2. Physics and the Quantum Mechanical Model
Neon advertising signs are
formed from glass tubes bent
in various shapes. An electric
current passing through the
gas in each glass tube makes
the gas glow with its own
characteristic color. You will
learn why each gas glows
with a specific color of light.
5.3
4. Light
The amplitude of a wave is
the wave’s height from zero
to the crest.
The wavelength,
represented by (the Greek
letter lambda), is the distance
between the crests.
5.3
5. Light
The frequency, represented
by (the Greek letter nu), is
the number of wave cycles to
pass a given point per unit of
time.
The SI unit of cycles per
second is called a hertz (Hz).
5.3
7. Light
The product of the frequency
and wavelength always
equals a constant (c), the
speed of light.
5.3
8. Light
According to the wave model, light
consists of electromagnetic waves.
Electromagnetic radiation
includes radio waves,
microwaves, infrared waves,
visible light, ultraviolet waves, X-
rays, and gamma rays.
All electromagnetic waves travel
in a vacuum at a speed of 2.998
108 m/s.
5.3
9. Light
Sunlight consists of light with a
continuous range of wavelengths
and frequencies.
When sunlight passes through a
prism, the different frequencies
separate into a spectrum of
colors.
In the visible spectrum, red light
has the longest wavelength and
the lowest frequency.
5.3
17. Atomic Spectra
When atoms absorb
energy, electrons move
into higher energy
levels. These electrons
then lose energy by
emitting light when they
return to lower energy
levels.
5.3
18. A prism separates light into the colors it
contains. When white light passes through a
prism, it produces a rainbow of colors.
5.3
19. When light from a helium lamp passes through
a prism, discrete lines are produced.
5.3
20. The frequencies of light emitted by an element
separate into discrete lines to give the atomic
emission spectrum of the element.
5.3
Mercury Nitrogen
21. An Explanation of
Atomic Spectra
An Explanation of
Atomic Spectra
How are the
frequencies of light
an atom emits
related to changes of
electron energies?
5.3
22. An Explanation of Atomic Spectra
In the Bohr model, the lone electron in the
hydrogen atom can have only certain
specific energies.
When the electron has its lowest
possible energy, the atom is in its
ground state.
Excitation of the electron by absorbing
energy raises the atom from the ground
state to an excited state.
A quantum of energy in the form of light
is emitted when the electron drops back
to a lower energy level.
5.3
23. An Explanation of Atomic
Spectra
The light emitted by an
electron moving from a
higher to a lower energy
level has a frequency
directly proportional to the
energy change of the
electron.
5.3
24. The three groups of lines in the hydrogen
spectrum correspond to the transition of electrons
from higher energy levels to lower energy levels.
5.3
26. Quantum Mechanics
In 1905, Albert Einstein successfully
explained experimental data by
proposing that light could be
described as quanta of energy.
The quanta behave as if they
were particles.
Light quanta are called photons.
In 1924, De Broglie developed an
equation that predicts that all
moving objects have wavelike
behavior.
5.3
27. Quantum Mechanics
Today, the wavelike properties of beams of electrons are
useful in magnifying objects. The electrons in an electron
microscope have much smaller wavelengths than visible
light. This allows a much clearer enlarged image of a very
small object, such as this mite.
5.3
28. Quantum Mechanics
Classical mechanics
adequately describes the
motions of bodies much larger
than atoms, while quantum
mechanics describes the
motions of subatomic particles
and atoms as waves.
5.3
29. Quantum Mechanics
The Heisenberg uncertainty principle
states that it is impossible to know exactly
both the velocity and the position of a
particle at the same time.
This limitation is critical in dealing with
small particles such as electrons.
This limitation does not matter for
ordinary-sized object such as cars or
airplanes.
5.3
31. Models of the Atom
The scale model shown is a physical model. However, not
all models are physical. In fact, several theoretical models of
the atom have been developed over the last few hundred
years. You will learn about the currently accepted model of
how electrons behave in atoms.
5.1
32. The Development of Atomic
Models
The Development of
Atomic Models
What was inadequate
about Rutherford’s
atomic model?
5.1
33. The Development of Atomic
Models
Rutherford’s atomic model could not
explain the chemical properties of
elements.
Rutherford’s atomic model could not
explain why objects change color when
heated.
5.1
34. The Development of Atomic
Models
The timeline shoes the development of
atomic models from 1803 to 1911.
5.1
35. The Development of Atomic
Models
The timeline shows the development of
atomic models from 1913 to 1932.
5.1
36. The Bohr Model
The Bohr Model
What was the new
proposal in the
Bohr model of the
atom?
5.1
37. The Bohr Model
Bohr proposed that
an electron is found
only in specific
circular paths, or
orbits, around the
nucleus.
