Download 06 Electronic Structure

Lecture Presentation
Chapter 6
Electronic
Structure of Atoms
© 2015 Pearson Education, Inc.
James F. Kirby
Quinnipiac University
Hamden, CT
Electronic
Structure
of Atoms
LESSON 1
6-1 The Wave Nature of
Light
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Electronic Structure
• This chapter is all about electronic
structure—the arrangement and
energy of electrons.
• It may seem odd to start by talking
about waves. However, extremely small
particles have properties that can only
be explained in this manner!
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Waves
• To understand the electronic structure of
atoms, one must understand the nature of
electromagnetic radiation.
• The distance between corresponding points
on adjacent waves is the wavelength ().
© 2015 Pearson Education, Inc.
Electronic
Structure
of Atoms
Properties of Waves
Wavelength () is the distance between identical points on
successive waves (meters).
Amplitude is the vertical distance from the midline of a
wave to the peak or trough.
Properties of Waves
Frequency (n) is the number of waves that pass through a
particular point in 1 second (Hz = 1 cycle/s).
The speed (u) of the wave =  x n
Waves
• The number of waves
passing a given point per unit
of time is the frequency (n).
• For waves traveling at the
same velocity, the longer the
wavelength, the smaller the
frequency.
• If the time associated with
the lines to the left is one
second, then the frequencies
would be 2 s–1 and 4 s–1,
Electronic
respectively.
Structure
of Atoms
© 2015 Pearson Education, Inc.
Maxwell (1873), proposed that visible light consists of
electromagnetic waves.
Electromagnetic
radiation is the emission
and transmission of energy
in the form of
electromagnetic waves.
Speed of light (c) in vacuum = 3.00 x 108 m/s
All electromagnetic radiation
xn=c
Red Light (~700nm) has a longer wavelength than
violet (~400nm)
Electromagnetic Radiation
• All electromagnetic radiation travels at the same
velocity: The speed of light (c) is 3.00  108 m/s.
c = n
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Light and the Electromagnetic
Spectrum
Wavelength x Frequency = Speed

m
x
n
1
s
=
c
m
s
c is defined to be the rate of travel of
all electromagnetic energy in a
vacuum and is a constant value—
speed of light.
m
8
c = 3.00 x 10
s
Light and the Electromagnetic
Spectrum
The light blue glow given off by mercury
streetlamps has a frequency of 6.88 x 1014 s-1 (or,
Hz). What is the wavelength in nanometers?
=
c
n
m
8
3.00 x 10
s
1 x 109 nm
m
=
6.88 x
= 436 nm
1014
1
s
A photon has a frequency of 6.0 x 104 Hz. Convert
this frequency into wavelength (nm). Does this frequency
fall in the visible region?
Hwk: page 249 - 254:
2, GF6.3, GF6.4, 13, 14, 17, 18, 19, 21,
22, 84
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
LESSON 2
6-2 Quantized Energy and
Photons
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Nature of Energy
The wave nature of light
does not explain how
an object can glow
when its temperature
increases.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Nature of Energy—Quanta
Max Planck
explained it by
assuming that
energy comes
in packets
called quanta
(singular:
quantum).
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Photoelectric Effect
• Einstein used quanta to explain
the photoelectric effect.
• Each metal has a different
energy at which it ejects
electrons. At lower energy,
electrons are not emitted.
• He concluded that energy is
proportional to frequency:
E = hn
where h is Planck’s constant,
6.626  10−34 J∙s.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Particle Nature of Light
Investigations carried out by Max Planck (blackbody
radiation) and Albert Einstein (photoelectric effect)
discredited the notion that all properties of light could be
explained in terms of its wave nature.
Today we considered light to be generated as a
stream of particles called photons, whose energy is
given by the Einstein equation
E=hν=hc/λ
h = 6.626 x 10-34 J·s
Quantum theory is used to explain any interaction of energy
and matter. Light is made of packets of energy called
photons.
When copper is bombarded with high-energy electrons,
X rays are emitted. Calculate the energy (in joules)
associated with the photons if the wavelength of the X
rays is 0.154 nm.
