.
B) All aqueous solutions are electrically conductive.
C) Solutions are electrically conductive if the solutes are electrolytes.
D) Solutions are electrically conductive regardless of the type of solute.

A

C) Solutions are electrically conductive if the solutes are electrolytes.

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63
Q

Which of the following is NOT a category of chemical reactions in aqueous solutions?

A) Precipitation reactions
B) Combustion reactions
C) Acid-base reactions
D) Oxidation-reduction reactions

A

B) Combustion reactions

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64
Q

In aqueous solutions, what is the expected result when an ionic compound is insoluble?

A) It will dissolve completely.
B) It will form a precipitate.
C) It will produce gases.
D) It will conduct electricity.

A

B) It will form a precipitate.

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65
Q

According to Arrhenius, acids and bases ionize in water to produce which ions, respectively?

A) H⁺ ions and OH⁻ ions
B) H2 ions and O⁻ ions
C) OH⁻ ions and H⁺ ions
D) OH2 ions and H ions

A

A) H⁺ ions and OH⁻ ions

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66
Q

Brønsted acids and bases are defined by which of the following characteristics?

A) Brønsted acids donate electrons, and Brønsted bases accept electrons.
B) Brønsted acids accept protons, and Brønsted bases donate protons.
C) Brønsted acids donate protons, and Brønsted bases accept protons.
D) Brønsted acids donate electrons, and Brønsted bases accept electrons.

A

C) Brønsted acids donate protons, and Brønsted bases accept protons.

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67
Q

What is the reaction called when an acid and a base combine?

A) Decomposition
B) Oxidation
C) Neutralization
D) Precipitation

A

C) Neutralization

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68
Q

In redox reactions, oxidation is characterized by which of the following?

A) Gain of electrons
B) Loss of electrons
C) Formation of ions
D) Neutralization of charges

A

B) Loss of electrons

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69
Q

Which of the following best describes an oxidation number?

A) The total number of electrons in an atom
B) A measure of an atom’s ability to lose protons
C) A tool to track charge distribution in a compound
D) The amount of energy released in a reaction

A

C) A tool to track charge distribution in a compound

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70
Q

Which reaction types can redox reactions be subclassified into?

A) Combination, decomposition, combustion, displacement, disproportionation
B) Precipitation, acid-base, neutralization, displacement
C) Sublimation, combustion, crystallization, evaporation
D) Dissolution, neutralization, decomposition, precipitation

A

A) Combination, decomposition, combustion, displacement, disproportionation

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71
Q

Molarity is defined as which of the following?

A) Number of grams of solute in 1 liter of solution
B) Number of moles of solute in 1 liter of solvent
C) Number of moles of solute in 1 liter of solution
D) Number of atoms of solute in 1 liter of solution

A

C) Number of moles of solute in 1 liter of solution

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72
Q

Dilution involves which of the following changes in a solution?

A) Increase in concentration
B) Decrease in concentration without changing the number of moles of solute
C) Increase in both concentration and volume
D) Decrease in both concentration and volume

A

B) Decrease in concentration without changing the number of moles of solute

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73
Q

Gravimetric analysis is commonly used for which of the following purposes?

A) Determining the identity or concentration of a solution by measuring its mass
B) Determining the boiling point of a solution
C) Measuring the color change in a solution
D) Calculating the pH of a solution

A

A) Determining the identity or concentration of a solution by measuring its mass

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74
Q

In acid-base titration, the equivalence point is defined as:

A) The point where the solution changes color
B) The point at which the reaction is complete
C) The point where the acid is fully dissolved
D) The point where the pH is exactly 7

A

B) The point at which the reaction is complete

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75
Q

Redox titrations are similar to acid-base titrations in that:

A) They both involve measuring temperature changes
B) They both reach an equivalence point where the reaction is complete
C) They both require gravimetric analysis
D) They both end when the solution reaches pH 7

A

B) They both reach an equivalence point where the reaction is complete

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76
Q

At 25°C and 1 atm, which of the following statements is true about the state of elements and molecular compounds?

A) All elements exist as gases.
B) Ionic compounds are gases under atmospheric conditions.
C) Many elements and molecular compounds exist as gases, but ionic compounds are usually solids.
D) No elements exist as gases under these conditions.

A

C) Many elements and molecular compounds exist as gases, but ionic compounds are usually solids.

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77
Q

Why do gases exert pressure on surfaces?

A) Because gas molecules are stationary.
B) Because gas molecules move randomly and collide with surfaces.
C) Due to the high density of gas molecules.
D) Because gas molecules expand when heated.

A

B) Because gas molecules move randomly and collide with surfaces.

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78
Q

Which of the following units is equivalent to 1 atmosphere (atm) of pressure?

A) 100 mmHg
B) 500 torr
C) 760 mmHg or 760 torr
D) 1 mmHg

A

C) 760 mmHg or 760 torr

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79
Q

Boyle’s Law describes the relationship between which two properties of an ideal gas?

A) Temperature and pressure
B) Volume and temperature
C) Volume and pressure, with volume being inversely proportional to pressure
D) Volume and number of molecules

A

C) Volume and pressure, with volume being inversely proportional to pressure

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80
Q

Charles’s and Gay-Lussac’s Law describe the relationship between which two properties of an ideal gas?

A) Volume and pressure
B) Volume and temperature, with volume directly proportional to temperature
C) Temperature and pressure, with pressure directly proportional to temperature
D) Volume and amount of gas

A

B) Volume and temperature, with volume directly proportional to temperature

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81
Q

What is absolute zero?

A) The highest temperature a gas can reach
B) The lowest temperature theoretically attainable, at -273.15°C
C) The freezing point of water
D) The temperature at which gases condense

A

B) The lowest temperature theoretically attainable, at -273.15°C

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82
Q

Which of the following laws states that equal volumes of gases contain equal numbers of molecules under the same temperature and pressure?

A) Boyle’s Law
B) Charles’s Law
C) Avogadro’s Law
D) Dalton’s Law

A

C) Avogadro’s Law

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83
Q

The ideal gas equation, PV=nRT, is derived from the combination of which laws?

A) Boyle’s, Charles’s, and Avogadro’s Laws
B) Charles’s, Dalton’s, and Graham’s Laws
C) Boyle’s, Graham’s, and Dalton’s Laws
D) Avogadro’s and Graham’s Laws

A

A) Boyle’s, Charles’s, and Avogadro’s Laws

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84
Q

Dalton’s Law of Partial Pressures states that each gas in a mixture:

A) Has no impact on the total pressure
B) Exerts the same pressure as it would if it occupied the entire volume alone
C) Contributes more pressure if it has a higher molar mass
D) Has zero interaction with other gases in the mixture

A

B) Exerts the same pressure as it would if it occupied the entire volume alone

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85
Q

Which of the following statements is NOT an assumption of the kinetic molecular theory of gases?

A) Gas molecules are separated by distances far greater than their own dimensions.
B) Gas molecules have negligible volume and are in constant motion.
C) Gas molecules attract and repel one another strongly.
D) Gas molecules frequently collide with each other.

A

C) Gas molecules attract and repel one another strongly.

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86
Q

A Maxwell speed distribution curve is used to show:

A) How the pressure of a gas changes with temperature
B) The number of gas molecules moving at various speeds at a given temperature
C) How gas molecules lose energy over time
D) The density of a gas at different temperatures

A

B) The number of gas molecules moving at various speeds at a given temperature

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87
Q

In diffusion, gas molecules:

A) Move through a small opening under pressure
B) Mix gradually with each other
C) Are condensed into a liquid form
D) Remain stationary

A

B) Mix gradually with each other

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88
Q

The process of effusion involves:

A) Two gases mixing gradually
B) Gas molecules moving through a small opening under pressure
C) Liquefaction of gases under high pressure
D) Gas molecules reacting chemically

A

B) Gas molecules moving through a small opening under pressure

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89
Q

The van der Waals equation is a modification of the ideal gas equation that accounts for:

A) High temperatures only
B) High pressures and the fact that real gases occupy volume and exert forces on each other
C) The diffusion rate of gases
D) The kinetic energy of gas molecules

A

B) High pressures and the fact that real gases occupy volume and exert forces on each other

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90
Q

The law of conservation of energy states that:

A) Energy is only created and cannot be destroyed.
B) The total amount of energy in the universe is constant.
C) Energy cannot be converted from one form to another.
D) Energy can only exist in thermal and electrical forms.