5.1
38. The Bohr Model
Each possible electron orbit in
Bohr’s model has a fixed energy.
The fixed energies an electron
can have are called energy
levels.
A quantum of energy is the
amount of energy required to
move an electron from one energy
level to another energy level.
5.1
39. Like the rungs of the
strange ladder, the
energy levels in an atom
are not equally spaced.
The higher the energy
level occupied by an
electron, the less energy
it takes to move from
that energy level to the
next higher energy level.
5.1
40. The Quantum Mechanical
Model
The Quantum
Mechanical Model
What does the
quantum mechanical
model determine
about the electrons
in an atom?
5.1
41. The Quantum Mechanical
Model
The quantum
mechanical model
determines the allowed
energies an electron
can have and how
likely it is to find the
electron in various
locations around the
nucleus.
5.1
42. The Quantum Mechanical
Model
Austrian physicist Erwin Schrödinger
(1887–1961) used new theoretical
calculations and results to devise and
solve a mathematical equation
describing the behavior of the electron
in a hydrogen atom.
The modern description of the electrons
in atoms, the quantum mechanical
model, comes from the mathematical
solutions to the Schrödinger equation.
5.1
43. The Quantum Mechanical
Model
The propeller blade has the same probability of being anywhere in the
blurry region, but you cannot tell its location at any instant. The
electron cloud of an atom can be compared to a spinning airplane
propeller.
5.1
44. The Quantum Mechanical
Model
In the quantum mechanical model, the probability of finding an electron
within a certain volume of space surrounding the nucleus can be
represented as a fuzzy cloud. The cloud is more dense where the
probability of finding the electron is high.
5.1
46. Atomic Orbitals
An atomic orbital is often
thought of as a region of
space in which there is a
high probability of finding
an electron.
Each energy sublevel
corresponds to an orbital
of a different shape,
which describes where
the electron is likely to be
found.
5.1
47. Atomic Orbitals
Different atomic orbitals are
denoted by letters. The s orbitals
are spherical, and p orbitals are
dumbbell-shaped.
5.1
48. Atomic Orbitals
Four of the five d orbitals have
the same shape but different
orientations in space.
5.1
49. The numbers and kinds of atomic orbitals
depend on the energy sublevel.
5.1
50. Atomic Orbitals
The number of
electrons allowed in
each of the first four
energy levels are
shown here.
5.1
51. Electron Arrangement
in Atoms
If this rock were to tumble
over, it would end up at a
lower height. It would have
less energy than before, but
its position would be more
stable. You will learn that
energy and stability play an
important role in
determining how electrons
are configured in an atom.
5.2
53. Electron Configurations
The ways in which electrons are
arranged in various orbitals around
the nuclei of atoms are called
electron configurations.
Three rules—the aufbau principle,
the Pauli exclusion principle, and
Hund’s rule—tell you how to find
the electron configurations of
atoms.
5.2
54. Aufbau Principle
According to the aufbau principle, electrons occupy the
orbitals of lowest energy first. In the aufbau diagram
below, each box represents an atomic orbital.
5.2
55. Electron Configurations
Pauli Exclusion Principle
According to the Pauli exclusion
principle, an atomic orbital may
describe at most two electrons.
To occupy the same orbital, two
electrons must have opposite
spins; that is, the electron spins
must be paired.
5.2
56. Electron Configurations
Hund’s Rule
Hund’s rule states that
electrons occupy orbitals of the
same energy in a way that
makes the number of electrons
with the same spin direction as
large as possible.
5.2
63. Exceptional Electron
Configurations
Some actual electron
configurations differ from those
assigned using the aufbau
principle because half-filled
sublevels are not as stable as filled
sublevels, but they are more stable
than other configurations.
5.2
64. Exceptional Electron
Configurations
Exceptions to the aufbau
principle are due to subtle
electron-electron
interactions in orbitals with
very similar energies.
Copper has an electron
configuration that is an
exception to the aufbau
principle.
5.2
66. 6.1
Organizing the Elements
In a self-service store, the
products are grouped
according to similar
characteristics. With a
logical classification system,
finding and comparing
products is easy. You will
learn how elements are
arranged in the periodic
table and what that
arrangement reveals about
the elements.
67. Searching For an Organizing
Principle
Searching For an
Organizing Principle
How did chemists begin
to organize the known
elements?
6.1
68. Searching For an Organizing
Principle
Chemists used the
properties of elements to
sort them into groups.
6.1
69. Searching For an Organizing
Principle
Chlorine, bromine, and iodine have very
similar chemical properties.
6.1
71. Mendeleev’s Periodic Table
Mendeleev arranged the
elements in his periodic table in
order of increasing atomic
mass.
