1. Electronic
Structure
of Atoms
Electronic Structure of Atoms
Resources
• Our TB: Ch. 6 of Chemistry: The Central
Science AP version (10th edition)
• Powerpoint * (from pearson) and in-class work
• POGIL activities: (1) Analysis of Spectral
Lines and (2) Interaction of Radiation and
Matter
• Online resources for our TB (in particular
online quiz)
• Chem tours from ch. 7 of the W.W.
Norton online book by Gilbert:
• http://www.wwnorton.com/college/chemistry/gi
lbert2/contents/ch07/studyplan.asp
• Animations from Glencoe site:
http://glencoe.mcgraw-
hill.com/sites/0023654666/student_view0/chapter7/
• Extra quizzes from Glencoe
http://glencoe.mcgraw-
hill.com/sites/0023654666/student_view0/chapter7/
• Video lectures from chem guy
http://www.kentchemistry.com/moviesfiles/chemguy/AP/C
hemguyAtomicTheory.htm
• Handouts and practice problems from M.
Brophy’s web site
2. Electronic
Structure
of Atoms
Chapter 6
Electronic Structure
of Atoms
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
John D. Bookstaver
St. Charles Community College
St. Peters, MO
2006, Prentice Hall, Inc.
3. Electronic
Structure
of Atoms
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 ().
4. 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.
6. Electronic
Structure
of Atoms
The Nature of Energy
• The wave nature of light
does not explain how
an object can glow
when its temperature
increases.
• Max Planck explained it
by assuming that
energy comes in
packets called quanta.
7. Electronic
Structure
of Atoms
The Nature of Energy
• Einstein used this
assumption to explain the
photoelectric effect.
• He concluded that energy
is proportional to
frequency:
E = h
where h is Planck’s
constant, 6.63 10−34 J-s
(i.e. units for h are J•s)
8. Electronic
Structure
of Atoms
The Nature of Energy
• Therefore, if one knows the
wavelength of light, one
can calculate the energy in
one photon, or packet, of
that light:
c =
E = h
9. Electronic
Structure
of Atoms
For electromagnetic radiation
animation and problems see:
http://www.wwnorton.com/coll
ege/chemistry/gilbert2/tutorial
s/interface.asp?chapter=chap
ter_07&folder=frequency_wa
velength
For All Chem tours for the electrons
in atoms and periodic properties
topic see:
http://www.wwnorton.co
m/college/chemistry/gilb
ert2/contents/ch07/study
plan.asp
Recommeded chem tours
animations:
Electromagnetic radiation
Light Emission and Absorbtion
Bohr Model of the Atom
De Broglie Wavelngth
Quantum numbers
Electron configuration
11. Electronic
Structure
of Atoms
The Nature of Energy
• One does not observe
a continuous
spectrum, as one gets
from a white light
source.
• Only a line spectrum of
discrete wavelengths
is observed.
12. Electronic
Structure
of Atoms
Go To Glencoe
Animation
http://glencoe.com/sites/common_a
ssets/advanced_placement/chemist
ry_chang9e/animations/chang_7e_
esp/pem1s3_1.swf
POGIL activity on Spectral Lines
(To Complete)
13. Electronic
Structure
of Atoms
The Nature of Energy
• 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).
14. Electronic
Structure
of Atoms
The Nature of Energy
• Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this
way:
2. Electrons in permitted orbits
have specific, “allowed”
energies; these energies will
not be radiated from the atom.
15. Electronic
Structure
of Atoms
The Nature of Energy
• Niels Bohr adopted
Planck’s assumption and
explained these
phenomena in this way:
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
16. Electronic
Structure
of Atoms
The Nature of Energy
The energy absorbed or emitted
from the process of electron
promotion or demotion can be
calculated by the equation:
E = −RH ( )
1
nf
2
1
ni
2
-
where RH is the Rydberg
constant, 2.18 10−18 J, and ni
and nf are the initial and final
energy levels of the electron.
17. Electronic
Structure
of Atoms
Go To Glencoe and
Norton Animations
http://glencoe.com/sites/common_a
ssets/advanced_placement/chemist
ry_chang9e/animations/chang_7e_
esp/pem1s3_1.swf
POGIL activity on Interaction of
Radiation and Matter
(To Complete)
Go to Chem tour for Bohr Model of
atom (and Rydberg equation)
http://www.wwnorton.com/college/c
hemistry/gilbert2/tutorials/interface.
asp?chapter=chapter_07&folder=hy
drogen_energies
18. Electronic
Structure
of Atoms
The Wave Nature of Matter
• Louis de Broglie posited that if light can
have material properties, matter should
exhibit wave properties.
• He demonstrated that the relationship
between mass and wavelength was
=
h
mv
19. Electronic
Structure
of Atoms
The Uncertainty Principle
• Heisenberg showed that the more precisely
the momentum of a particle is known, the less
precisely is its position known:
• In many cases, our uncertainty of the
whereabouts of an electron is greater than the
size of the atom itself!
(x) (mv)
h
4
20. Electronic
Structure
of Atoms
Quantum Mechanics
• Erwin Schrödinger
developed a
mathematical treatment
into which both the
wave and particle nature
of matter could be
incorporated.
• It is known as quantum
mechanics.
21. Electronic
Structure
of Atoms
The Quantum Mechanical Model
• Energy is quantized - It comes in chunks.
• A quantum is the amount of energy needed to
move from one energy level to another.
• Since the energy of an atom is never “in
between” there must be a quantum leap in
energy.