A

B) The total amount of energy in the universe is constant.

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91
Q

A process that gives off heat to the surroundings is termed:

A) Endothermic
B) Exothermic
C) Isothermal
D) Adiabatic

A

B) Exothermic

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92
Q

Which of the following describes a state function?

A) It is defined by the path taken to reach a final state.
B) It depends on initial and final states only, not the path.
C) It changes based on the speed of the reaction.
D) It is independent of the properties of the system.

A

B) It depends on initial and final states only, not the path.

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93
Q

Which of the following are considered state functions?

A) Work and energy
B) Energy and pressure
C) Work and heat
D) Work and pressure

A

B) Energy and pressure

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94
Q

According to the first law of thermodynamics, energy:

A) Cannot be converted to another form.
B) Cannot be created or destroyed, only converted from one form to another.
C) Exists only in mechanical form.
D) Is a concept exclusive to chemistry.

A

B) Cannot be created or destroyed, only converted from one form to another.

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95
Q

The change in enthalpy (ΔH) for a constant-pressure process is equal to:

A) ΔE + PΔV
B) ΔE - PΔV
C) ΔE × P
D) ΔE/P

A

A) ΔE + PΔV

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96
Q

Enthalpy (H) is a state function that represents:

A) The change in internal energy of the system.
B) The heat of reaction at constant pressure.
C) The amount of heat at constant volume.
D) The kinetic energy of particles.

A

B) The heat of reaction at constant pressure.

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97
Q

Which types of calorimeters are used to measure heat changes in physical and chemical processes?

A) Constant-temperature and constant-pressure calorimeters
B) Constant-volume and constant-pressure calorimeters
C) Isothermal and adiabatic calorimeters
D) Differential and standard calorimeters

A

B) Constant-volume and constant-pressure calorimeters

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98
Q

Hess’s Law states that the overall enthalpy change in a reaction:

A) Depends on the path taken by the reaction.
B) Is the sum of the enthalpy changes of individual steps.
C) Cannot be measured directly.
D) Is equal to the difference in enthalpy between reactants and products only.

A

B) Is the sum of the enthalpy changes of individual steps.

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99
Q

The standard enthalpy of a reaction can be calculated from:

A) The sum of bond energies in reactants and products.
B) The difference in entropy of the reactants and products.
C) The standard enthalpies of formation of reactants and products.
D) The molar masses of reactants and products.

A

C) The standard enthalpies of formation of reactants and products.

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100
Q

Theheat of solution of an ionic compound in water is the sum of:

A) The lattice energy of the compound and the heat of dilution.
B) The bond dissociation energy and lattice energy.
C) The lattice energy of the compound and the heat of hydration.
D) The bond formation energy and the heat of hydration.

A

C) The lattice energy of the compound and the heat of hydration.

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101
Q

**The heat of dilution is defined as:

A) The heat required to dissolve a compound in water.
B) The heat absorbed or evolved when a solution is diluted.
C) The energy needed to break bonds in a compound.
D) The heat released during an exothermic reaction.

A

B) The heat absorbed or evolved when a solution is diluted.

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102
Q

The quantum theory, developed by Planck, explains the emission of radiation by heated solids by stating that radiant energy is:

A) Emitted in a continuous range.
B) Emitted in small, discrete amounts called quanta.
C) Emitted only at certain wavelengths.
D) Emitted only in the visible spectrum.

A

B) Emitted in small, discrete amounts called quanta.

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103
Q

Einstein used quantum theory to propose which of the following ideas?

A) Light behaves as a continuous wave.
B) Light behaves as a stream of particles called photons.
C) Electrons travel in fixed paths around the nucleus.
D) Energy levels in atoms are continuous.

A

B) Light behaves as a stream of particles called photons.

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104
Q

The line spectrum of hydrogen was explained by Bohr using quantum theory by:

A) Describing the electron’s path as a continuous orbit.
B) Quantizing the electron’s energy to specific values defined by a principal quantum number.
C) Assuming the electron has no defined position.
D) Using the concept of protons.

A

B) Quantizing the electron’s energy to specific values defined by a principal quantum number.

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105
Q

In the Bohr model of the atom, an electron emits a photon when:

A) It moves from a lower-energy state to a higher-energy state.
B) It moves from a higher-energy state to a lower-energy state.
C) It remains in its ground state.
D) It becomes neutralized.

A

B) It moves from a higher-energy state to a lower-energy state.

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106
Q

According to de Broglie’s theory, the wavelength of a moving particle is given by:

A) λ = mv
B) λ = E/h
C) λ = h/mv
D) λ = hE

A

C) λ = h/mv

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107
Q

The Schrödinger equation describes:

A) The velocity of large particles.
B) The motions and energies of submicroscopic particles, like electrons.
C) The energy levels in a nucleus.
D) The magnetic field of an atom.

A

B) The motions and energies of submicroscopic particles, like electrons.

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108
Q

Which of the following best describes an atomic orbital?

A) A specific path an electron follows around the nucleus.
B) A function defining the probability distribution of electron density in space.
C) A fixed point where electrons are located.
D) A region where electrons have zero probability of being found.

A

B) A function defining the probability distribution of electron density in space.

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109
Q

How many quantum numbers characterize each electron in an atom?

A) One
B) Two
C) Three
D) Four

A

D) Four

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110
Q

The single s orbital for each energy level is:

A) Spherical and centered on the nucleus.
B) Dumbbell-shaped and oriented at right angles.
C) Disk-shaped and located around the nucleus.
D) Tetrahedral and not centered on the nucleus.

A

A) Spherical and centered on the nucleus.

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111
Q

In a hydrogen atom, the energy of the electron is determined solely by:

A) Its angular momentum.
B) Its magnetic quantum number.
C) Its principal quantum number.
D) Its spin.

A

C) Its principal quantum number.

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112
Q

The Pauli Exclusion Principle states that:

A) Two electrons can have the same four quantum numbers.
B) No two electrons in the same atom can have the same four quantum numbers.
C) Electrons fill orbitals of lowest energy first.
D) The electron’s energy increases with distance from the nucleus.

A

B) No two electrons in the same atom can have the same four quantum numbers.

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113
Q

According to Hund’s Rule, the most stable arrangement of electrons in a subshell is:

A) With all electrons paired.
B) With the maximum number of parallel spins.
C) With each electron occupying a different subshell.
D) With the least number of unpaired electrons.

A

B) With the maximum number of parallel spins.

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114
Q

Which of the following describes diamagnetic atoms?

A) Atoms with one or more unpaired electrons.
B) Atoms in which all electrons are paired.
C) Atoms with only one electron in the outer shell.
D) Atoms with more protons than electrons.

A

B) Atoms in which all electrons are paired.

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115
Q

The Aufbau Principle provides a guideline for:

A) The arrangement of protons and neutrons in the nucleus.
B) The order in which electrons fill atomic orbitals based on energy.
C) The distribution of isotopes in nature.
D) The arrangement of ions in a crystal lattice.

A

B) The order in which electrons fill atomic orbitals based on energy.

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116
Q

In early versions of the periodic table, elements were arranged in increasing order of:

A) Atomic numbers
B) Atomic masses
C) Electronegativity
D) Melting points

A

B) Atomic masses

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117
Q

The modern periodic table arranges elements in order of their:

A) Atomic masses
B) Electronegativity
C) Atomic numbers
D) Atomic radii

A

C) Atomic numbers

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118
Q

The properties of elements are determined primarily by their:

A) Atomic radius
B) Atomic mass
C) Electron configuration
D) Electronegativity

A

C) Electron configuration

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119
Q

Periodic variations in the physical properties of elements are due to:

A) Differences in atomic structure
B) Variations in atomic mass
C) Changes in melting and boiling points
D) Differences in color and density

A

A) Differences in atomic structure

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120
Q

How does metallic character vary across the periodic table?