Calculate the energy, in joules, of a photon emitted
by an oxygen atom with 557.7 nm.
Calculate the energy, in kilojoules, of a mole of
such photons.
Hwk: page 249 - 254:
23, 25, 27, 29, 33
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
LESSON 3
6-3 Line Spectra and the
Bohr Model
Electronic
Structure
of Atoms
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Atomic Emissions
Another mystery in the early twentieth century
involved the emission spectra observed from
energy emitted by atoms and molecules.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Hydrogen Spectrum
• Johann Balmer (1885) discovered a
simple formula relating the four lines to
integers.
• Johannes Rydberg advanced this
formula.
• Neils Bohr explained why this
mathematical relationship works.
© 2015 Pearson Education, Inc.
Electronic
Structure
of Atoms
Continuous vs. Line Spectra
• For atoms and molecules,
one does not observe a
continuous spectrum
(the “rainbow”), as one
gets from a white light
source.
• Only a line spectrum of
discrete wavelengths is
observed. Each element
has a unique line
spectrum.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Line Emission Spectrum of Hydrogen Atoms
The Bohr Model
•
Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this way:
1. Electrons in an atom can
only occupy certain orbits
(corresponding to certain
energies).
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Bohr Model
2. Electrons in permitted orbits
have specific, “allowed”
energies; these energies will not
be radiated from the atom.
3. Energy is only absorbed or
emitted in such a way as to
move an electron from one
“allowed” energy state to
another; the energy is defined by
E = hn
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Bohr Model
The energy absorbed or emitted
from the process of electron
promotion or demotion can be
calculated by the equation
E = −hcRH (
1
1
–
n f2
ni2
)
where RH is the Rydberg
constant, 1.097  107 m−1, and ni
and nf are the initial and final
energy levels of the electron.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Limitations of the Bohr Model
• It only works for hydrogen!
• Classical physics would result in an
electron falling into the positively
charged nucleus. Bohr simply assumed
it would not!
• Circular motion is not wave-like in
nature.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Important Ideas from the
Bohr Model
Points that are incorporated into the
current atomic model include the
following:
1) Electrons exist only in certain discrete
energy levels.
2) Energy is involved in the transition of
an electron from one level to another.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Bohr’s Model of
the Atom (1913)
1. e- can only have specific
(quantized) energy
values
2. light is emitted as emoves from one energy
level to another
En = -RH (
1
n2
)
n (principal quantum number) = 1,2,3,…
RH (Rydberg constant) = 2.18 x 10-18J
E = hn
E = hn
ni = 3
ni = 3
ni = 2
nf = 2
nnf f==11
Ephoton = E = Ef - Ei
1
Ef = -RH ( 2
nf
1
Ei = -RH ( 2
ni
1
E = RH( 2
ni
)
)
1
n2f
)
Calculate the wavelength (in nm) of a photon
emitted by a hydrogen atom when its electron
drops from the n = 5 state to the n = 3 state.
LESSON 4
6-4 The Wave Behavior of
Matter
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Wave Nature of Matter
• Louis de Broglie theorized
that if light can have material
properties, matter should
exhibit wave properties.
• He demonstrated that the
relationship between mass
and wavelength was
The wave nature of light
is used to produce this
electron micrograph.
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h
 = mv
Electronic
Structure
of Atoms
Why is e- energy quantized?
De Broglie (1924) reasoned
that e- is both particle and
wave.
2pr = n
 = h/mu
u = velocity of em = mass of e-
What is the de Broglie wavelength (in nm)
associated with a 2.5 g Ping-Pong ball
traveling at 15.6 m/s?
The Uncertainty Principle
Heisenberg showed
that the more precisely
the momentum of a
particle is known, the
less precisely is its
position is known:
h
(x) (mv) 
4p
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Hwk: page 249 - 254:
37, 39, 47, 49
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
LESSON 5
6-5 Quantum Mechanics
and Atomic Orbitals
Electronic
Structure
of Atoms
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Quantum Mechanics
• Erwin Schrödinger
developed a mathematical
treatment into which both
the wave and particle
nature of matter could be
incorporated.