• In 1926, Erwin Schrodinger derived an
equation that described the energy and
position of the electrons in an atom
• (this slide from: J. Hushen’s presentation on Atomic Structure at
http://teachers.greenville.k12.sc.us/sites/jhushen/Pages/AP%20Chemistry.aspx)
22. Electronic
Structure
of Atoms
Schrodinger’s Wave Equation
2
2
2 2
8
d
h E
V
m dx
Equation for the
probability of a single
electron being found
along a single axis (x-axis)
Erwin Schrodinger
(this slide from: J. Hushen’s presentation on Atomic Structure at
http://teachers.greenville.k12.sc.us/sites/jhushen/Pages/AP%20Chemistry.aspx)
23. Electronic
Structure
of Atoms
Quantum Mechanics
• The wave equation is
designated with a lower
case Greek psi ().
• The square of the wave
equation, 2, gives a
probability density map of
where an electron has a
certain statistical likelihood
of being at any given instant
in time.
24. Electronic
Structure
of Atoms
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.
26. Electronic
Structure
of Atoms
Azimuthal 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.
28. Electronic
Structure
of Atoms
Magnetic Quantum Number, ml
• Describes the three-dimensional
orientation of the orbital.
• Values 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, etc.
32. Electronic
Structure
of Atoms
s Orbitals
Observing a graph of
probabilities of finding
an electron versus
distance from the
nucleus, we see that s
orbitals possess n−1
nodes, or regions
where there is 0
probability of finding an
electron.
34. Electronic
Structure
of Atoms
d Orbitals
• Value of l is 2.
• Four of the
five orbitals
have 4 lobes;
the other
resembles a p
orbital with a
doughnut
around the
center.
35. Electronic
Structure
of Atoms
Energies of Orbitals
• For a one-electron
hydrogen atom,
orbitals on the same
energy level have
the same energy.
• That is, they are
degenerate.
36. Electronic
Structure
of Atoms
Energies of Orbitals
• As the number of
electrons increases,
though, so does the
repulsion between
them.
• Therefore, in many-
electron atoms,
orbitals on the same
energy level are no
longer degenerate.
37. Electronic
Structure
of Atoms
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.
38. Electronic
Structure
of Atoms
Spin Quantum Number, ms
• This led to a fourth
quantum number, the
spin quantum number,
ms.
• The spin quantum
number has only 2
allowed values: +1/2
and −1/2.
39. Electronic
Structure
of Atoms
Pauli Exclusion Principle
• No two electrons in the
same atom can have
exactly the same energy.
• For example, no two
electrons in the same
atom can have identical
sets of quantum
numbers.
40. Electronic
Structure
of Atoms
Go To
www.ptable.com
IMPORTANT
Use periodic Table to help you write
electron configurations of atoms (and
ions)
Dynamic Periodic Table and
Investigate (play with) the Orbitals
option (on Top Tabs) for quantum
numbers, orbitals and electron
configurations of various elements
Go To
Glencoe site for animations on
electron configuration
http://glencoe.mcgraw-
hill.com/sites/0023654666/student_
view0/chapter7/animations_center.
html#
43. Electronic
Structure
of Atoms
Electron Configurations
• Distribution of all
electrons in an atom.
• Consist of
Number denoting the
energy level.
Letter denoting the type
of orbital.
Superscript denoting the
number of electrons in
those orbitals.
46. Electronic
Structure
of Atoms
Periodic Table
• We fill orbitals in
increasing order of
energy.
• Different blocks on
the periodic table,
then correspond to
different types of
orbitals.
50. Electronic
Structure
of Atoms
ELECTRON SPIN
•1920--chemists realized that since electrons interact
with a magnetic field, there must be one more concept
to explain the behavior of electrons in atoms.
•ms--the 4th quantum number; accounts for the reaction
of electrons in a magnetic field
MAGNETISM
•magnetite--Fe3O4, natural magnetic oxide of iron
•1600--William Gilbert concluded the earth is also a large spherical magnet with magnetic south at the north pole (Santa's habitat).
•NEVER FORGET: opposites attract & likes repel
PARAMAGNETISM AND UNPAIRED ELECTRONS
•diamagnetic--not magnetic [magnetism dies]; in fact they are slightly repelled. All electrons are PAIRED.
•paramagnetic--attracted to a magnetic field; lose their magnetism when removed from the magnetic field; HAS ONE OR MORE
UNPAIRED ELECTRONS
•ferromagnetic--retain magnetism upon introduction to, then removal from a magnetic field
•All of these are explained by electron spins
•Each electron has a magnetic field with N & S poles
•electron spin is quantized such that, in an external magnetic field, only two orientations of the electron magnet and its spin are
possible
•+/- 1/2
•H is paramagnetic; He is diamagnetic, WHY?
•H has one unpaired electron
•He has NO unpaired electrons; all spins offset and cancel each other out
•(Taken from summary notes posted on M. Brophy’s website)
51. Electronic
Structure
of Atoms
•What about ferromagnetic?
clusters of atoms have their unpaired electrons aligned
within a cluster, clusters are more or less aligned and
substance acts as a magnet. Don't drop it!!
•When all of the domains, represented by these arrows
are aligned, it behaves as a magnet. This is what happens
if you drop it! The domains go indifferent directions and it
no longer operates as a magnet.
(Taken from summary notes posted on M. Brophy’s website)
52. Electronic
Structure
of Atoms
Activities and Problem set __
TB ch. 6 – all sections required for
SAT II and AP exams and most
are required for regents exam
View and take notes on the
recommended animations
POGIL activities on (1) Analysis of
Spectral Lines and (2) Interaction
of Radiation and Matter
Online practice quiz due by ______
• Ch 6 Problems: write out questions (or
photocopy them) ; write out answers &
show work
• First carefully study the sample
exercises in chapter 6 (you don’t have to
copy them out) and then DO all in-
chapter practice exercises according to
the directions above.
• Do all GIST, and Visualizing concepts,
problems
• end of chapter 6 exercises: _________
