# Atomic Structure Electron Configurations Periodicity

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```					Atomic Structure
Electron Configurations & Periodicity

Chemistry 40S
Unit 1
M. Patenaude
GPHS Science Dept

1
Introduction
   Atomic structure explains chemical properties
and patterns of chemical reactivity.
   Chemical reactions involve electrons. Knowing
where the electrons are, how many, and what
their energy levels helps explain many chemical
phenomena.
   Spectroscopy is used to explore atomic
discussion of the nature of light.

Slide 2
The Nature of Light
Particles or Waves??

   Light must be made of
particles because it…
   travels in a vacuum
   reflects off of objects
   exerts force (on the
tails of comets)
   Light must consist of
waves because it…
   reflects like waves
By the end of the 19th century,
   refracts and diffracts    scientists had concluded that
   exhibits interference     light is composed of WAVES!

Slide 3
   Visible light is just one form of electromagnetic
   Light propagates in space as a wave
   In vacuum, speed of light is constant and given
the symbol “c”, c = 3.00 x 108 m/s
   Light waves have amplitude, frequency, and
wavelength.
     = wavelength - distance between consecutive
crests,  = frequency, # of crests that pass a
given point in one second (SI Unit is s-1 or Hz)
    c = and  = c/
Slide 4
James Maxwell proposed in the mid-1800’s that light
is composed of perpendicular electric and magnetic waves
Slide 5
are distances – S.I. unit is the meter

Which of the above waves has
the higher frequency?

Slide 6
The 7 Regions of the E.M. Spectrum

Slide 7
Slide 8
Sample Calculation –
wavelength/frequency conversion

   Calculate the frequency of visible light having a
wavelength of 485 nm?

Remember to use S.I. units in your calculations!

     =c/
= (3.00 x 108 m/s) ÷ (485 x 10-9 m)
= 6.55 x 1014 s-1

   What colour of visible light is this?

Slide 9
Sample Calculation –
wavelength & frequency conversion
   CO2 absorbs light with a wavelength of 0.018 mm.
Which frequency is this?
c 3.00  108 ms1
                  1.7  1013 s1
 0.018  103 m
Remember that 1 Hertz (Hz)  1 s-1

   What is the wavelength of WFAE at 90.7 MHz?
c 3.00  108 ms1
           6 1
 3.31 m
  90.7  10 s
Slide 10
Particle Theory of Light
   In 1900, Max Planck turned the world of physics on
its head by resurrecting the particle theory of light.
   Planck was trying to explain the “ultraviolet
   He proposed that light is composed of particles
(quanta) each carrying a fixed amount of energy.
   The amount of energy per quantum is directly
proportional to the frequency of the light:
hc                        34
E  h    or   E        , h  6.626  10         J s


Slide 11
Planck’s Equation – sample problem
   The wavelength of maximum visual acuity in
humans is 550 nm. (green light)
   What is the energy of a single photon having this
wavelength?

   Ans. 3.61 x 10-19 J per photon

   What is the energy of a mole of photons at 550
nm?

   Ans. 218 kJ/mol

Slide 12
Energy, wavelength, frequency

a)   X-Rays
b)   Red light
c)   Green light

       Rank the above in order of …
    increasing wavelength
    decreasing energy
    increasing frequency

Slide 13
Atomic Emission Spectra
“Line Spectra” vs “Continuous Spectra”
    Historic work (1800’s) involving light emitted by pure
elements in gas phase, subjected to very high voltages.
    The light was passed through a prism or diffraction
grating to produce a spectrum.
    Each element emitted only certain wavelengths of light –
a LINE SPECTRUM instead of the more familiar
“continuous” spectrum seen in the rainbow.
    Balmer and Rydberg described the lines
mathematically, for hydrogen.
    Hydrogen’s emission spectrum has several lines in the
visible region – called the “Balmer series”

Slide 14
A Continuous Spectrum
contains all wavelengths of light – a rainbow

Slide 15
Slide 16
Bohr’s Explanation of Line Spectra

   Neils Bohr developed a mathematical model that could
explain the observation of atomic line spectra.
   He proposed that electrons orbit the nucleus in certain
“allowed” orbits - or energy levels. That is, the electron’s
energy is QUANTIZED (not continuous).
   Working on a model of single-electron atoms, Bohr derived
an equation to calculate the energy of an electron in the nth
orbit of such an atom:

18  Z2 
En  2.18  10  2 
where Z = atomic number

n              and n = electron energy level

Slide 17
Bohr’s Quantum Model for
Hydrogen
   The electron in hydrogen occupies
discrete energy levels.
   The atom does not radiate energy
when the electron is in an energy
level.
   When an electron falls to a lower
energy level, a quantum of radiation
is released with energy equal to the
difference between energy levels.
energy levels if the atom absorbs a
energy.