The periodic table can be used
to predict the properties of
undiscovered elements.
6.1
74. The Periodic Law
In the modern periodic table, elements are
arranged in order of increasing atomic number.
6.1
75. The Periodic Law
The periodic law: When elements are
arranged in order of increasing atomic
number, there is a periodic repetition of
their physical and chemical properties.
The properties of the elements within a
period change as you move across a
period from left to right.
The pattern of properties within a period
repeats as you move from one period to
the next.
6.1
77. Metals, Nonmetals, and
Metalloids
Three classes of elements
are metals, nonmetals, and
metalloids.
Across a period, the
properties of elements
become less metallic and
more nonmetallic.
6.1
86. Metals, Nonmetals, and
Metalloids
Nonmetals
In general, nonmetals are poor
conductors of heat and electric
current.
Most nonmetals are gases at
room temperature.
A few nonmetals are solids,
such as sulfur and phosphorus.
One nonmetal, bromine, is a
dark-red liquid.
6.1
87. Metals, Nonmetals, and
Metalloids
Metalloids
A metalloid generally has
properties that are similar to
those of metals and
nonmetals.
The behavior of a metalloid
can be controlled by changing
conditions.
6.1
88. Metals, Nonmetals, and
Metalloids
If a small amount of boron is mixed with silicon,
the mixture is a good conductor of electric
current. Silicon can be cut into wafers, and
used to make computer chips.
6.1
90. 6.2
Classifying the Elements
A coin may contain much
information in a small
space—its value, the year it
was minted, and its country
of origin. Each square in a
periodic table also contains
information. You will learn
what types of information are
usually listed in a periodic
table.
91. 6.2
Squares in the Periodic Table
Squares in the
Periodic Table
What type of
information can be
displayed in a
periodic table?
92. Squares in the Periodic Table
The periodic table displays the symbols and names
of the elements, along with information about the
structure of their atoms.
6.2
93. Squares in the Periodic Table
The background colors in the
squares are used to distinguish
groups of elements.
The Group 1A elements are
called alkali metals.
The Group 2A elements are
called alkaline earth metals.
The nonmetals of Group 7A are
called halogens.
6.2
96. Electron Configurations in
Groups
Elements can be sorted into
noble gases, representative
elements, transition metals, or
inner transition metals based
on their electron configurations.
6.2
98. Electron Configurations in
Groups
The Noble Gases
The noble gases are the elements in
Group 8A of the periodic table. The
electron configurations for the first four
noble gases in Group 8A are listed
below.
6.2
99. Electron Configurations in
Groups
The Representative Elements
Elements in groups 1A through 7A are often
referred to as representative elements
because they display a wide range of
physical and chemical properties.
The s and p sublevels of the highest
occupied energy level are not filled.
The group number equals the number of
electrons in the highest occupied energy
level.
6.2
106. Transition Elements
Transition Elements
There are two types of transition
elements—transition metals and
inner transition metals. They are
classified based on their
electron configurations.
6.2
107. Transition Elements
In atoms of a transition metal,
the highest occupied s sublevel
and a nearby d sublevel contain
electrons.
In atoms of an inner transition
metal, the highest occupied s
sublevel and a nearby f sublevel
generally contain electrons.
6.2
113. Periodic Trends
Sodium chloride (table
salt) produced the
geometric pattern in the
photograph. Such a
pattern can be used to
calculate the position of
nuclei in a solid. You will
learn how properties such
as atomic size are related
to the location of elements
in the periodic table.
6.3
114. Trends in Atomic Size
Trends in Atomic Size
What are the trends
among the elements
for atomic size?
6.3
115. Trends in Atomic Size
The atomic radius is one half of the distance
between the nuclei of two atoms of the same
element when the atoms are joined.
6.3
116. Trends in Atomic Size
Group and Periodic Trends in
Atomic Size
In general, atomic size
increases from top to bottom
within a group and
decreases from left to right
across a period.
6.3
122. Ions
Some compounds are composed
of particles called ions.
An ion is an atom or group of
atoms that has a positive or
negative charge.
A cation is an ion with a positive
charge.
An anion is an ion with a
negative charge.
6.3
123. Trends in Ionization Energy
Trends in Ionization Energy
What are the trends among
the elements for first
ionization energy, ionic size,
and electronegativity?
6.3
124. Trends in Ionization Energy
The energy required to remove an
electron from an atom is called
ionization energy.
The energy required to remove
the first electron from an atom is
called the first ionization energy.
The energy required to remove an
electron from an ion with a 1+
charge is called the second
ionization energy.
