Chemistry A molecular approach

Question 1

Atoms and Molecules

Answer 1

■ All matter is composed of atoms and molecules.
■ Chemistry is the science that investigates the properties of matter
by examining the atoms and molecules that compose it.

Question 2

The Scientific Approach to Knowledge

Answer 2

■ Science begins with the observation of the physical world. A number
of related observations can be summarized in a statement or
generalization called a scientific law.
■ A hypothesis is a tentative interpretation or an explanation of observations.
One or more well-established hypotheses may prompt
the development of a scientific theory, a model for nature that explains
the underlying reasons for observations and laws.
■ Laws, hypotheses, and theories all give rise to predictions that
can be tested by experiments, carefully controlled procedures
designed to produce critical new observations. If scientists cannot
confirm the predictions, they must modify or replace the law,
hypothesis, or theory.

Question 3

The Classification of Matter

Answer 3

■ We classify matter according to its state (solid, liquid, or gas) or
according to its composition (pure substance or mixture).
■ A pure substance can either be an element, which cannot be
chemically broken down into simpler substances, or a compound,
which is composed of two or more elements in fixed proportions.
■ A mixture can be either homogeneous, with the same composition
throughout, or heterogeneous, with different compositions
in different regions.

Question 4

The Properties of Matter

Answer 4

■ We classify the properties of matter into two types: physical and
chemical. Matter displays its physical properties without changing
its composition.
■ Changes in matter in which composition does not change are
physical changes. Changes in matter in which composition does
change are chemical changes.

Question 5

Energy

Answer 5

■ In chemical and physical changes, matter often exchanges energy
with its surroundings. In these exchanges, the total energy is
always conserved; energy is neither created nor destroyed.
■ Systems with high potential energy tend to change in the
direction of lower potential energy, releasing energy into the surroundings.

Question 6

The Units of Measurement and Significant Figures

Answer 6

■ Scientists use SI units, which are based on the metric system. The
SI base units include the meter (m) for length, the kilogram (kg)
for mass, the second (s) for time, and the kelvin (K) for temperature.
■ Derived units are formed from a combination of other units.
Common derived units include those for volume (cm^3 or m^3) and
density (g/cm^3).
■ The number of digits in a reported measurement reflects the
uncertainty in the measurement. Significant figures are the
non–place-holding digits in a reported number.

Question 7

Brownian Motion

Answer 7

■ Brownian motion is the erratic, jittery motion of small particles
that was first observed by Robert Brown in 1827. The description of
Brownian motion by Einstein in 1905 and confirmation by Perrin
in 1908 removed any lingering doubt about the particulate nature
of matter.

Question 8

The Atomic Theory

Answer 8

■ Each element is composed of indestructible particles called
atoms.
■ All atoms of a given element have the same mass and other
properties.
■ Atoms combine in simple, whole-number ratios to form compounds.
■ Atoms of one element cannot change into atoms of another element.
In a chemical reaction, atoms change the way that they are
bound together with other atoms to form a new substance.

Question 9

The Electron

Answer 9

■ J. J. Thomson discovered the electron in the late 1800s through
experiments with cathode rays. He deduced that electrons are
negatively charged, and he measured their charge-to-mass ratio.
■ Robert Millikan measured the charge of the electron, which—in
conjunction with Thomson’s results—led to the calculation of the
mass of an electron.

Question 10

The Nuclear Atom

Answer 10

■ In 1909, Ernest Rutherford probed the inner structure of the atom
by working with a form of radioactivity called alpha radiation and
developed the nuclear theory of the atom.
■ Nuclear theory states that the atom is mainly empty space, with
most of its mass concentrated in a tiny region called the nucleus
and most of its volume occupied by relatively light electrons.

Question 11

Subatomic Particles

Answer 11

■ Atoms are composed of three fundamental particles: the proton
(1 amu, +1 charge), the neutron (1 amu, 0 charge), and the electron
(~0 amu, -1 charge).
■ The number of protons in the nucleus of the atom is its atomic
number (Z) and defines the element.
■ The sum of the number of protons and neutrons is the mass number
(A).
■ Atoms of an element that have different numbers of neutrons
(and therefore different mass numbers) are isotopes.
■ Atoms that lose or gain electrons become charged and are ions.
Cations are positively charged and anions are negatively charged.