Slide 18
The Hydrogen Line Spectrum

Slide 19
Bohr’s Model: Calculations

   Show that an              Step 1: Calculate E4 & E2
electron transition           E4 = -1.36 x 10-19 J
from n = 4 to n = 2           E2 = - 5.45 x 10-19 J
results in the
emission of visible
light. Calculate the      Step 2: Find DE42
wavelength of the             DE = E2 – E4 = 4.09 x 10-19 J
light emitted.
   Step 3: Calculate 
   = 4.86 x 10-7 m = 486 nm

Slide 20
What’s Wrong with Bohr’s Model??

   Although it works great        Ultimately, the failure of
for single-electron             Bohr’s model lay in the
atoms, Bohr’s model             fact that he treated the
fails for atoms with 2 or       electron as a charged
more electrons!                 particle orbiting the
   It was a huge leap              nucleus like a planet
forward, but was                around the sun.
fundamentally flawed.          Electrons are more
complicated …

Slide 21
Wave-Particle Dual Nature of
Electrons
   Einstein’s famous equation, E = mc2, suggests that energy
and mass are related – one can convert matter directly into
energy.
   Louis de Broglie made the leap that if light can behave as
wave/particles, then so can matter!
   He showed that the wavelength of a baseball is negligible
(as expected), but that the wavelength of an electron was
on the order of magnitude equal to that of electromagnetic
   This is what Bohr had missed – an atomic model must
make use of the wave-nature of electrons to be complete!

Slide 22
Visualizing the Wave-Nature of
Electrons

Slide 23
Heisenberg’s “Uncertainty Principle”
The more precisely the position is
determined, the less precisely the
momentum is known in this instant,
and vice versa.

   It is impossible to determine BOTH the
position and momentum (velocity) of an
electron - the act of observing the
electron changes it.
   The electron exists in nature as a wave-
particle (whatever that is) until an
experimenter observes it and imposes
one of the two natures upon it!

Slide 24
"When it comes to atoms,
language can only be used as in
poetry. The poet too, is not nearly
so concerned with describing
facts as with creating images."
- Niels Bohr to Werner Heisenberg

25
The Heisenburg uncertainty principle is the catch 22 of
physics. Most modern physics and chemistry is largely based on
the study of electrons. In order to view an electron we must use
light to observe it, but in the process of observing we change
the electron’s momentary existence because the light necessary
to view it takes the form of a photon particle. The photon particle
travels at approximately the same velocity as the electron
particle, so that
when it impacts the electron it alters its
position as when two asteroids collide
into each other at similar velocities but
different angles.

Therefore, according to Heisenburg,
it is impossible to know decisively an
electron’s position and momentum.

Slide 26
Heisenburg proposed that this inability to perceive reality on an atomic
level is not the result of technological ignorance, but rather a fact of the
laws of nature. In other words, Heisenburg believed that no matter how
technologically advanced we become we would never be able to fully
perceive and record atomic reality. This means that an electron’s orbital
path around the nucleus of an atom can only be statistically
approximated.

This is why the quantum atomic model has a cloud of electrons, not a few
electrons following exact elliptical orbits as in Bohr's atomic model. Each
electron in the quantum atomic model is a calculated probability, so the
cloud describes not the exact orbit of millions of electrons, but rather the
probability of the position of only a few electrons at any given time.
Slide 27
The Schrödinger Wave Equation
   Schrödinger applied de Broglie’s concept of
matter waves to the electron
   He explained the quantum energies of the
electron orbits in the Bohr model of the atom as
vibration frequencies of electron "matter
waves" around the atom's nucleus.
   He provided a mathematical equation (wave
function) to describe electron waves.
   Physicists had difficulty explaining the
MEANING of the wave function itself, but it
turns out that its square gives the probability of
finding the electron in a particular region of
space in the atom.

Slide 28
Schrodinger Equation, cont.
   These regions of probability are called orbitals, and this
concept replaces that of the electron “orbit”.
   Instead of electrons orbiting like planets around the sun,
Schrodinger’s picture shows an “electron cloud” surrounding
the nucleus and doesn’t state anything about the path or
position of electrons within that cloud.
   Orbital “pictures” shown in texts represent regions of space
where the probability of finding an electron is 90% or greater.

Slide 29
Quantum Numbers
   Schrodinger's wave equation used three constants
(“quantum numbers”) that were used to describe orbitals
– regions of space where an electron has a probability of
being found.
   His equation used 3 quantum numbers to describe the
size & energy, shape, and orientations of the orbitals in
atoms.
   The three quantum numbers: n, l and ml .
   This set of three numbers provides us with an “orbital
address” for an electron within an atom – we can
describe the orbital in which an electron is found if we
know these three quantum numbers.