• This is known as
quantum mechanics.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Quantum Mechanics
• The solution of Schrödinger’s
wave equation is designated with
a lowercase Greek psi ().
• The square of the wave equation,
2, gives the electron density, or
probability of where an electron is
likely to be at any given time.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Schrodinger Wave Equation
In 1926 Schrodinger wrote an equation that
described both the particle and wave nature of the eWave function (Y) describes:
1. energy of e- with a given Y
2. probability of finding e- in a volume of space
Schrodingers equation can only be solved exactly
for the hydrogen atom. Must approximate its
solution for multi-electron systems.
Quantum Numbers
• Solving the wave equation gives a set of
wave functions, or orbitals, and their
corresponding energies.
• Each orbital describes a spatial
distribution of electron density.
• An orbital is described by a set of three
quantum numbers.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Principal Quantum Number (n)
• The principal quantum number, n,
describes the energy level on which the
orbital resides.
• The values of n are integers ≥ 1.
• These correspond to the values in the
Bohr model.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
principal quantum number n
n = 1, 2, 3, 4, ….
distance of e- from the nucleus
2n2
n=1
n=2
n=3
Where 90% of the
e- density is found
for the 1s orbital
e- density (1s orbital) falls off rapidly
as distance from nucleus increases
Angular Momentum Quantum
Number (l)
• This quantum number defines the shape of
the orbital.
• Allowed values of l are integers ranging
from 0 to n − 1.
• We use letter designations to communicate
the different values of l and, therefore, the
shapes and types of orbitals.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Angular Momentum Quantum
Number (l)
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
l=0
l=1
l=2
l=3
s orbital
p orbital
d orbital
f orbital
Shape of the “volume” of space that the e- occupies
l = 0 (s orbitals)
l = 1 (p orbitals)
l = 2 (d orbitals)
Magnetic Quantum Number (ml)
• The magnetic quantum number describes the
three-dimensional orientation of the orbital.
• Allowed values of ml are integers ranging
from −l to l:
−l ≤ ml ≤ l
• Therefore, on any given energy level, there
can be up to 1 s orbital, 3 p orbitals, 5 d
orbitals, 7 f orbitals, and so forth.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Magnetic Quantum Number (ml)
• Orbitals with the same value of n form an electron
shell.
• Different orbital types within a shell are subshells.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
magnetic quantum number ml
for a given value of l
ml = -l, …., 0, …. +l
if l = 1 (p orbital), ml = -1, 0, or 1
if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2
orientation of the orbital in space
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
spin quantum number ms
ms = +½ or -½
ms = +½
ms = -½
LESSON 6
6-6 Representations of
Orbitals
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
s Orbitals
• The value of l for s orbitals is 0.
• They are spherical in shape.
• The radius of the sphere increases with the
value of n.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
s Orbitals
• For an ns orbital, the
number of peaks is n.
• For an ns orbital, the
number of nodes (where
there is zero probability
of finding an electron) is
n – 1.
• As n increases, the
electron density is more
spread out and there is
a greater probability of
finding an electron
further from the nucleus.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
p Orbitals
• The value of l for p orbitals is 1.
• They have two lobes with a node between them.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
d Orbitals
• The value of l for a
d orbital is 2.
• Four of the five d
orbitals have four
lobes; the other
resembles a p
orbital with a
doughnut around
the center.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
ml = -1
ml = -2
ml = 0
ml = -1
ml = 0
ml = 1
ml = 1
ml = 2
f Orbitals
• Very complicated shapes (not shown
in text)
• Seven equivalent orbitals in a sublevel
• l=3
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Hwk: page 249 - 254:
55, 57, 59, 61
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
LESSON 7
6-7 Many-Electron Atoms
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Energies of Orbitals—Hydrogen
• For a one-electron
hydrogen atom,
orbitals on the same
energy level have
the same energy.