Question 12

The Periodic Table

Answer 12

■ The periodic table tabulates all known elements in order of
increasing atomic number.
■ The periodic table is arranged so that similar elements are grouped
together in columns.
■ Elements on the left side and in the center of the periodic table are
metals and tend to lose electrons in chemical changes.
■ Elements on the upper right side of the periodic table are nonmetals
and tend to gain electrons in chemical changes.
■ Elements located on the boundary between metals and nonmetals
are metalloids.

Question 13

Atomic Mass and the Mole

Answer 13

■ The atomic mass of an element, listed directly below its symbol
in the periodic table, is a weighted average of the masses of the
naturally occurring isotopes of the element.
■ One mole of an element is the amount of that element that contains
Avogadro’s number (6.022 * 10^23) of atoms.
■ Any sample of an element with a mass (in grams) that equals its
atomic mass contains one mole of the element. For example, the
atomic mass of carbon is 12.011 amu; therefore, 12.011 g of carbon
contains 1 mol of carbon atoms.

Question 14

Chemical Bonds

Answer 14

■ Chemical bonds, the forces that hold atoms together in compounds,
arise from the interactions between nuclei and electrons
in atoms.
■ In an ionic bond, one or more electrons are transferred from one
atom to another, forming a cation (positively charged) and an
anion (negatively charged). The two ions are drawn together by
the attraction between the opposite charges.
■ In a covalent bond, one or more electrons are shared between two
atoms. The atoms are held together by the attraction between
their nuclei and the shared electrons.

Question 15

Representing Molecules and Compounds

Answer 15

■ A compound is represented with a chemical formula, which indicates
the elements present and the number of atoms of each.
■ An empirical formula gives only the relative number of atoms,
while a molecular formula gives the actual number of atoms present
in the molecule.
■ Structural formulas show how atoms are bonded together, while
molecular models portray the geometry of the molecule.
■ Compounds can be divided into two types: molecular compounds,
formed between two or more covalently bonded nonmetals, and
ionic compounds, usually formed between a metal ionically
bonded to one or more nonmetals. The smallest identifiable unit
of a molecular compound is a molecule, and the smallest identifiable
unit of an ionic compound is a formula unit: the smallest
electrically neutral collection of ions.
■ Elements can also be divided into two types: molecular elements,
which occur as (mostly diatomic) molecules, and atomic elements,
which occur as individual atoms.

Question 16

Formula Mass and Mole Concept for Compounds

Answer 16

■ The formula mass of a compound is the sum of the atomic masses
of all the atoms in the chemical formula. Like the atomic masses
of elements, the formula mass characterizes the average mass of a
molecule (or a formula unit).
■ The mass of one mole of a compound is the molar mass of that
compound and equals its formula mass (in grams).

Question 17

Chemical Composition

Answer 17

■ The mass percent composition of a compound indicates each
element’s percentage of the total compound’s mass. We can
determine
the mass percent composition from the compound’s
chemical formula and the molar masses of its elements.
■ The chemical formula of a compound provides the relative number
of atoms (or moles) of each element in a compound, and we
can therefore use it to determine numerical relationships between
moles of the compound and moles of its constituent elements. We
can extend this relationship to mass by using the molar masses of
the compound and its constituent elements.
■ If the mass percent composition and molar mass of a compound are
known, we can determine its empirical and molecular formulas.

Question 18

Organic Compounds

Answer 18

■ Organic compounds are composed of carbon, hydrogen, and a few
other elements such as nitrogen, oxygen, and sulfur.
■ The simplest organic compounds are hydrocarbons, compounds
composed of only carbon and hydrogen.
■ Hydrocarbons are categorized into three types based on the bonds
they contain: alkanes contain single bonds, alkenes contain double
bonds, and alkynes contain triple bonds.
■ All other organic compounds can be thought of as hydrocarbons
with one or more functional groups—characteristic atoms or
groups of atoms.
■ Common functionalized hydrocarbons include alcohols, ethers,
aldehydes, ketones, carboxylic acids, esters, and amines.

Question 19

Climate Change and the Combustion of
Fossil Fuels

Answer 19

■ Greenhouse gases warm Earth by trapping some of the sunlight that
penetrates Earth’s atmosphere. Global warming, resulting from rising
atmospheric carbon dioxide levels, is potentially harmful.
■ The largest atmospheric carbon dioxide source is the burning of
fossil fuels. This can be verified by reaction stoichiometry.