A) It increases across a period from left to right.
B) It decreases across a period from metals to metalloids to nonmetals.
C) It stays constant across all periods.
D) It is highest in noble gases.

A

B) It decreases across a period from metals to metalloids to nonmetals.

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121
Q

Atomic radius generally:

A) Increases from left to right across a period and decreases from top to bottom within a group.
B) Decreases from left to right across a period and increases from top to bottom within a group.
C) Remains constant across a period but increases from top to bottom.
D) Decreases both across a period and within a group.

A

B) Decreases from left to right across a period and increases from top to bottom within a group.

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122
Q

Which of the following statements about ionization energy is correct?

A) It measures the attraction of an atom for additional electrons.
B) It decreases as the attraction between the nucleus and electron increases.
C) It measures the tendency of an atom to resist the loss of an electron.
D) It is usually higher for metals than for nonmetals.

A

C) It measures the tendency of an atom to resist the loss of an electron.

123
Q

Elements with high electron affinities tend to:

A) Readily lose electrons.
B) Readily gain electrons.
C) Have low ionization energies.
D) Have large atomic radii.

A

B) Readily gain electrons.

124
Q

In the periodic table, noble gases are stable because:

A) They have high ionization energies.
B) Their outer ns and np subshells are completely filled.
C) They readily lose electrons.
D) They have incomplete valence shells.

A

B) Their outer ns and np subshells are completely filled.

125
Q

Which group of elements tends to lose electrons to become isoelectronic with the noble gases that precede them in the periodic table?

A) Noble gases
B) Nonmetals in Groups 5A, 6A, and 7A
C) Metals in Groups 1A, 2A, and 3A
D) Transition metals

A

C) Metals in Groups 1A, 2A, and 3A

126
Q

Nonmetals in Groups 5A, 6A, and 7A tend to:

A) Accept electrons to become isoelectronic with the noble gases that follow them.
B) Lose electrons to match the electron configuration of noble gases.
C) Remain neutral and not react.
D) Gain protons to reach a stable configuration.

A

A) Accept electrons to become isoelectronic with the noble gases that follow them.

127
Q

A Lewis dot symbol is primarily used to represent which of the following?

A) Atomic mass of an element
B) The number of valence electrons in an atom of a representative element
C) The arrangement of protons and neutrons in the nucleus
D) The ionization energy of an atom

A

B) The number of valence electrons in an atom of a representative element

128
Q

Elements likely to form ionic compounds generally have:

A) Low ionization energies or high electron affinities
B) High ionization energies and low electron affinities
C) High atomic masses
D) Low atomic radii

A

A) Low ionization energies or high electron affinities

129
Q

An ionic bond is formed by:

A) The sharing of electrons between two atoms
B) The electrostatic attraction between positive and negative ions
C) The interaction between two identical atoms
D) The repulsion between valence electrons

A

B) The electrostatic attraction between positive and negative ions

130
Q

What is lattice energy?

A) The measure of the bond length in a covalent bond
B) The measure of the stability of an ionic solid
C) The amount of energy needed to ionize an atom
D) The energy associated with electron affinity

A

B) The measure of the stability of an ionic solid

131
Q

In a covalent bond, two atoms share:

A) Protons to form a strong nuclear bond
B) A pair of electrons
C) Neutrons to create an isotope
D) Their atomic mass units

A

B) A pair of electrons

132
Q

The arrangement of bonding electrons and lone pairs in a molecule is represented by:

A) A Lewis structure
B) The periodic table
C) Atomic mass units
D) The Born-Haber cycle

A

A) A Lewis structure

133
Q

Electronegativity is best described as:

A) The tendency of an atom to donate electrons
B) A measure of an atom’s ability to attract electrons in a chemical bond
C) The size of an atom
D) The number of valence electrons in an atom

A

B) A measure of an atom’s ability to attract electrons in a chemical bond

134
Q

According to the octet rule, atoms generally form covalent bonds to:

A) Surround themselves with exactly six electrons
B) Surround themselves with eight electrons in their valence shell
C) Match the electron configuration of hydrogen
D) Maximize their atomic radius

A

B) Surround themselves with eight electrons in their valence shell

135
Q

Which of the following is true about resonance structures?

A) Resonance structures represent an element’s atomic mass.
B) Only one resonance structure can exist for a given molecule.
C) They represent a molecule or ion more accurately than any single Lewis structure does.
D) They are less stable than a single Lewis structure.

A

C) They represent a molecule or ion more accurately than any single Lewis structure does.

136
Q

The strength of a covalent bond can be measured in terms of:

A) Bond enthalpy
B) Electron affinity
C) Ionization energy
D) Atomic radius

A

A) Bond enthalpy

137
Q

The VSEPR model for predicting molecular geometry is based on which of the following assumptions?

A) Valence-shell electron pairs are stationary.
B) Valence-shell electron pairs repel one another and tend to stay as far apart as possible.
C) Only bonding electron pairs determine the molecular shape.
D) Valence-shell electrons attract each other.

A

B) Valence-shell electron pairs repel one another and tend to stay as far apart as possible.

138
Q

According to the VSEPR model, lone pairs of electrons:

A) Do not affect molecular geometry.
B) Repel each other less forcefully than bonding pairs.
C) Repel bonding pairs more forcefully, distorting bond angles.
D) Are only found in nonpolar molecules.

A

C) Repel bonding pairs more forcefully, distorting bond angles.

139
Q

The dipole moment of a molecule is a measure of:

A) The total number of electrons in the molecule.
B) The mass of the molecule.
C) The charge separation in a molecule due to differences in electronegativity.
D) The bond angles in the molecule.

A

C) The charge separation in a molecule due to differences in electronegativity.

140
Q

Which two quantum mechanical theories explain covalent bond formation?

A) Molecular orbital theory and resonance theory
B) Valence bond theory and molecular orbital theory
C) VSEPR theory and valence bond theory
D) Atomic theory and kinetic theory

A

B) Valence bond theory and molecular orbital theory

141
Q

In sp hybridization, the two hybrid orbitals lie:

A) At the corners of a tetrahedron
B) In a straight line
C) In a trigonal planar arrangement
D) Around the nucleus

A

B) In a straight line

142
Q

In sp³d hybridization, the five hybrid orbitals are directed:

A) Towards the corners of a square
B) Towards the corners of a trigonal bipyramid
C) Towards the corners of a hexagon
D) In a straight line

A

B) Towards the corners of a trigonal bipyramid

143
Q

An sp²-hybridized carbon atom can form:

A) Two sigma bonds and one pi bond
B) One sigma bond and one pi bond
C) Three pi bonds
D) A single bond only

A

A) Two sigma bonds and one pi bond

144
Q

Molecular orbital theory describes bonding in terms of:

A) The attraction between ions
B) The combination and rearrangement of atomic orbitals into molecular orbitals
C) Electrons being shared only between atoms of the same element
D) Covalent bonds being weaker than ionic bonds

A

B) The combination and rearrangement of atomic orbitals into molecular orbitals

145
Q

Which of the following statements is true about bonding and antibonding molecular orbitals?

A) Bonding molecular orbitals are higher in energy than individual atomic orbitals.
B) Antibonding molecular orbitals have regions of high electron density between the nuclei.
C) Bonding molecular orbitals increase electron density between the nuclei and are lower in energy than individual atomic orbitals.
D) Antibonding molecular orbitals stabilize the molecule.

A

C) Bonding molecular orbitals increase electron density between the nuclei and are lower in energy than individual atomic orbitals.

146
Q

The number of molecular orbitals in a molecule is:

A) Equal to the number of atoms in the molecule.
B) Equal to the number of atomic orbitals combined.
C) Always greater than the number of atomic orbitals combined.
D) Determined solely by the atomic mass.

A

B) Equal to the number of atomic orbitals combined.

147
Q

Molecules are considered stable if:

A) The number of electrons in bonding molecular orbitals is equal to that in antibonding molecular orbitals.
B) The number of electrons in bonding molecular orbitals is greater than in antibonding molecular orbitals.
C) The molecule has only antibonding electrons.
D) There are no electrons in molecular orbitals.

A

B) The number of electrons in bonding molecular orbitals is greater than in antibonding molecular orbitals.