• Chemists call them
degenerate orbitals.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Energies of Orbitals—
Many-electron Atoms
• As the number of electrons
increases, so does the
repulsion between them.
• Therefore, in atoms with
more than one electron, not
all orbitals on the same
energy level are degenerate.
• Orbital sets in the same
sublevel are still degenerate.
• Energy levels start to overlap
in energy (e.g., 4s is lower
Electronic
in energy than 3d.)
Structure
of Atoms
© 2015 Pearson Education, Inc.
Spin Quantum Number, ms
• In the 1920s, it was discovered that
two electrons in the same orbital do
not have exactly the same energy.
• The “spin” of an electron describes
its magnetic field, which affects its
energy.
• This led to the spin quantum
number, ms.
• The spin quantum number has only
two allowed values, +½ and –½.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Pauli Exclusion Principle
• No two electrons in the same atom can
have exactly the same energy.
• Therefore, no two electrons in the same
atom can have identical sets of quantum
numbers.
• This means that every electron in an atom
must differ by at least one of the four
quantum number values: n, l, ml, and ms.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Schrodinger Wave Equation
Y = fn(n, l, ml, ms)
Existence (and energy) of electron in atom is described
by its unique wave function Y.
Pauli exclusion principle - no two electrons in an atom
can have the same four quantum numbers.
Each seat is uniquely identified (E, R12, S8)
Each seat can hold only one individual at a
time
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
LESSON 8
6-8 Electron Configurations
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Electron Configurations
• The way electrons are distributed in an
5 atom is called its electron configuration.
• The most stable organization is the lowest
possible energy, called the ground state.
• Each component consists of
– a number denoting the energy level;
4p
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Electron Configurations
5
4p
• The way electrons are distributed in an
atom is called its electron configuration.
• The most stable organization is the lowest
possible energy, called the ground state.
• Each component consists of
– a number denoting the energy level;
– a letter denoting the type of orbital;
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Electron Configurations
5
4p
• The way electrons are distributed in an
atom is called its electron configuration.
• The most stable organization is the lowest
possible energy, called the ground state.
• Each component consists of
– a number denoting the energy level;
– a letter denoting the type of orbital;
– a superscript denoting the number of
electrons in those orbitals.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Orbital Diagrams
• Each box in the
diagram represents
one orbital.
• Half-arrows represent
the electrons.
• The direction of the
arrow represents the
relative spin of the
electron.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
“Fill up” electrons in lowest energy orbitals (Aufbau principle)
??
B 1s22s22p1
B 5 electrons
Be 1s22s2
Be 4 electrons
Li 1s22s1
Li 3 electrons
He 1s2
He 2 electrons
H 1s1
H 1 electron
Order of orbitals (filling) in multi-electron atom
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s
Outermost subshell being filled with electrons
The most stable arrangement of electrons
in subshells is the one with the greatest
number of parallel spins (Hunds rule).
C 6 electrons
C 1s22s22p2
N 7 electrons
N 1s22s22p3
O 8 electrons
O 1s22s22p4
F 9 electrons
F 1s22s22p5
Ne 10 electrons
Ne 1s22s22p6
Hund’s Rule
“For degenerate
orbitals, the
lowest energy is
attained when
the number of
electrons with
the same spin is
maximized.”
 This means that, for a set of orbitals in the same
sublevel, there must be one electron in each orbital
before pairing and the electrons have the same spin,
as much as possible.
© 2015 Pearson Education, Inc.
Electronic
Structure
of Atoms
What is the electron configuration of Mg?
What are the possible quantum numbers for the
last (outermost) electron in Cl?
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Condensed Electron Configurations
• Elements in the same group of the
periodic table have the same number
of electrons in the outer most shell.
These are the valence electrons.
• The filled inner shell electrons are
called core electrons. These include
completely filled d or f sublevels.