6.3
125. Trends in Ionization Energy
Group and Periodic Trends in
Ionization Energy
First ionization energy tends to
decrease from top to bottom
within a group and increase
from left to right across a
period.
6.3
129. Trends in Ionic Size
Trends in Ionic Size
During reactions between
metals and nonmetals, metal
atoms tend to lose electrons,
and nonmetal atoms tend to
gain electrons. The transfer
has a predictable effect on the
size of the ions that form.
6.3
130. Trends in Ionic Size
Cations are always smaller
than the atoms from which
they form. Anions are always
larger than the atoms from
which they form.
6.3
131. Trends in Ionic Size
Relative Sizes of Some Atoms and
Ions
6.3
132. Trends in Ionic Size
Trends in Ionic Size
6.3Sizegenerallyincreases
133. Trends in Electronegativity
Trends in Electronegativity
Electronegativity is the ability of an
atom of an element to attract
electrons when the atom is in a
compound.
In general, electronegativity values
decrease from top to bottom within a
group. For representative elements,
the values tend to increase from left
to right across a period.
6.3
138. Ions
Pyrite (FeS2), a common mineral
that emits sparks when struck
against steel, is often mistaken
for gold—hence its nickname,
“fool’s gold.” Pyrite is an example
of a crystalline solid. In this
chapter, you will learn about
crystalline solids composed of
ions that are bonded together.
But first you need to understand
how ions form from neutral
atoms.
7.1
140. Valence Electrons
Valence electrons are the
electrons in the highest
occupied energy level of an
element’s atoms.
The number of valence
electrons largely determines
the chemical properties of an
element.
7.1
141. Valence Electrons
To find the number of
valence electrons in an
atom of a representative
element, simply look at its
group number.
7.1
144. The Octet Rule
The Octet Rule
Atoms of which elements
tend to gain electrons?
Atoms of which elements
tend to lose electrons?
7.1
145. The Octet Rule
Noble gases, such as neon and argon,
are unreactive in chemical reactions. In
1916, chemist Gilbert Lewis used this
fact to explain why atoms form certain
kinds of ions and molecules.
He called his explanation the octet rule:
In forming compounds, atoms tend to
achieve the electron configuration of a
noble gas.
7.1
146. The Octet Rule
Atoms of metals tend to lose
their valence electrons, leaving
a complete octet in the next-
lowest energy level. Atoms of
some non-metals tend to gain
electrons or to share electrons
with another nonmetal to
achieve a complete octet.
7.1
148. Formation of Cations
An atom’s loss of valence
electrons produces a cation,
or a positively charged ion.
7.1
149. Formation of Cations
The most common cations are those produced
by the loss of valence electrons from metal
atoms.
You can represent the electron loss, or
ionization, of the sodium atom by drawing the
complete electron configuration of the atom and
of the ion formed.
7.1
150. Formation of Cations
The electron configuration of the sodium
ion is the same as that of a neon atom.
7.1
153. Formation of CationsA magnesium atom attains the electron configuration of
neon by losing both valence electrons. The loss of
valence electrons produces a magnesium cation with a
charge of 2+.
7.1
154. Formation of Cations
Walnuts are a good dietary
source of magnesium.
Magnesium ions (Mg2+) aid
in digestive processes.
7.1
155. Formation of Cations
Cations of Group 1A
elements always have a
charge of 1+. Cations of
group 2A elements always
have a charge of 2+.
7.1
156. Formation of Cations A copper atom can ionize to form a 1+ cation (Cu+). By losing its lone 4s
electron, copper attains a pseudo noble-gas electron configuration.
7.1
158. Formation of Anions
The gain of negatively charged
electrons by a neutral atom
produces an anion.
An anion is an atom or a group of
atoms with a negative charge.
The name of an anion typically
ends in -ide.
7.1
159. Formation of Anions
The figure shows the symbols of
anions formed by some elements
in Groups 5A, 6A, and 7A.
7.1
160. Formation of Anions
A gain of one electron gives chlorine an octet and converts a
chlorine atom into a chloride ion. It has the same electron
configuration as the noble gas argon.
7.1
161. Formation of Anions
Both a chloride ion and the argon atom have an
octet of electrons in their highest occupied energy
levels.
7.1
162. Formation of Anions
In this equation, each dot in the electron dot
structure represents an electron in the valence
shell in the electron configuration diagram.
7.1
163. Formation of Anions
The negatively
charged ions in
seawater—the
anions—are mostly
chloride ions.
7.1
164. Formation of Anions
The ions that are produced when atoms
of chlorine and other halogens gain
electrons are called halide ions.
All halogen atoms have seven valence
electrons.
All halogen atoms need to gain only one
electron to achieve the electron
configuration of a noble gas.
7.1