148
Q

Delocalized molecular orbitals:

A) Are found only in polar molecules.
B) Allow electrons to be free to move around multiple nuclei or groups of atoms.
C) Are confined between two specific atoms.
D) Represent regions with no electron density.

A

B) Allow electrons to be free to move around multiple nuclei or groups of atoms.

148
Q

Which of the following is the primary difference between the condensed state and the gaseous state of a substance?

A) Temperature
B) Molecular shape
C) Distance separating the molecules
D) Color of the substance

A

C) Distance separating the molecules

149
Q

What are intermolecular forces?

A) Attractive forces within a single molecule
B) Attractive forces between molecules or between molecules and ions
C) Forces that cause a molecule to emit light
D) Forces that increase as temperature increases

A

B) Attractive forces between molecules or between molecules and ions

150
Q

Which forces attract polar molecules with dipole moments to other polar molecules or ions?

A) Dispersion forces
B) Dipole-dipole and ion-dipole forces
C) Covalent bonds
D) Ionic bonds

A

B) Dipole-dipole and ion-dipole forces

151
Q

Dispersion forces arise due to:

A) Permanent dipoles in nonpolar molecules
B) Temporary dipoles in ordinarily nonpolar molecules
C) Attraction between ions and molecules
D) Hydrogen bonding in water

A

B) Temporary dipoles in ordinarily nonpolar molecules

152
Q

Hydrogen bonding is a strong type of:

A) Ionic bond
B) Dipole-dipole interaction involving a hydrogen atom and an electronegative atom (O, N, or F)
C) Covalent bond between hydrogen atoms
D) Metallic bond

A

B) Dipole-dipole interaction involving a hydrogen atom and an electronegative atom (O, N, or F)

153
Q

Surface tension is:

A) A property that minimizes the volume of a liquid
B) The energy required to expand a liquid’s surface area
C) The force that reduces the temperature of a liquid
D) A property that decreases as intermolecular forces increase

A

B) The energy required to expand a liquid’s surface area

154
Q

Viscosity, a measure of a liquid’s resistance to flow, generally:

A) Increases with increasing temperature
B) Decreases with increasing temperature
C) Remains constant with temperature change
D) Is unaffected by molecular interactions

A

B) Decreases with increasing temperature

155
Q

The open structure of water in its solid state (ice) is due to:

A) Dispersion forces between molecules
B) Covalent bonding
C) Hydrogen bonding forming a three-dimensional network
D) Ionic interactions

A

C) Hydrogen bonding forming a three-dimensional network

156
Q

Water plays a critical role in Earth’s climate due to:

A) Its high specific heat and hydrogen bonding
B) Its ability to absorb carbon dioxide
C) Its high viscosity
D) Its metallic bonding

A

A) Its high specific heat and hydrogen bonding

157
Q

All solids are classified as either:

A) Ionic or covalent
B) Crystalline or amorphous
C) Metallic or nonmetallic
D) Magnetic or nonmagnetic

A

B) Crystalline or amorphous

158
Q

The unit cell is:

A) A representation of the smallest repeating unit in a crystalline solid
B) The largest component of an amorphous solid
C) Only found in gases
D) The unit of measurement for liquid density

A

A) A representation of the smallest repeating unit in a crystalline solid

159
Q

The four types of crystals are:

A) Ionic, metallic, covalent, and molecular
B) Amorphous, liquid, metallic, and ionic
C) Crystalline, amorphous, molecular, and ionic
D) Gaseous, liquid, solid, and plasma

A

A) Ionic, metallic, covalent, and molecular

160
Q

In a closed container, a liquid establishes a dynamic equilibrium between:

A) Freezing and melting
B) Evaporation and condensation
C) Sublimation and deposition
D) Dissolution and crystallization

A

B) Evaporation and condensation

161
Q

The molar heat of vaporization is:

A) The energy required to melt one mole of a solid
B) The energy required to vaporize one mole of a liquid
C) The energy required to heat one mole of water
D) The energy released when a gas condenses

A

B) The energy required to vaporize one mole of a liquid

162
Q

For every substance, the temperature at which the gas phase cannot be liquefied, regardless of pressure, is called:

A) Boiling point
B) Freezing point
C) Critical temperature
D) Vapor pressure

A

C) Critical temperature

163
Q

A phase diagram shows:

A) The melting points of elements
B) The equilibrium phases of a single substance and the boundaries between them
C) The magnetic properties of a material
D) The specific heat of different phases of a substance

A

B) The equilibrium phases of a single substance and the boundaries between them

164
Q

A solution is best described as:

A) A homogeneous mixture of two or more substances
B) A heterogeneous mixture of two or more substances
C) A pure substance in the liquid state
D) A mixture where the solute and solvent are immiscible

A

A) A homogeneous mixture of two or more substances

165
Q

The ease of dissolving a solute in a solvent is primarily governed by:

A) Temperature
B) Intermolecular forces and the resulting disorder
C) The size of the solute particles
D) The amount of solvent present

A

B) Intermolecular forces and the resulting disorder

166
Q

Which of the following is NOT a common way to express the concentration of a solution?

A) Percent by mass
B) Molarity
C) Density
D) Mole fraction

A

C) Density

167
Q

Increasing temperature generally:

A) Increases the solubility of gases in water
B) Decreases the solubility of solids in water
C) Increases the solubility of solids and liquids, but decreases the solubility of gases in water
D) Does not affect solubility at all

A

C) Increases the solubility of solids and liquids, but decreases the solubility of gases in water

168
Q

According to Henry’s Law, the solubility of a gas in a liquid is:

A) Inversely proportional to the temperature
B) Directly proportional to the partial pressure of the gas over the solution
C) Independent of the gas’s partial pressure
D) Only applicable to nonpolar gases

A

B) Directly proportional to the partial pressure of the gas over the solution

169
Q

Raoult’s law states that the partial pressure of a substance A over a solution is equal to:

A) The vapor pressure of A in the solution
B) The mole fraction of A times the vapor pressure of pure A
C) The molarity of A times the volume of the solution
D) The density of A divided by the mole fraction

A

B) The mole fraction of A times the vapor pressure of pure A

170
Q

Which of the following is NOT a colligative property?

A) Boiling-point elevation
B) Freezing-point depression
C) Osmotic pressure
D) Solubility

A

D) Solubility

171
Q

In electrolyte solutions, the van’t Hoff factor provides:

A) The pH of the solution
B) A measure of the extent of dissociation of electrolytes in solution
C) The density of the solution
D) The boiling point of the solution

A

B) A measure of the extent of dissociation of electrolytes in solution

172
Q

A colloid is defined as:

A) A homogeneous solution with particles smaller than 1 pm
B) A dispersion of particles between 1 × 10³ pm and 1 × 10⁶ pm in another substance
C) A mixture with particles that settle over time
D) A solution that does not exhibit the Tyndall effect

A

B) A dispersion of particles between 1 × 10³ pm and 1 × 10⁶ pm in another substance

173
Q

The Tyndall effect is used to distinguish:

A) Solutions from colloids
B) Solvents from solutes
C) Gases from liquids
D) Electrolytes from nonelectrolytes

A

A) Solutions from colloids

174
Q

The rate of a chemical reaction is defined as:

A) The time taken for a reaction to complete
B) The change in concentration of reactants or products over time
C) The amount of energy released during the reaction
D) The total mass of reactants used

A

B) The change in concentration of reactants or products over time

175
Q

The rate constant (k) of a reaction changes with:

A) The concentration of reactants
B) The presence of a catalyst
C) Temperature only
D) The pressure of the system

A

C) Temperature only

176
Q

Reaction order refers to:

A) The sequence in which reactants are mixed
B) The power to which the concentration of a reactant is raised in the rate law
C) The amount of product formed in a reaction
D) The number of steps in a reaction mechanism

A

B) The power to which the concentration of a reactant is raised in the rate law

177
Q

The half-life of a reaction is:

A) The time required for the concentration of a reactant to decrease by half
B) The energy needed to initiate a reaction
C) The time taken for a catalyst to be consumed
D) A measure of the overall reaction order