• We write a shortened version of an
electron configuration using brackets
around a noble gas symbol and listing
only valence electrons.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
ns2np6
ns2np5
ns2np4
ns2np3
ns2np2
ns2np1
d10
d5
d1
ns2
ns1
Ground State Electron Configurations of the Elements
4f
5f
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Electronic
Structure
of Atoms
Do examples of electron configuration, orbital
notation, and shorthand method.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Paramagnetic
unpaired electrons
2p
Diamagnetic
all electrons paired
2p
Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
Atoms gain electrons
so that anion has a
noble-gas outer
electron
configuration.
Atoms lose electrons so
that cation has a noblegas outer electron
configuration.
H 1s1
H- 1s2 or [He]
F 1s22s22p5
F- 1s22s22p6 or
[Ne]
O2- 1s22s22p6 or [Ne]
O 1s22s22p4
N 1s22s22p3 N3- 1s22s22p6 or [Ne]
-1
-2
-3
+3
+2
+1
Cations and Anions Of Representative Elements
Na+: [Ne]
Al3+: [Ne]
O2-: 1s22s22p6 or [Ne]
F-: 1s22s22p6 or
[Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
Electron Configurations of Cations of Transition Meta
When a cation is formed from an atom of a transition
metal, electrons are always removed first from the ns
orbital and then from the (n – 1)d orbitals.
Fe:
[Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn:
[Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
LESSON 9
6-9 Electron Configurations
and the Periodic Table
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Periodic Table
• We fill orbitals in increasing order of energy.
• Different blocks on the periodic table correspond to
different types of orbitals: s = blue, p = pink (s and p
are representative elements); d = orange (transition
elements); f = tan (lanthanides and actinides, or
inner transition elements)
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Some Anomalies
 Some irregularities
occur when there
are enough
electrons to half-fill
s and d orbitals on
a given row.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Chromium as an Anomaly
• For instance, the electron configuration
for chromium is
[Ar] 4s1 3d5
rather than the expected
[Ar] 4s2 3d4.
• This occurs because the 4s and 3d
orbitals are very close in energy.
• These anomalies occur in f-block atoms
Electronic
with f and d orbitals, as well.
Structure
of Atoms
© 2015 Pearson Education, Inc.
Hwk: page 249 - 254:
71, 72, 74, 75, 77, 78, 79, 80
Quiz to follow
Electronic
Structure
of Atoms
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Review Questions
Chapter 6
Stoichiometry
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The wavelength of a photon
multiplied by its frequency
equals
a.
b.
c.
d.
c, the speed of light.
h, Planck’s constant.
Avogadro’s Number.
4.184.
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The wavelength of a photon
multiplied by its frequency
equals
a.
b.
c.
d.
c, the speed of light.
h, Planck’s constant.
Avogadro’s Number.
4.184.
© 2015 Pearson Education, Inc.
The energy of a photon
divided by its frequency
equals
a.
b.
c.
d.
c, the speed of light.
h, Planck’s constant.
Avogadro’s Number.
4.184.
© 2015 Pearson Education, Inc.
The energy of a photon
divided by its frequency
equals
a.
b.
c.
d.
c, the speed of light.
h, Planck’s constant.
Avogadro’s Number.
4.184.
© 2015 Pearson Education, Inc.
The “rainbow of colors”
produced by sunlight striking a
prism is called
a.
b.
c.
d.
a continuous spectrum.
monochromatic light.
a line spectrum.
a Balmer series.
© 2015 Pearson Education, Inc.
The “rainbow of colors”
produced by sunlight striking a
prism is called
a.
b.
c.
d.
a continuous spectrum.
monochromatic light.
a line spectrum.
a Balmer series.
© 2015 Pearson Education, Inc.
The lowest energy state of a
hydrogen atom is called its
_______ state.
a.
b.
c.
d.
bottom
ground
fundamental
original
© 2015 Pearson Education, Inc.
The lowest energy state of a
hydrogen atom is called its
_______ state.
a.
b.
c.
d.
bottom
ground
fundamental
original
© 2015 Pearson Education, Inc.
When an electron moves from
the n = 3 orbit to the n = 2
orbit of a hydrogen atom, what
wavelength of light is emitted?
a.
b.
c.
d.
410 nm
434 nm
486 nm
656 nm
© 2015 Pearson Education, Inc.
