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Lecture Presentation
Chapter 6
Electronic
Structure of Atoms
© 2015 Pearson Education, Inc.
James F. Kirby
Quinnipiac University
Hamden, CT
Electronic
Structure
of Atoms
Introduction to chapter 6 -- Electronic Structure
• This chapter is all about electronic structure—the arrangement and
energy of electrons. The arrangement of electrons is dependent on
the electromagnetic forces between the electrons and the nucleus of
the atom, which is explained using Coulomb’s Law
• Coulomb’s law: the force between two charged particles is
proportional to the magnitude of each of the two charges (q1 and q2),
and inversely proportional to the square of the distance, r, between
them. If the two charges are of opposite sign, the force between them
is attractive; if they are of the same sign, the force is repulsive.
• What this has to do with an atom:
 The closer an electron is to the nucleus, the more attraction it has to the
nucleus
 The closer electrons are to each other, the more they repel one another
• It may seem odd to start by talking about waves. However, extremely
small particles have properties that can only be explained in this Electronic
manner!
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.1 – The Wave Nature of Light
• Electromagnetic Radiation = radiant energy that
exhibits wavelike behavior and travels through
space at the speed of light in a vacuum (3.00x108
m/s). Doesn’t require a medium for traveling. Ex.
The light from the sun, microwaves, x-rays
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
• Properties of electromagnetic radiation:
a. wavelength (  ) = the distance between 2 consecutive
crests or troughs of a wave. Common units are nm and
m.
b. frequency (  ) = the # of waves that pass a given point in
space per second. SI unit = s-1 (hertz or Hz)
for electromagnetic radiation, c =  x 
(c=the speed of light (3 x 108m/s), λ is in m)
So, λ and ν are inversely related
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.2 - Quantized Energy and Photons (the
particle nature of light)
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.
Max Planck
explained it by
assuming that
energy comes
in packets
called quanta
(singular:
quantum).
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
• The wave nature of light also does not explain the
following:
• Photoelectric effect = the ability of certain materials to
emit electrons from their surfaces when struck by
electromagnetic radiation of a minimum frequency. This
demonstrates a particle-like quality of light. The particles
are called photons. These photons transfer their
energies to electrons, which are then ejected.
Ephoton = h , where h=Planck’s
constant (6.626 x 10-34 J s), and
 is frequency.
• energy can be gained or lost only
in whole-number multiples of h.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.3 - Line Spectra and the Bohr Model
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.
• This cannot be explained by
the wave theory
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
The Bohr Model
•
Niels Bohr explained the line specta in
this way:
1. Electrons in an atom can
only occupy certain orbits
(corresponding to certain energies).
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
• When electrons go from a lower to a
higher energy level, they absorb
energy
• When electrons go from a higher to
a lower energy level, they emit the
excess quantum of energy as a
photon of light with an amount of
energy equal to the difference in
energy between the higher and
lower energy level.
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.
6.4 - The Wave Behavior of Matter
• Louis de Broglie theorized that if light can have
material properties, matter should exhibit wave
properties.
• He said that a moving electron has a wavelength
(very small)
• Heisenberg (with his Uncertainty Principle) showed
that the more precisely the momentum of a particle is
known, the less precisely is its position is known
• The work of these 2 scientists led to looking at the
location of electrons in terms of probabilities instead
of exact locations and paths, and describing electron
behavior in terms of waves
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.5 - Quantum Mechanics
• Erwin Schrödinger developed a
mathematical equation into which both
the wave and particle nature of matter
could be incorporated.
• This approach to electronic structure of
an atom is known as quantum
mechanics.
• The solutions of his wave equation are
designated with a lowercase Greek psi
() (called a wave function).
• The square of the wave function, 2,
gives the electron density, or
probability of where an electron is
likely to be at any given time.
© 2015 Pearson Education, Inc.
Electronic
Structure
of Atoms
• Atomic Orbitals – mathematically derived, they allow us
to visualize a 3-dimensional region in which there’s a
significant probability of finding an electron
• n = the main energy level occupied by the electron
(called the Principal Quantum Number). As n, the eenergy and distance from the nucleus .
Electronic
• n2 = the total # of orbitals in a main energy level.
Structure
of Atoms
© 2015 Pearson Education, Inc.
Sublevels – in each main energy level, there are n
sublevels. Each sublevel has a different orbital shape
and energy. The different shapes (from lowest energy to
highest) are:
s – spherical (s sublevels consist of 1orbital)
p – dumb-bell shaped (p sublevels consist of 3
orbitals)
d- complex shape (consists of 5 orbitals)
f- complex shape (consists of 7 orbitals)
orbital shapes
Each atomic orbital is represented by the principal
quantum number followed by the letter of the sublevel.
Ex. 1s, 2p
• Each orbital holds 2 electrons.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.6 – Representation of Orbitals
s Orbitals
• They are spherical in shape.
• The radius of the sphere increases with the
value of n.
• s orbitals are found in every main energy
level
© 2015 Pearson Education, Inc.
Electronic
Structure
of Atoms
p Orbitals
• They have two lobes with a node between them.
• p orbitals begin in the 2nd energy level
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
d Orbitals
• Four of the five d
orbitals have four
lobes; the other
resembles a p
orbital with a
doughnut around
the center.
• d orbitals begin in
the 3rd energy
level
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
f Orbitals
• Very complicated shapes (not shown
in text)
• Seven equivalent orbitals in a sublevel
• f orbitals begin in the 4th energy level
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.7 – Many-Electron Atoms
• 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
• Orbital energies are lower (more
negative) than in hydrogen (because of
the increased nuclear attractive force)
• 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
in energy than 3d.)
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
6.8 - Electron Configurations
• 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 of the electron configuration consists of
1) a number denoting the energy level
2) a letter denoting the type of orbital
3) a superscript denoting the number of electrons in
those orbitals
Example: 4p5 means 5 electrons in the p sublevel of
Electronic
energy level 4
Structure
of Atoms
© 2015 Pearson Education, Inc.
• The Rules for Electron
Configurations
• 1. Aufbau principle:
Electrons occupy
orbitals of lowest
energy first
• Exceptions: One
reason for exceptions
is that filled and halffilled orbitals are more
stable than partially
filled.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
• 2. Pauli exclusion principle: An orbital only holds 2 electrons, and
the electrons in the same orbital have opposite spins
• 3. Hund’s rule: When entering orbitals of equal energy, the espread out when possible. Electrons in half-filled orbitals have the
same spins.
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.
Types of Electron Configurations:
• spdf notation – shows how many e- are in each sublevel,
not how they are distributed in that sublevel
• ex. Nitrogen = 1s22s22p3
• spdf notation (expanded) – shows the distribution in the
last sublevel
• ex. Nitrogen = 1s22s22px12py12pz1
• noble-gas-core abbreviated electron configuration
(condensed electron configuration): replace part of the
config. with a noble gas in brackets.
• Ex. Nitrogen = He 2s22p3
© 2015 Pearson Education, Inc.
Electronic
Structure
of Atoms
6.9 Electron Configurations and the Periodic
Table
• 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 inner shell electrons are called core
electrons. These include electrons in d or f
sublevels, even though they might be of
higher energy than the valence electrons.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.
Periodic Table
• Different blocks on the periodic table correspond to
the filling of 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.
Exceptions to the Aufbau rule: exceptions sometimes
occur when orbitals can be half-filled or filled by moving
electrons up in energy (for sublevels that are close in
energy)
• For example, 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 with f and d
orbitals, as well.
Electronic
Structure
of Atoms
© 2015 Pearson Education, Inc.