A

A) The time required for the concentration of a reactant to decrease by half

178
Q

According to collision theory, a reaction occurs when:

A) Molecules collide with sufficient energy to break bonds
B) The temperature is raised to its boiling point
C) The products are more stable than the reactants
D) Molecules align perfectly with each other

A

A) Molecules collide with sufficient energy to break bonds

179
Q

The complete series of elementary steps for a reaction is known as:

A) The rate law
B) The reaction mechanism
C) The activation energy
D) The reaction order

A

B) The reaction mechanism

180
Q

The rate-determining step in a reaction mechanism is:

A) The first step in a sequence of reactions
B) The slowest step in the reaction mechanism
C) The step with the highest concentration of products
D) The step where the most energy is released

A

B) The slowest step in the reaction mechanism

181
Q

A catalyst speeds up a reaction by:

A) Increasing the concentration of reactants
B) Lowering the activation energy (Ea) needed for the reaction
C) Raising the temperature of the system
D) Changing the reaction order

A

B) Lowering the activation energy (Ea) needed for the reaction

182
Q

In heterogeneous catalysis:

A) The catalyst and reactants are in the same phase
B) The catalyst is in a different phase from the reactants
C) The reaction takes place in a homogeneous mixture
D) The catalyst gets consumed during the reaction

A

B) The catalyst is in a different phase from the reactants

183
Q

Dynamic equilibria between phases are called:

A) Chemical equilibria
B) Physical equilibria
C) Heterogeneous equilibria
D) Homogeneous equilibria

A

B) Physical equilibria

184
Q

Chemical equilibrium in a reaction is reached when:

A) The forward reaction rate is greater than the reverse reaction rate
B) The rates of the forward and reverse reactions are equal, and concentrations of reactants and products remain constant
C) The concentrations of reactants are zero
D) Only products are present

A

B) The rates of the forward and reverse reactions are equal, and concentrations of reactants and products remain constant

185
Q

In a chemical equilibrium expression, pure solids and pure liquids:

A) Appear with a concentration of zero
B) Are included with their actual concentrations
C) Are considered constants and do not appear in the equilibrium expression
D) Are doubled in the expression

A

C) Are considered constants and do not appear in the equilibrium expression

186
Q

A homogeneous equilibrium is a reaction in which:

A) All reactants and products are in different phases
B) All reactants and products are in the same phase
C) Only gases are involved
D) Only solids are involved

A

B) All reactants and products are in the same phase

187
Q

The equilibrium constant for a reaction that is the sum of two or more reactions is:

A) The sum of the equilibrium constants of the individual reactions
B) The difference between the equilibrium constants of the individual reactions
C) The product of the equilibrium constants of the individual reactions
D) The reciprocal of the equilibrium constant for each reaction

A

C) The product of the equilibrium constants of the individual reactions

188
Q

The equilibrium constant for the reverse of a particular reaction is:

A) The negative of the equilibrium constant for the forward reaction
B) The same as the forward reaction
C) The reciprocal of the equilibrium constant for the forward reaction
D) The square of the equilibrium constant for the forward reaction

A

C) The reciprocal of the equilibrium constant for the forward reaction

189
Q

The equilibrium constant K is defined as:

A) The sum of the concentrations of products and reactants
B) The ratio of the rate constant for the forward reaction to the rate constant for the reverse reaction
C) The sum of rate constants of forward and reverse reactions
D) The difference between rate constants of forward and reverse reactions

A

B) The ratio of the rate constant for the forward reaction to the rate constant for the reverse reaction

190
Q

The reaction quotient Q is used to:

A) Determine the overall reaction rate
B) Predict the direction of the reaction to reach equilibrium
C) Measure the activation energy of the reaction
D) Replace the equilibrium constant

A

B) Predict the direction of the reaction to reach equilibrium

191
Q

Le Châtelier’s principle states that if an external stress is applied to a system at equilibrium:

A) The system will shift to eliminate all reactants
B) The system will shift to form only products
C) The system will adjust to partially offset the stress
D) The equilibrium constant will double

A

C) The system will adjust to partially offset the stress

192
Q

Which of the following changes will affect the value of the equilibrium constant K for a reaction?

A) A change in concentration of reactants or products
B) A change in temperature
C) The addition of a catalyst
D) An increase in volume

A

B) A change in temperature

193
Q

According to Brønsted-Lowry theory, an acid is defined as a substance that:

A) Accepts protons
B) Donates protons
C) Accepts electrons
D) Donates electrons

A

B) Donates protons

194
Q

The pH of an aqueous solution is defined as:

A) The square root of the hydrogen ion concentration
B) The negative logarithm of the hydrogen ion concentration (in mol/L)
C) The molarity of the solution
D) The ionization constant of the solution

A

B) The negative logarithm of the hydrogen ion concentration (in mol/L)

195
Q

At 25°C, a solution with a pH of 5 is:

A) Neutral
B) Basic
C) Acidic
D) Amphoteric

A

C) Acidic

196
Q

Which of the following is classified as a strong acid?

A) HClO₄
B) CH₃COOH
C) H₂CO₃
D) NH₄OH

A

A) HClO₄

197
Q

The acid ionization constant (Ka) is:

A) Inversely proportional to acid strength
B) Constant for all acids
C) Directly proportional to acid strength
D) Only applicable to bases

A

C) Directly proportional to acid strength

198
Q

Percent ionization is:

A) A measure of an acid’s color
B) Another measure of the strength of an acid
C) Always constant regardless of acid concentration
D) Used only for strong acids

A

B) Another measure of the strength of an acid

199
Q

The product of the ionization constant of an acid and the ionization constant of its conjugate base is equal to:

A) The ion-product constant of water
B) The pH of the solution
C) The molarity of the acid
D) The percent ionization

A

A) The ion-product constant of water

200
Q

The relative strengths of acids can be explained qualitatively by:

A) Their solubility in water
B) Their molecular structures
C) Their physical state
D) Their color

A

B) Their molecular structures

201
Q

In salt hydrolysis, the reaction of salt ions with water can:

A) Produce neutral solutions only
B) Produce acidic or basic solutions
C) Produce only acidic solutions
D) Produce only basic solutions

A

B) Produce acidic or basic solutions

202
Q

Highly charged metal ions like Al³⁺ and Fe³⁺ :

A) Yield neutral solutions
B) Hydrolyze to yield acidic solutions
C) Are always amphoteric
D) Are strong bases

A

B) Hydrolyze to yield acidic solutions

203
Q

Metal oxides can be classified as:

A) Basic, acidic, or amphoteric
B) Only basic
C) Only acidic
D) Only neutral

A

A) Basic, acidic, or amphoteric

204
Q

Lewis acids are defined as substances that:

A) Accept pairs of electrons
B) Donate pairs of electrons
C) Contain ionizable hydrogen atoms
D) Act only in nonpolar solvents

A

A) Accept pairs of electrons

205
Q

The common ion effect tends to:

A) Increase the ionization of a weak acid or weak base
B) Suppress the ionization of a weak acid or weak base
C) Increase the solubility of a salt
D) Have no effect on ionization

A

B) Suppress the ionization of a weak acid or weak base

206
Q

A buffer solution is a combination of:

A) A strong acid and its conjugate base
B) A weak acid and its weak conjugate base or a weak base and its weak conjugate acid
C) Two strong acids
D) A salt and a strong acid

A

B) A weak acid and its weak conjugate base or a weak base and its weak conjugate acid

207
Q

The pH at the equivalence point of a strong acid–strong base titration is:

A) Greater than 7
B) Less than 7
C) Exactly 7
D) Variable, depending on the acid concentration

A

C) Exactly 7

208
Q

Acid-base indicators are:

A) Strong acids that do not change color
B) Weak organic acids or bases that change color near the equivalence point in an acid-base neutralization reaction
C) Inert substances used to detect pH
D) Strong bases that turn colorless when neutralized

A

B) Weak organic acids or bases that change color near the equivalence point in an acid-base neutralization reaction

209
Q

The solubility product (𝐾𝑠𝑝) expresses:

A) The equilibrium between a gas and a liquid
B) The equilibrium between a solid and its ions in solution
C) The rate of dissolution of a substance
D) The pH of a buffer solution