Question 20

Writing and Balancing Chemical Equations

Answer 20

■ In chemistry, we represent chemical reactions with chemical
equations. The substances on the left-hand side of a chemical
equation are the reactants, and the substances on the right-hand
side are the products.
■ Chemical equations are balanced when the number of each type
of atom on the left side of the equation is equal to the number on
the right side.

Question 21

Reaction Stoichiometry

Answer 21

■ Reaction stoichiometry refers to the numerical relationships
between the reactants and products in a balanced chemical equation.
■ Reaction stoichiometry allows us to predict, for example, the
amount of product that can be formed for a given amount of reactant,
or how much of one reactant is required to react with a given
amount of another.

Question 22

Limiting Reactant, Theoretical Yield, and Percent
Yield

Answer 22

■ When a chemical reaction actually occurs, the reactants are usually
not present in the exact stoichiometric ratios specified by the
balanced chemical equation. The limiting reactant is the one that
is available in the smallest stoichiometric quantity—it will be
completely consumed in the reaction, and it limits the amount of
product that can be made.
■ Any reactant that does not limit the amount of product is in excess.
■ The amount of product that can be made from the limiting reactant
is the theoretical yield.
■ The actual yield—always equal to or less than the theoretical
yield—is the amount of product that is actually made when the
reaction is carried out.
■ The percentage of the theoretical yield that is actually produced
when the reaction is carried out is the percent yield.

Question 23

Combustion, Alkali Metals, and Halogens

Answer 23

■ In a combustion reaction, a substance reacts with oxygen—emitting
heat and forming one or more oxygen-containing products. The
alkali
metals react with nonmetals, losing electrons in the process.
■ The halogens react with many metals to form metal halides. They
also react with hydrogen to form hydrogen halides and with one
another to form interhalogen compounds.

Question 24

Solution Concentration and Stoichiometry

Answer 24

■ An aqueous solution is a homogeneous mixture of water (the
solvent) with another substance (the solute).
■ We express the concentration of a solution in molarity, the
number of moles of solute per liter of solution.
■ We can use the molarities and volumes of reactant solutions to
predict the amount of product that forms in an aqueous reaction.

Question 25

Aqueous Solutions and Precipitation
Reactions

Answer 25

■ Solutes that completely dissociate (or completely ionize in the
case of the strong acids) to ions in solution are strong electrolytes,
and their solutions are good conductors of electricity.
■ Solutes that only partially dissociate (or partially ionize) are weak
electrolytes.
■ Solutes that do not dissociate (or ionize) are nonelectrolytes.
■ A substance that dissolves in water to form a solution is soluble.
■ In a precipitation reaction, we mix two aqueous solutions and
a solid (precipitate) forms.
■ The solubility rules are an empirical set of guidelines that help predict
the solubilities of ionic compounds; these rules are especially
useful when determining whether or not a precipitate will form.

Question 26

Equations for Aqueous Reactions

Answer 26

■ We can represent an aqueous reaction with a molecular equation,
which shows the complete neutral formula for each compound in
the reaction.
■ We can also represent an aqueous reaction with a complete ionic
equation, which shows the dissociated nature of strong electrolytes.
■ A third representation of an aqueous reaction is the net ionic
equation, in which the spectator ions—those that do not change
in the course of the reaction—are left out of the equation.

Question 27

Acid–Base and Gas-Evolution Reactions

Answer 27

■ In an acid–base reaction, an acid, a substance that produces H+
in solution, reacts with a base, a substance that produces OH- in
solution, and the two neutralize each other, producing water (or
in some cases a weak electrolyte).
■ An acid–base titration is a laboratory procedure in which a
reaction is carried to its equivalence point—the point at which the
reactants are in exact stoichiometric proportions; titrations are
useful in determining the concentrations of unknown solutions.
■ In gas-evolution reactions, two aqueous solutions combine, and a
gas is produced.