When an electron moves from
the n = 3 orbit to the n = 2
orbit of a hydrogen atom, what
wavelength of light is emitted?
a.
b.
c.
d.
410 nm
434 nm
486 nm
656 nm
© 2015 Pearson Education, Inc.
When an electron moves from
the n = 4 orbit to the n = 2
orbit of a hydrogen atom, what
wavelength of light is emitted?
a.
b.
c.
d.
410 nm
434 nm
486 nm
656 nm
© 2015 Pearson Education, Inc.
When an electron moves from
the n = 4 orbit to the n = 2
orbit of a hydrogen atom, what
wavelength of light is emitted?
a.
b.
c.
d.
410 nm
434 nm
486 nm
656 nm
© 2015 Pearson Education, Inc.
“It is impossible to
simultaneously know both the
position and the momentum of
an electron in an atom” is
a.
b.
c.
d.
Hund’s rule.
deBroglie’s hypothesis.
Pauli’s exclusion principle.
Heisenberg’s uncertainty
principle.
© 2015 Pearson Education, Inc.
“It is impossible to
simultaneously know both the
position and the momentum of
an electron in an atom” is
a.
b.
c.
d.
Hund’s rule.
deBroglie’s hypothesis.
Pauli’s exclusion principle.
Heisenberg’s uncertainty
principle.
© 2015 Pearson Education, Inc.
“No two electrons in an atom
may have the same values for
all four quantum numbers” is
a.
b.
c.
d.
Hund’s rule.
deBroglie’s hypothesis.
Pauli’s exclusion principle.
Heisenberg’s uncertainty principle.
© 2015 Pearson Education, Inc.
“No two electrons in an atom
may have the same values for
all four quantum numbers” is
a.
b.
c.
d.
Hund’s rule.
deBroglie’s hypothesis.
Pauli’s exclusion principle.
Heisenberg’s uncertainty principle.
© 2015 Pearson Education, Inc.
The line spectrum of hydrogen
includes lines at 447, 502,
587, and 668 nm. The line at
___ nm represents the most
energetic transition.
a. 447
c. 587
© 2015 Pearson Education, Inc.
b. 502
d. 668
The line spectrum of hydrogen
includes lines at 447, 502,
587, and 668 nm. The line at
___ nm represents the most
energetic transition.
a. 447
c. 587
© 2015 Pearson Education, Inc.
b. 502
d. 668
Which sequence lists types of
electromagnetic energy in
order of increasing energy?
a.
b.
c.
d.
microwave, IR, visible, UV
IR, microwave, UV, visible
UV, visible, IR, microwave
visible, UV, microwave, IR
© 2015 Pearson Education, Inc.
Which sequence lists types of
electromagnetic energy in
order of increasing energy?
a.
b.
c.
d.
microwave, IR, visible, UV
IR, microwave, UV, visible
UV, visible, IR, microwave
visible, UV, microwave, IR
© 2015 Pearson Education, Inc.
n and l are the principal and
angular momentum quantum
numbers. When n = 3, the
allowed values of l are
a.
b.
c.
d.
1, 2, and 3.
1 and 2.
0, 1, 2, and 3.
0, 1, and 2.
© 2015 Pearson Education, Inc.
n and l are the principal and
angular momentum quantum
numbers. When n = 3, the
allowed values of l are
a.
b.
c.
d.
1, 2, and 3.
1 and 2.
0, 1, 2, and 3.
0, 1, and 2.
© 2015 Pearson Education, Inc.
Which set of quantum
numbers correctly describes
an electron in the outermost
orbital of a sulfur atom?
a.
b.
c.
d.
n = 3, l = 2, ml = –2
n = 2, l = 1, ml = –1
n = 2, l = 0, ml = 0
n = 3, l = 1, ml = –1
© 2015 Pearson Education, Inc.
Which set of quantum
numbers correctly describes
an electron in the outermost
orbital of a sulfur atom?
a.
b.
c.
d.
n = 3, l = 2, ml = –2
n = 2, l = 1, ml = –1
n = 2, l = 0, ml = 0
n = 3, l = 1, ml = –1
© 2015 Pearson Education, Inc.
