Download Chapter 6

Lecture Presentation
Chapter 6
Electronic
Structure of Atoms
© 2015 Pearson Education, Inc.
James F. Kirby
Quinnipiac University
Hamden, CT
Electronic
Structure
of Atoms
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
Waves
• The number of waves
passing a given point per unit
of time is the frequency ().
• 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.
Electromagnetic Radiation
• All electromagnetic radiation travels at the same
velocity: The speed of light (c) is 3.00  108 m/s.
c = 
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Sample Exercise 6.1 Concepts of Wavelength and
Frequency
Two electromagnetic waves are represented in the margin. (a) Which wave has
the higher frequency? (b) If one wave represents visible light and the other
represents infrared radiation, which wave is which?
Solution
(a) Wave 1 has a longer wavelength (greater distance between peaks). The longer the wavelength,
the lower the frequency (ν = c ⁄ λ). Thus, Wave 1 has the lower frequency, and Wave 2 has the
higher frequency.
(b) The electromagnetic spectrum (Figure 6.4) indicates that infrared radiation has a longer
wavelength than visible light. Thus, Wave 1 would be the infrared radiation.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Sample Exercise 6.2 Calculating Frequency from
Wavelength
The yellow light given off by a sodium vapor lamp used for public lighting has a wavelength of
589 nm. What is the frequency of this radiation?
Solution
Analyze We are given the wavelength, λ, of the radiation and asked to calculate its frequency, ν.
Plan The relationship between the wavelength and the frequency is given by Equation 6.1. We can
solve for ν and use the values of λ and c to obtain a numerical answer. (The speed of light, c, is 3.00
× 108 m ⁄ s to three significant figures.)
Solve Solving Equation 6.1 for frequency gives (ν = c ⁄ λ). When we insert the values for c and λ, we
note that the units of length in these two quantities are different. We can convert the Wavelength
from nanometers to meters, so the units cancel:
Check The high frequency is reasonable because of the short wavelength. The units are proper
because frequency has units of “per second,” or s–1.
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 = h
where h is Planck’s constant,
6.626  10−34 J∙s.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Sample Exercise 6.3 Energy of a Photon
Calculate the energy of one photon of yellow light that has a wavelength of 589 nm.
Solution
Analyze Our task is to calculate the energy, E, of a photon, given λ = 589 nm.
Plan We can use Equation 6.1 to convert the wavelength to frequency:
ν=c⁄λ
We can then use Equation 6.3 to calculate energy:
E = hν
Solve The frequency, ν, is calculated from the given wavelength, as shown in Sample Exercise 6.2:
ν = c ⁄ λ = 5.09 × 1014 s–1
The value of Planck constant, h, is given both in the text and in the table of physical constants on
the inside back cover of the text, and so we can easily calculate E:
E = (6.626 × 10–34 J-s)(5.09 × 1014 s–1) = 3.37 × 10–19 J
Comment If one photon of radiant energy supplies 3.37 × 10–19 J, we calculate that one mole of
these photons will supply:
(6.02 × 1023 photons ⁄ mol)(3.37 × 10–19 J ⁄ photon) = 2.03 × 105 J ⁄ mol
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
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.
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.
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
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 = h
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.
Sample Exercise 6.4 Electronic Transitions in the
Hydrogen Atom
Using Figure 6.12, predict which of these electronic
transitions produces the spectral line having the
longest wavelength: n = 2 to n = 1, n = 3 to n = 2, or n
= 4 to n = 3.
Solution
The wavelength increases as frequency decreases (λ = c ⁄ ν). Hence, the longest wavelength will
be associated with the lowest frequency. According to Planck’s equation, E = hv, the lowest
frequency is associated with the lowest energy. In Figure 6.12 the shortest vertical line
represents the smallest energy change. Thus, the n = 4 to n = 3 transition produces the longest
Electronic
wavelength (lowest frequency) line.
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.
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.
© 2015 Pearson Education, Inc.
h
 = mv
Electronic
Structure
of Atoms
Sample Exercise 6.5 Matter Waves
What is the wavelength of an electron moving with a speed of 5.97 × 106 m ⁄ s? The mass
of the electron is 9.11 × 10–31 kg.
Solution
Analyze We are given the mass, m, and velocity, v, of the electron, and we must calculate its de
Broglie wavelength, λ.
Plan The wavelength of a moving particle is given by Equation 6.8, so λ is calculated by inserting the
known quantities h, m, and v. In doing so, however, we must pay attention to units.
Solve Using the value of Planck constant:
h = 6.626 × 10–34 J-s
we have the following:
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
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) 
4
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
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.
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.
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.
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.
Sample Exercise 6.6 Subshells of the Hydrogen Atom
(a) Without referring to Table 6.2, predict the number of subshells in the fourth shell, that is, for n
= 4. (b) Give the label for each of these subshells. (c) How many orbitals are in each of these
subshells?
Solution
Analyze and Plan We are given the value of the principal quantum number, n. We need to determine
the allowed
values of l and ml for this given value of n and then count the number of orbitals in each subshell.
Solve There are four subshells in the fourth shell, corresponding to the four possible values of l (0,
1, 2, and 3).
These subshells are labeled 4s, 4p, 4d, and 4f. The number given in the designation of a
subshell is the principal quantum number, n; the letter designates the value of the angular
momentum quantum number, l: for l = 0, s; for l = 1, p; for l = 2, d; for l = 3, f.
There is one 4s orbital (when l = 0, there is only one possible value of ml: 0). There are three 4p
orbitals (when
l = 1, there are three possible values of ml: 1, 0, −1). There are five 4d orbitals (when l = 2, there are
five allowed values of ml: 2, 1, 0, −1, −2). There are seven 4f orbitals (when l = 3, there are seven
permitted values of ml: 3, 2, 1, 0, −1, −2, −3).
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.
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.
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.
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.
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
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.
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.