Question 28

Oxidation–Reduction Reactions

Answer 28

■ In oxidation–reduction reactions, one substance transfers
electrons to another substance.
■ In oxidation–reduction reactions, the substance that loses
electrons is oxidized, and the substance that gains them is reduced.
■ An oxidation state is a fictitious charge given to each atom in an
oxidation–reduction reaction by assigning all shared electrons
to the atom with the greater attraction for those electrons. An
oxidation state is an imposed electronic bookkeeping scheme, not
an actual physical state.
■ The oxidation state of an atom increases upon oxidation and
decreases upon reduction.
■ The activity series of metals can be used to predict spontaneous
redox reaction. Any half-reaction in the series is spontaneous
when paired with any reverse half-reaction below it.

Question 29

Pressure

Answer 29

■ Gas pressure is the force per unit area that results from gas particles
colliding with the surfaces around them. We use various units to
measure pressure, including mmHg, torr, Pa, psi, in Hg, and atm.

Question 30

The Simple Gas Laws

Answer 30

■ The simple gas laws express relationships between pairs of variables
when other variables are constant.
■ Boyle’s law states that the volume of a gas is inversely proportional
to its pressure.
■ Charles’s law states that the volume of a gas is directly
proportional to its temperature.
■ Avogadro’s law states that the volume of a gas is directly proportional
to the amount (in moles).

Question 31

The Ideal Gas Law and Its Applications

Answer 31

■ The ideal gas law, PV = nRT, describes the relationship among all
four gas variables and contains the simple gas laws within it.
■ We can use the ideal gas law to find one of the four variables given
the other three. We can use it to calculate the molar volume
of an ideal gas, which is 22.4 L at STP, and to calculate the density
and molar mass of a gas.

Question 32

Mixtures of Gases and Partial Pressures

Answer 32

■ In a mixture of gases, each gas acts independently of the others
so that any overall property of the mixture is the sum of the
properties of the individual components.
■ The pressure of any individual component is its partial
pressure.

Question 33

Gas Stoichiometry

Answer 33

■ In reactions involving gaseous reactants and products, we often
report quantities in volumes at specified pressures and temperatures.
We can convert these quantities to amounts (in moles)
using the ideal gas law. Then we can use the stoichiometric coefficients
from the balanced equation to determine the stoichiometric
amounts of other reactants or products.
■ The general form for these types of calculations is: volume A → amount A (in moles) → amount B (in moles) → quantity of B
(in desired units).
■ In cases where the reaction is carried out at STP, we can use the
molar volume at STP (22.4 L = 1 mol) to convert between volume
in liters and amount in moles.

Question 34

Kinetic Molecular Theory and Its Applications

Answer 34

■ Kinetic molecular theory is a quantitative model for gases. The
theory has three main assumptions: (1) gas particles are negligibly
small, (2) the average kinetic energy of a gas particle is proportional
to the temperature in kelvins, and (3) the collision of one gas
particle with another is completely elastic (the particles do not stick
together). The gas laws all follow from the kinetic molecular theory.
■ We can use kinetic molecular theory to derive the expression
for the root mean square velocity of gas particles. This velocity
is inversely proportional to the molar mass of the gas, and
therefore—at a given temperature—smaller gas particles are (on
average) moving more quickly than larger ones.
■ The kinetic molecular theory also allows us to predict the
mean free path of a gas particle (the distance it travels between
collisions) and relative rates of diffusion or effusion.

Question 35

Real Gases

Answer 35

■ Real gases differ from ideal gases to the extent that they do not
always fit the assumptions of kinetic molecular theory.
■ These assumptions tend to break down at high pressures where the
volume is higher than predicted for an ideal gas because the particles
are no longer negligibly small compared to the space between them.
■ The assumptions also break down at low temperatures where the
pressure is lower than predicted because the attraction between
molecules combined with low kinetic energies causes partially
inelastic collisions.
■ The van der Waals equation predicts gas properties under
nonideal
conditions.

Question 36

The Nature of Energy and Thermodynamics

Answer 36

■ Energy, which is measured in the SI unit of joules (J), is the capacity
to do work.
■ Work is the result of a force acting through a distance.
■ Many different kinds of energy exist, including kinetic energy,
thermal energy, potential energy, and chemical energy, a type of
potential energy associated with the relative positions of electrons
and nuclei in atoms and molecules.
■ The first law of thermodynamics states that energy can be converted
from one form to another, but the total amount of energy is
always conserved.
■ The internal energy (E) of a system is the sum of all of its kinetic
and potential energy. Internal energy is a state function, which
means that it depends only on the state of the system and not on
the pathway by which it got to that state.
■ A chemical system exchanges energy with its surroundings
through heat (the transfer of thermal energy caused by a temperature
difference) or work. The total change in internal energy is the
sum of these two quantities.