Which set of values is not
correct for an electron
occupying a 4d orbital?
a.
b.
c.
d.
n = 4, l = 2, ml = 0
n = 4, l = 2, ml = –1/2
n = 3, l = 4, ml = 1
n = 3, l = 1, ml = 1
© 2015 Pearson Education, Inc.
Which set of values is not
correct for an electron
occupying a 4d orbital?
a.
b.
c.
d.
n = 4, l = 2, ml = 0
n = 4, l = 2, ml = –1/2
n = 3, l = 4, ml = 1
n = 3, l = 1, ml = 1
© 2015 Pearson Education, Inc.
The only allowed values for
the spin magnetic quantum
number are
a.
b.
c.
d.
0 and 1.
0 and +1/2.
+1/2 and 1.
+1/2 and –1/2.
© 2015 Pearson Education, Inc.
The only allowed values for
the spin magnetic quantum
number are
a.
b.
c.
d.
0 and 1.
0 and +1/2.
+1/2 and 1.
+1/2 and –1/2.
© 2015 Pearson Education, Inc.
s orbitals are shaped like
a.
b.
c.
d.
four-leaf clovers.
dumbbells.
spheres.
triangles.
© 2015 Pearson Education, Inc.
s orbitals are shaped like
a.
b.
c.
d.
four-leaf clovers.
dumbbells.
spheres.
triangles.
© 2015 Pearson Education, Inc.
p orbitals are shaped like
a.
b.
c.
d.
four-leaf clovers.
dumbbells.
spheres.
triangles.
© 2015 Pearson Education, Inc.
p orbitals are shaped like
a.
b.
c.
d.
four-leaf clovers.
dumbbells.
spheres.
triangles.
© 2015 Pearson Education, Inc.
At a node, the probability of
finding an electron is ___ %.
a.
b.
c.
d.
0
1
50
100
© 2015 Pearson Education, Inc.
At a node, the probability of
finding an electron is ___ %.
a.
b.
c.
d.
0
1
50
100
© 2015 Pearson Education, Inc.
The electron configuration of a
carbon atom is
a.
b.
c.
d.
[He]2s22p6.
[He]2s22p4.
[He]2s22p2.
[He]2s2.
© 2015 Pearson Education, Inc.
The electron configuration of a
carbon atom is
a.
b.
c.
d.
[He]2s22p6.
[He]2s22p4.
[He]2s22p2.
[He]2s2.
© 2015 Pearson Education, Inc.
The electron configuration of a
germanium atom is
a.
b.
c.
d.
[Ar]4s24p2.
[Ar]4s23d104p2.
[Kr]4s23d104p2.
[Kr]4s23d104p2.
© 2015 Pearson Education, Inc.
The electron configuration of a
germanium atom is
a.
b.
c.
d.
[Ar]4s24p2.
[Ar]4s23d104p2.
[Kr]4s23d104p2.
[Kr]4s23d104p2.
© 2015 Pearson Education, Inc.
The electron configuration of a
copper atom is
a.
b.
c.
d.
[Ar]4s23d9.
[Ar]4s13d10.
[Ar]4s23d10.
[Ar]4s23d7.
© 2015 Pearson Education, Inc.
The electron configuration of a
copper atom is
a.
b.
c.
d.
[Ar]4s23d9.
[Ar]4s13d10.
[Ar]4s23d10.
[Ar]4s23d7.
© 2015 Pearson Education, Inc.
The valence electron
configuration of elements in
column 6A(16) of the Periodic
Table is
a.
b.
c.
d.
np6.
ns0np6.
ns2np4.
impossible to predict because
each element is unique.
© 2015 Pearson Education, Inc.
The valence electron
configuration of elements in
column 6A(16) of the Periodic
Table is
a.
b.
c.
d.
np6.
ns0np6.
ns2np4.
impossible to predict because
each element is unique.
© 2015 Pearson Education, Inc.