A

B) The equilibrium between a solid and its ions in solution

210
Q

The presence of a common ion:

A) Increases the solubility of a slightly soluble salt
B) Decreases the solubility of a slightly soluble salt
C) Has no effect on solubility
D) Neutralizes the salt in solution

A

B) Decreases the solubility of a slightly soluble salt

211
Q

The solubility of slightly soluble salts containing basic anions increases when:

A) The hydrogen ion concentration decreases
B) The hydrogen ion concentration increases
C) The pH is exactly 7
D) The solution temperature decreases

A

B) The hydrogen ion concentration increases

212
Q

Complex ions are formed by:

A) The combination of a metal cation with a Lewis base
B) The combination of a nonmetal with a Lewis base
C) The dissolution of a salt in water
D) A reduction reaction

A

A) The combination of a metal cation with a Lewis base

213
Q

Qualitative analysis involves:

A) The identification of the amount of a substance in a solution
B) The identification of cations and anions in a solution
C) Measuring the pH of a solution
D) Determining the solubility product of a salt

A

B) The identification of cations and anions in a solution

214
Q

Earth’s atmosphere is primarily composed of:

A) Oxygen and carbon dioxide
B) Nitrogen and oxygen
C) Argon and nitrogen
D) Hydrogen and helium

A

B) Nitrogen and oxygen

215
Q

Which factors influence the chemical processes in the atmosphere?

A) Solar radiation, volcanic eruptions, and human activities
B) Only human activities
C) Volcanic eruptions and the Earth’s core temperature
D) Only volcanic eruptions

A

A) Solar radiation, volcanic eruptions, and human activities

216
Q

Auroras are caused by:

A) Reflections of sunlight on Earth’s atmosphere
B) Bombardment of molecules and atoms by solar particles in the outer atmosphere
C) The greenhouse effect
D) Volcanic eruptions

A

B) Bombardment of molecules and atoms by solar particles in the outer atmosphere

217
Q

The ozone layer in the stratosphere absorbs harmful:

A) Infrared radiation
B) Visible light
C) UV radiation in the 200-300 nm range
D) Microwaves

A

C) UV radiation in the 200-300 nm range

218
Q

Chlorofluorocarbons (CFCs) have contributed to:

A) The formation of auroras
B) The protection of the ozone layer
C) The depletion of the ozone layer
D) Increased oxygen levels in the atmosphere

A

C) The depletion of the ozone layer

219
Q

Volcanic eruptions can lead to:

A) Increased oxygen levels in the atmosphere
B) Air pollution, ozone depletion, and climate effects
C) A reduction in greenhouse gases
D) The formation of auroras

A

B) Air pollution, ozone depletion, and climate effects

220
Q

Carbon dioxide contributes to global warming because:

A) It absorbs ultraviolet radiation
B) It absorbs infrared radiation and traps heat from Earth
C) It reflects sunlight away from the Earth’s surface
D) It prevents the formation of ozone

A

B) It absorbs infrared radiation and traps heat from Earth

221
Q

Sulfur dioxide, and to a lesser extent nitrogen oxides, are responsible for:

A) The greenhouse effect
B) The formation of acid rain
C) The formation of ozone
D) Increased oxygen levels

A

B) The formation of acid rain

222
Q

Photochemical smog is primarily caused by:

A) Volcanic eruptions
B) The reaction of automobile exhaust with sunlight
C) CFC emissions
D) Acid rain

A

B) The reaction of automobile exhaust with sunlight

223
Q

Indoor air pollution can be caused by:

A) Radon, carbon monoxide, carbon dioxide, and formaldehyde
B) Only radon
C) Only carbon monoxide
D) Dust particles

A

A) Radon, carbon monoxide, carbon dioxide, and formaldehyde

224
Q

Entropy is best described as:

A) The amount of energy required to start a reaction
B) A measure of the different ways a system can disperse its energy
C) The concentration of a substance in solution
D) A constant value for all substances

A

B) A measure of the different ways a system can disperse its energy

225
Q

According to the second law of thermodynamics, any spontaneous process must:

A) Increase the entropy of the universe
B) Decrease the entropy of the universe
C) Increase the entropy of the system only
D) Maintain constant entropy in the universe

A

A) Increase the entropy of the universe

226
Q

The standard entropy of a chemical reaction can be calculated from:

A) The molecular weight of the reactants and products
B) The absolute entropies of the reactants and products
C) The reaction rate
D) The pH of the solution

A

B) The absolute entropies of the reactants and products

227
Q

The third law of thermodynamics states that the entropy of a perfect crystalline substance is:

A) Infinite at 0 K
B) Zero at 0 K
C) 1 at 0 K
D) Equal to the energy of the substance at 0 K

A

B) Zero at 0 K

228
Q

Under conditions of constant temperature and pressure, the free-energy change (Δ𝐺) is:

A) Less than zero for a nonspontaneous process
B) Greater than zero for a spontaneous process
C) Less than zero for a spontaneous process and greater than zero for a nonspontaneous process
D) Always equal to zero

A

C) Less than zero for a spontaneous process and greater than zero for a nonspontaneous process

229
Q

The equation ΔG=ΔH−TΔS is used to:

A) Calculate the equilibrium constant of a reaction
B) Determine the spontaneity of a process
C) Measure the rate of a reaction
D) Find the concentration of products

A

B) Determine the spontaneity of a process

230
Q

The standard free-energy change (ΔGᵒ) for a reaction can be calculated from:

A) The absolute entropies of the reactants
B) The standard free energies of formation of reactants and products
C) The pH of the reactants
D) The concentration of products only

A

B) The standard free energies of formation of reactants and products

231
Q

The relationship between the equilibrium constant (K) and the standard free-energy change (ΔGᵒ) is given by the equation:

A) ΔGᵒ = RT lnK
B) ΔGᵒ = -RT ln K
C) ΔGᵒ = K/RT
D) ΔGᵒ = RT/K

A

B) ΔGᵒ = -RT ln K

232
Q

Many biological reactions are nonspontaneous but are driven by:

A) High temperatures
B) The hydrolysis of ATP, for which ΔGᵒ is negative
C) The release of oxygen
D) The absorption of carbon dioxide

A

B) The hydrolysis of ATP, for which ΔGᵒ is negative

233
Q

Redox reactions involve:

A) The transfer of protons
B) The transfer of electrons
C) The sharing of neutrons
D) Only the formation of covalent bonds

A

B) The transfer of electrons

234
Q

All electrochemical reactions involve:

A) The formation of ionic bonds
B) The transfer of electrons, making them redox reactions
C) Only spontaneous reactions
D) Only nonspontaneous reactions

A

B) The transfer of electrons, making them redox reactions

235
Q

In a galvanic cell, the process that generates electricity is:

A) A nonspontaneous reaction
B) A spontaneous redox reaction
C) A neutralization reaction
D) A reaction without any electron transfer

A

B) A spontaneous redox reaction

236
Q

In a galvanic cell, the two halves where reactions occur are known as:

A) Electrodes
B) Electrolytes
C) Half-cells
D) Batteries

A

C) Half-cells

237
Q

The electromotive force (emf) of a cell is defined as:

A) The number of electrons in the cell
B) The voltage difference between the two electrodes
C) The concentration of ions in the solution
D) The resistance of the circuit

A

B) The voltage difference between the two electrodes

238
Q

In the external circuit of a galvanic cell, electrons flow from:

A) Cathode to anode
B) Anode to cathode
C) Solution to electrodes
D) Electrolyte to salt bridge

A

B) Anode to cathode

239
Q

The quantity of electricity carried by 1 mole of electrons is:

A) 1,000 C
B) 96,500 C
C) 10,000 C
D) 500 C

A

B) 96,500 C

240
Q

Standard reduction potentials are used to predict:

A) The type of metal formed in a reaction
B) The redox products, direction, and spontaneity of reactions
C) The boiling point of solutions
D) The pH of a solution

A

B) The redox products, direction, and spontaneity of reactions

241
Q

In a spontaneous redox reaction, the decrease in free energy is equal to:

A) The energy absorbed by the system
B) The electrical work done by the system on the surroundings
C) The temperature change in the reaction
D) The increase in concentration of ions

A

B) The electrical work done by the system on the surroundings

242
Q

The Nernst equation relates:

A) The pH of a solution to its temperature
B) The emf of a cell to the concentrations of reactants and products under non-standard conditions
C) The resistance of a conductor to its length
D) The solubility of a salt to its temperature

A

B) The emf of a cell to the concentrations of reactants and products under non-standard conditions

243
Q

Batteries consist of one or more galvanic cells and are widely used as power sources. A common example of a battery used in automobiles is:

A) The Leclanché cell
B) The mercury battery
C) The lead storage battery
D) The dry cell

A

C) The lead storage battery

244
Q

The corrosion of metals, such as the rusting of iron, is an example of:

A) A galvanic cell reaction
B) An electrochemical phenomenon
C) A physical change
D) A non-redox process

A

B) An electrochemical phenomenon

245
Q

In an electrolytic cell, electric current from an external source is used to:

A) Drive a spontaneous reaction
B) Drive a nonspontaneous chemical reaction
C) Prevent electron flow
D) Measure the concentration of ions

A

B) Drive a nonspontaneous chemical reaction

246
Q

Depending on their reactivities, metals exist in nature in either:

A) Only the combined state
B) Only the free state
C) The free or combined state
D) Only in solid compounds

A

C) The free or combined state

247
Q

The process of recovering a metal from its ore generally involves:

A) Only purifying the metal
B) Preparing, separating (usually through reduction), and purifying the metal
C) Melting the metal ore directly
D) Only heating the ore

A

B) Preparing, separating (usually through reduction), and purifying the metal

248
Q

Common methods used for purifying metals include:

A) Melting and sublimation
B) Electrolysis, distillation, and zone refining
C) Freezing and filtration
D) Crystallization and sublimation

A

B) Electrolysis, distillation, and zone refining

249
Q

Metallic bonds are formed by:

A) Attraction between metal ions and electrons in a “sea” of electrons
B) Sharing of electrons between nonmetal atoms
C) Ionic bonds between cations and anions
D) Hydrogen bonds between metal atoms

A

A) Attraction between metal ions and electrons in a “sea” of electrons

250
Q

In insulators, the energy gap between the valence band and conduction band is:

A) So large that electrons cannot be promoted to the conduction band
B) Small, allowing electrons to flow freely
C) Negligible, causing high conductivity
D) Zero, allowing free movement of protons

A

A) So large that electrons cannot be promoted to the conduction band

251
Q

In semiconductors, conductivity increases with higher temperatures because:

A) Electrons cannot cross the energy gap at any temperature
B) The energy gap between the valence band and conduction band decreases
C) More electrons gain energy to reach the conduction band
D) Electrons move to the valence band

A

C) More electrons gain energy to reach the conduction band

252
Q

n-Type semiconductors and p-Type semiconductors differ in that:

A) n-Type semiconductors accept electrons, while p-Type donate electrons
B) n-Type contain donor impurities with extra electrons, while p-Type contain acceptor impurities with “positive holes”
C) Both types have the same impurities
D) Only n-Type semiconductors conduct electricity

A

B) n-Type contain donor impurities with extra electrons, while p-Type contain acceptor impurities with “positive holes”

253
Q

The alkali metals are:

A) The least reactive metals in the periodic table
B) The most reactive of all metallic elements, with an oxidation number of +1 in their compounds
C) The most stable elements, rarely reacting
D) Always forming uninegative ions

A

B) The most reactive of all metallic elements, with an oxidation number of +1 in their compounds

254
Q

The alkaline earth metals are generally:

A) More reactive than alkali metals
B) Less reactive than alkali metals, with an oxidation number of +2 in their compounds
C) Nonreactive and stable
D) Only reactive in the presence of acids

A

B) Less reactive than alkali metals, with an oxidation number of +2 in their compounds

255
Q

Aluminum does not react with water because:

A) It is an inert metal
B) It has a strong metallic bond
C) It forms a protective oxide layer; its hydroxide is amphoteric
D) It lacks valence electrons

A

C) It forms a protective oxide layer; its hydroxide is amphoteric

256
Q

A hydrogen atom typically contains:

A) One proton and no electrons
B) One proton and one electron
C) Two protons and one electron
D) One neutron and no protons

A

B) One proton and one electron

257
Q

The isotopes of hydrogen include:

A) Hydrogen-1, Hydrogen-2, and Hydrogen-3
B) Hydrogen, Helium, and Deuterium
C) Oxygen-16, Oxygen-17, and Oxygen-18
D) Only Hydrogen-1

A

A) Hydrogen-1, Hydrogen-2, and Hydrogen-3

258
Q

Important inorganic compounds of carbon include all of the following EXCEPT:

A) Cyanides
B) Sulfates
C) Carbonates and bicarbonates
D) Carbon monoxide

A

B) Sulfates

259
Q

Elemental nitrogen (N2):

A) Has a double bond and is highly reactive
B) Contains a triple bond and is very stable
C) Has a single bond and is very toxic
D) Is found mainly in compounds with sulfur

A

B) Contains a triple bond and is very stable

260
Q

Phosphorus exists in forms like white phosphorus (P4) and red phosphorus. White phosphorus is:

A) Non-toxic and stable
B) Highly toxic, reactive, and flammable
C) Only used in fertilizers
D) Always combined with metals

A

B) Highly toxic, reactive, and flammable

261
Q

The most important phosphorus compounds are:

A) Carbonates
B) Nitrates
C) Phosphates
D) Sulfates

A

C) Phosphates

262
Q

Elemental oxygen (O2) is:

A) Diamagnetic with no unpaired electrons
B) Paramagnetic with two unpaired electrons
C) Always found as ozone (O3)
D) Found only in superoxides

A

B) Paramagnetic with two unpaired electrons

263
Q

Sulfuric acid is produced by:

A) The electrolysis of sulfur dioxide
B) The contact process involving sulfur dioxide and sulfur trioxide
C) The fusion of sulfur and nitrogen
D) Decomposition of sulfur oxides

A

B) The contact process involving sulfur dioxide and sulfur trioxide

264
Q

The halogens are:

A) Non-toxic and stable elements
B) Only found as free elements in nature
C) Toxic and reactive elements found only in compounds with other elements
D) Inert and do not react with metals

A

C) Toxic and reactive elements found only in compounds with other elements

265
Q

The reactivity, toxicity, and oxidizing ability of halogens:

A) Increase from fluorine to iodine
B) Decrease from fluorine to iodine
C) Remain the same for all halogens
D) Vary randomly among the halogens

A

B) Decrease from fluorine to iodine

266
Q

Transition metals typically have:

A) Completely filled d subshells
B) Incompletely filled d subshells and a tendency to form complexes
C) Only s subshell electrons
D) No ability to form complexes

A

B) Incompletely filled d subshells and a tendency to form complexes

267
Q

The first-row transition metals (scandium to copper):

A) Are the least reactive of all transition metals
B) Are the most common transition metals with characteristic chemistry representing the entire group
C) Are only found in complex ions
D) Do not exhibit variable oxidation states

A

B) Are the most common transition metals with characteristic chemistry representing the entire group

268
Q

Complex ions consist of:

A) A metal ion surrounded by ions of the same element
B) A metal ion surrounded by ligands that donate electron pairs
C) Nonmetals bonded to other nonmetals
D) Only metal ions in isolation

A

B) A metal ion surrounded by ligands that donate electron pairs

269
Q

Coordination compounds can display:

A) Only one stable geometric structure
B) Geometric and/or optical isomerism
C) No isomerism
D) Only magnetic properties

A

B) Geometric and/or optical isomerism

270
Q

Crystal field theory explains bonding in complexes by:

A) Covalent bonding only
B) Electrostatic interactions and the splitting of d orbitals into higher and lower energy levels in an octahedral complex
C) The transfer of protons
D) Ionic interactions with no energy splitting

A

B) Electrostatic interactions and the splitting of d orbitals into higher and lower energy levels in an octahedral complex

271
Q

The difference in crystal field splitting between strong-field and weak-field ligands is that:

A) Strong-field ligands cause large splitting, while weak-field ligands cause small splitting
B) Weak-field ligands cause large splitting, while strong-field ligands cause small splitting
C) Both types of ligands cause the same amount of splitting
D) Only strong-field ligands cause any splitting

A

A) Strong-field ligands cause large splitting, while weak-field ligands cause small splitting

272
Q

Complex ions can undergo:

A) No reaction in solution
B) Ligand exchange reactions in solution
C) Combustion reactions
D) Only solid-state reactions

A

B) Ligand exchange reactions in solution

273
Q

Coordination compounds are useful in various applications, including:

A) Fuel production only
B) Antidotes for metal poisoning and in chemical analysis
C) Building materials
D) Enhancing photosynthesis in plants

A

B) Antidotes for metal poisoning and in chemical analysis

274
Q

For stable nuclei with low atomic numbers, the neutron-to-proton ratio is generally:

A) Greater than 2
B) Close to 1
C) Greater than 3
D) Close to 0

A

B) Close to 1

275
Q

All nuclei with 84 or more protons are:

A) Stable and non-radioactive
B) Unstable and radioactive
C) Non-reactive
D) Highly reactive but stable

A

B) Unstable and radioactive

276
Q

Nuclear binding energy is:

A) The amount of energy needed to ionize an atom
B) A measure of nuclear stability
C) Equal to the mass number of the nucleus
D) The kinetic energy of electrons in the nucleus

A

B) A measure of nuclear stability

277
Q

Radioactive nuclei emit:

A) Only alpha particles
B) Only gamma rays
C) Alpha particles, beta particles, positrons, or gamma rays
D) Neutrons only

A

C) Alpha particles, beta particles, positrons, or gamma rays

278
Q

The parent of a natural radioactive decay series that helps determine the age of rocks is:

A) Carbon-14
B) Uranium-238
C) Thorium-234
D) Potassium-40

A

B) Uranium-238

279
Q

Artificial radioactive elements are created by:

A) Heating elements to high temperatures
B) Bombarding elements with accelerated neutrons, protons, or alpha particles
C) Exposing elements to sunlight
D) Mixing elements with radioactive materials

A

B) Bombarding elements with accelerated neutrons, protons, or alpha particles

280
Q

Nuclear fission involves:

A) The combination of two small nuclei into one large nucleus
B) The splitting of a large nucleus into two smaller nuclei and one or more neutrons
C) The emission of gamma rays only
D) The decay of a single proton

A

B) The splitting of a large nucleus into two smaller nuclei and one or more neutrons

281
Q

Nuclear reactors use heat from nuclear fission to produce power. Important types of reactors include:

A) Only breeder reactors
B) Light water reactors, heavy water reactors, and breeder reactors
C) Cold fusion reactors
D) Only uranium reactors

A

B) Light water reactors, heavy water reactors, and breeder reactors

282
Q

Nuclear fusion, which occurs in the sun, is:

A) The process of splitting heavy nuclei
B) The combination of two light nuclei to form one heavy nucleus
C) An achievable, large-scale energy source on Earth
D) The process of electron emission in metals

A

B) The combination of two light nuclei to form one heavy nucleus

283
Q

Radioactive isotopes are valuable in medicine because:

A) They are stable and unreactive
B) They are easy to detect and useful as tracers in medical practices
C) They release no radiation
D) They are only used for imaging bones

A

B) They are easy to detect and useful as tracers in medical practices

284
Q

High-energy radiation damages living systems by:

A) Reducing temperature in cells
B) Causing ionization and forming free radicals
C) Strengthening cellular structures
D) Preventing cell division

A

B) Causing ionization and forming free radicals

285
Q

Carbon atoms can form a vast number of compounds because:

A) They only form single bonds.
B) They can link up with other carbon atoms in straight and branched chains.
C) They are highly reactive with metals.
D) They are the heaviest atoms on the periodic table.

A

B) They can link up with other carbon atoms in straight and branched chains.

286
Q

Organic compounds are derived from two main types of hydrocarbons:

A) Aromatic and saturated hydrocarbons
B) Aliphatic and aromatic hydrocarbons
C) Cyclic and branched hydrocarbons
D) Single-bonded and triple-bonded hydrocarbons

A

B) Aliphatic and aromatic hydrocarbons

287
Q

Methane (CH4) is the simplest member of which class of hydrocarbons?

A) Alkenes
B) Alkynes
C) Alkanes
D) Aromatic hydrocarbons

A

C) Alkanes

288
Q

Cyclopropane (C3H6) is the simplest member of:

A) Aromatic hydrocarbons
B) Cycloalkanes
C) Alkenes
D) Alkynes

A

B) Cycloalkanes

289
Q

Ethylene (CH2=CH2) is the simplest member of which class of hydrocarbons?

A) Alkanes
B) Aromatic hydrocarbons
C) Olefins or alkenes
D) Cycloalkanes

A

C) Olefins or alkenes

290
Q

The general formula CₙH₂ₙ−₂ represents which type of hydrocarbons?

A) Alkanes
B) Alkenes
C) Alkynes
D) Aromatic hydrocarbons

A

C) Alkynes

291
Q

Aromatic hydrocarbons are characterized by:

A) Having no double bonds
B) Containing one or more benzene rings
C) Only forming straight chains
D) Reacting only with water

A

B) Containing one or more benzene rings

292
Q

Functional groups in organic chemistry:

A) Do not affect the reactivity of molecules
B) Impart specific types of chemical reactivity to molecules
C) Are found only in aromatic hydrocarbons
D) Are limited to alcohols and esters only

A

B) Impart specific types of chemical reactivity to molecules

293
Q

Classes of compounds characterized by their functional groups include:

A) Alkanes and cycloalkanes
B) Alcohols, ethers, aldehydes, ketones, carboxylic acids, esters, and amines
C) Only alkenes and alkynes
D) Only hydrocarbons with benzene rings

A

B) Alcohols, ethers, aldehydes, ketones, carboxylic acids, esters, and amines

294
Q

Polymers are large molecules made up of small, repeating units called:

A) Ions
B) Monomers
C) Atoms
D) Radicals

A

B) Monomers

295
Q

Which of the following is NOT a natural polymer?

A) Cellulose
B) Nylon
C) Rubber
D) Nucleic acids

A

B) Nylon

296
Q

Organic polymers can be synthesized by:

A) Crystallization and filtration
B) Addition reactions or condensation reactions
C) Electrolysis
D) Distillation

A

B) Addition reactions or condensation reactions

297
Q

Stereoisomers of a polymer have different properties because:

A) They have different types of monomers
B) They are formed from non-repeating units
C) The starting units are joined in different ways
D) They do not contain any asymmetric monomers

A

C) The starting units are joined in different ways

298
Q

Synthetic rubbers include:

A) Polychloroprene and styrene-butadiene rubber
B) Polyethylene and polyvinyl chloride
C) Nylon and Dacron
D) Cellulose and rubber

A

A) Polychloroprene and styrene-butadiene rubber

299
Q

The function and properties of proteins are determined by their:

A) Solubility in water
B) Structure, which is influenced by hydrogen bonding and other intermolecular forces
C) Ability to conduct electricity
D) Size and weight alone

A

B) Structure, which is influenced by hydrogen bonding and other intermolecular forces

300
Q

The primary structure of a protein refers to:

A) Its three-dimensional folded arrangement
B) Its amino acid sequence
C) The shape defined by hydrogen bonds
D) Its ability to dissolve in water

A

B) Its amino acid sequence

301
Q

Nucleotides, the building blocks of DNA and RNA, contain:

A) Only a phosphate group and a ribose molecule
B) A purine or pyrimidine base, a sugar molecule, and a phosphate group
C) Only nitrogen and oxygen atoms
D) Only ribose molecules

A

B) A purine or pyrimidine base, a sugar molecule, and a phosphate group

302
Q

RNA differs from DNA in that:

A) RNA contains deoxyribose instead of ribose
B) RNA contains ribose and different bases than DNA
C) RNA contains the same bases and sugar as DNA
D) RNA is not involved in protein synthesis

A

B) RNA contains ribose and different bases than DNA

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