Question 37

Heat and Work

Answer 37

■ We quantify heat with the equation q = m * Cs * ΔT. In this
expression,
Cs is the specific heat capacity, the amount of heat
required
to change the temperature of 1 g of the substance by 1 °C.
Compared to most substances, water has a very high heat capacity—
it takes a lot of heat to change its temperature.
■ The type of work most characteristic of chemical reactions is
pressure–
volume work, which occurs when a gas expands against
an external pressure. Pressure–volume work can be quantified
with the equation w = -Pext ΔV.
■ The change in internal energy (ΔE) that occurs during a chemical
reaction is the sum of the heat (q) exchanged and the work (w)
done: ΔE = q + w.

Question 38

Enthalpy

Answer 38

■ The heat evolved in a chemical reaction occurring at constant
pressure is the change in enthalpy (ΔH) for the reaction. Like
internal energy, enthalpy is a state function.
■ An endothermic reaction has a positive enthalpy of reaction; an
exothermic reaction has a negative enthalpy of reaction.
■ We can use the enthalpy of reaction to stoichiometrically determine
the heat evolved when a specific amount of reactant reacts.

Question 39

Calorimetry

Answer 39

■ Calorimetry is a method of measuring ΔE or ΔH for a reaction.
■ In bomb calorimetry, the reaction is carried out under conditions
of constant volume, so ΔE = qv. We can therefore use the temperature
change of the calorimeter to calculate ΔE for the reaction.
■ When a reaction takes place at constant pressure, energy may be
released both as heat and as work.
■ In coffee-cup calorimetry, a reaction is carried out under atmospheric
pressure in a solution, so q = ΔH. We use the temperature
change of the solution to calculate ΔH for the reaction.

Question 40

Calculating 𝚫Hrxn

Answer 40

■ We can calculate the enthalpy of reaction (ΔHrxn) from known
thermochemical data using the following relationships: (a) when
a reaction is multiplied by a factor, ΔHrxn is multiplied by the
same factor; (b) when a reaction is reversed, ΔHrxn changes sign;
and (c) if a chemical reaction can be expressed as a sum of two or
more steps, ΔHrxn is the sum of the ΔH’s for the individual steps
(Hess’s law). We can use these relationships to determine the
enthalpy change of an unknown reaction from reactions with
known enthalpy changes.
■ A second method to calculate ΔHrxn from known thermochemical
data involves using tabulated standard enthalpies of formation for
the reactants and products of the reaction. These are usually tabulated
for substances in their standard states, and the enthalpy of
reaction
is called the standard enthalpy of reaction (ΔH °rxn). For any
reaction, we obtain ΔH °rxn by subtracting the sum of the enthalpies
of formation of the reactants multiplied by their stoichiometric
coefficients
from the sum of the enthalpies of formation
of the
products multiplied by their stoichiometric coefficients.

Question 41

Environmental Problems Associated with Fossil
Fuel Use

Answer 41

■ Fossil fuels are nonrenewable fuels; once humans consume them,
they cannot be replaced.
■ At current rates of consumption, natural gas and petroleum
reserves will be depleted in 50–100 years.
■ In addition to their limited supply, the products of the combustion
of fossil fuels—directly or indirectly formed—contribute to
environmental problems including air pollution, acid rain, and
global climate change, which involves an increase in Earth’s average
temperature caused by CO2 emission.

Question 42

The Realm of Quantum Mechanics

Answer 42

■ The theory of quantum mechanics explains the behavior of particles,
such as photons (particles of light) and electrons, in the
atomic and subatomic realms.
■ Since the electrons of an atom determine many of its chemical
and physical properties, quantum mechanics is foundational to
understanding chemistry.

Question 43

The Nature of Light

Answer 43

■ Light is a type of electromagnetic radiation—a form of energy
embodied in oscillating electric and magnetic fields that travels
through space at 3.00 * 10^8 m/s.
■ The wave nature of light is characterized by its wavelength—the
distance between wave crests—and its ability to experience interference
(constructive or destructive) and diffraction.
■ The particle nature of light is characterized by the specific quantity
of energy carried in each photon.
■ The electromagnetic spectrum includes all wavelengths of electromagnetic
radiation from gamma rays (high energy per photon,
short wavelength) to radio waves (low energy per photon, long
wavelength). Visible light is a tiny sliver in the middle of the electromagnetic
spectrum.

Question 44

Atomic Spectroscopy

Answer 44

■ Atomic spectroscopy is study of the light absorbed and emitted by
atoms when an electron makes a transition from one energy level
to another.
■ The wavelengths absorbed or emitted in atomic spectra depend
on the energy differences between the levels involved in the transition;
large energy differences result in short wavelengths, and
small energy differences result in long wavelengths.

Question 45

The Wave Nature of Matter

Answer 45

■ Electrons have a wave nature with an associated wavelength; the
de Broglie relation quantifies the wavelength of an electron.
■ The wave nature and particle nature of matter are complementary—
the more we know of one, the less we know of the other.
■ The wave–particle duality of electrons is quantified in
Heisenberg’s uncertainty principle, which states that there is a
limit to how well we can know both the position of an electron
(associated with the electron’s particle nature) and the velocity
times the mass of an electron (associated with the electron’s wave
nature)—the more accurately one is measured, the greater the uncertainty
in measurement of the other.
■ The inability to simultaneously know both the position and velocity
of an electron results in indeterminacy, the inability to predict
a trajectory for an electron. Consequently, electron behavior is
described differently than the behavior of everyday-sized particles.
■ The trajectory we normally associate with macroscopic objects
is replaced, for electrons, with statistical descriptions that show
not the electron’s path, but the region where it is most likely to be
found.

Question 46

The Quantum-Mechanical
Model of the Atom

Answer 46

■ The most common way to describe electrons in atoms according
to quantum mechanics is to solve the Schrödinger equation for
the energy states of the electrons within the atom. When the electron
is in these states, its energy is well defined but its position is
not. The position of an electron is described by a probability distribution
map called an orbital.
■ The solutions to the Schrödinger equation (including the energies
and orbitals) are characterized by four quantum numbers: n, l, ml,
and ms.
■ The principal quantum number (n) determines the energy of the
electron and the size of the orbital; the angular momentum quantum
number (l) determines the shape of the orbital; the magnetic
quantum number (ml) determines the orientation of the orbital;
and the spin quantum number (ms) specifies the orientation of
the spin of the electron.

Question 47

Periodic Properties and the Development of the
Periodic Table

Answer 47

■ In the nineteenth century, Dmitri Mendeleev arranged the elements
in an early version of the periodic table so that atomic mass
increased from left to right in a row and elements with similar
properties fell in the same columns.
■ Periodic properties are predictable based on an element’s position
within the periodic table. Periodic properties include atomic
radius, ionization energy, electron affinity, density, and metallic
character.
■ Quantum mechanics explains the periodic table by describing
how electrons fill the quantum-mechanical orbitals within the
atoms that compose the elements.

Question 48

Electron Configurations

Answer 48

An electron configuration for an atom shows which quantummechanical
orbitals the atom’s electrons occupy. For example, the
electron configuration of helium (1s^2) indicates that helium’s two
electrons occupy the 1s orbital.
■ The order of filling quantum-mechanical orbitals in multielectron
atoms is 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s.
■ According to the Pauli exclusion principle, each orbital can hold a
maximum of two electrons and those electrons have opposing spins.
■ According to Hund’s rule, orbitals of the same energy first fill
singly with electrons having parallel spins before pairing.

Question 49

Electron Configurations and the Periodic Table

Answer 49

■ Because quantum-mechanical orbitals fill sequentially with
increasing atomic number, we can infer the electron configuration
of an element from its position in the periodic table.
■ The most stable electron configurations are those with completely
full s and p sublevels. Therefore, the most stable and unreactive
elements—those with the lowest energy electron configurations—
are the noble gases.
■ Elements with one or two valence electrons are among the most
active metals, readily losing their valence electrons to attain noble
gas configurations.
■ Elements with six or seven valence electrons are among the most
active nonmetals, readily gaining enough electrons to attain a
noble gas configuration.