Slide 30
Principal Quantum Number, n
   n can have values > 1, integers only.
   For element atoms on the periodic table in their
ground states, 1  n  6
   As n increases, energy of electron increases and
size of orbital increases
   Larger n means greater probability of greater
distance from the nucleus.
   Electrons with same value of n are said to be in
the same “electron shell”

Slide 31
Shape Quantum Number, l
   The values for l (lower case letter L, script) are limited by
the principal energy level – they range from 0 to “n–1”
   Example: If n = 3, l can be 0, 1 or 2

   l gives information about shape of orbital
l =0        s orbital (found on all principle energy levels)
l =1        p orbitals (found on level 2 and higher)
l =2        d orbitals (found on level 3 and higher)
l =3        f orbitals (found on level 4 and higher)

Slide 32
Graphs show
Probability
vs Distance
from nucleus

s Orbitals
NODE

on three
energy
levels

Slide 33
p and d Orbital shapes

Slide 34
Slide 35
Orientation Quantum Number, ml
   The values of ml are limited by the value of l for
a given orbital. They include positive and
negative integers from -l through +l , inclusive

   For a p orbital, l = 1 , so ml can be: 1, 0, 1
   ml gives information about orientation of
orbitals in space
   The # of allowed values gives the number of
orbitals of this type on a given energy level. (3,
in the case of p orbitals)
Slide 36
Electron Spin Quantum Number, ms
   A fourth quantum number, ms
describes electron spin
(either +½ or –½)
   Each electron in atom has a
unique set of these four
quantum numbers.
   Electrons in orbitals with same
n and l values are said to be in
A spinning negative charge        the same subshell.
creates a magnetic field.        Electrons with all three numbers
The direction of spin d           the same, n, l , and ml , are in
determines the direction of
the same orbital.
the field.
Slide 37
Pauli Exclusion Principle
   Electrons have negative charge and repel each
other. How are the electrons in an atom
distributed?
   Wolfgang Pauli proposed that no two electrons
in a given atom can be described by the
same four quantum numbers!
   The first three quantum numbers determine an
orbital – therefore spins must be opposite!
   Practical result is that each orbital can hold a
maximum of two electrons, with opposite spins.

Slide 38
Concept Check
   Which of the following sets of       What is the maximum
quantum numbers are NOT               number of electrons in an
allowed in the hydrogen               atom that can have these
atoms? For those that are             quantum numbers?
incorrect, state what is wrong.
   n=4
n       l     ml      ms
   n = 5, ml = +1
2      1       -1     +½
   n = 5, ms = + ½
1      1       0      -½                n = 2, l = 1
8      7       -6     +½                n = 3, l = 2
1      0       2      -½                n = 3, ml = 4

Slide 39
Electron Configurations

   Electrons orbitals are defined by their quantum
numbers, n, l and ml .
   Each electron in an atom has a unique set of 4
quantum numbers.
   No two electrons can have the same “address”,
i.e., the same 4 quantum #’s
   Rules define how multiple electrons will be
distributed among the possible energy levels

Slide 40
Aufbau Principle
   In the ground state, the electrons occupy the lowest
available energy levels. An atom is in an excited
state if one or more electrons are in higher energy
orbitals.
   In atoms with more than one electron the lower
energy orbitals get filled by electrons first!
   This is the Aufbau principle, which is named after the
German word which means “to build up”.
   When one describes the locations of the electrons in
work up to the highest energy electron.

Slide 41
Hund’s Rule
   A set of orbitals is said to be “degenerate” if the orbitals
possess the same energies. For example, all three “2p”
orbitals on energy level 2 are degenerate. All five “3d”
orbitals on energy level 3 are also degenerate.
   When filling a set of degenerate orbitals, Hund said that
electrons should be left unpaired as long as possible so
as to minimize electron-electron repulsions within the
orbitals.
   When each degenerate orbital has one electron,
electrons will then pair, spins opposed, until that sub-
shell is filled.

Slide 42
Summary of Distribution Rules
   Electrons distribute to lower energy levels
until they are filled, before occupying higher
levels. (Aufbau Principle)
   Electrons will spread out as much as possible
within a sub-shell. (Hund’s Rule)
   Electrons will pair up, two to an orbital, spins
opposed, until that sub-shell is filled. (Pauli
Exclusion)
   Every electron will have unique set of 4
quantum numbers.

Slide 43
Electronic Configurations
   An electron configuration is a way of describing the
locations of electrons within an atom.
   Usually written for the ground state of an atom: when its
electrons occupy orbitals giving the lowest possible
overall energy for the atom
   Each subshell is designated by “n” and the type of
orbital. Such as: 1s 3p 4d
   The number of electrons in an orbital is shown as a
superscript: 1s2 3p3          4d9
   How does one determine the order of filling of the
orbitals in an atom?? Which orbitals get filled first??

Slide 44
Orbital Energies in Multielectron
Atoms

Slide 45
The Order of Filling of Orbitals
Simply follow the order of elements in the periodic table!

Slide 46
Another useful device…
Start at the top and

1s,
2s,
2p, 3s,
3p, 4s,
3d, 4p, 5s…

Slide 47
Examples of Electron Configurations

   Helium has 2 electrons in the 1s orbital, He: 1s2
   Carbon has 6 electrons, C: 1s2 2s2 2p2
   Calcium has 20 electrons, Ca: 1s2 2s2 2p6 3s2 3p6 4s2
   Calcium cation, Ca2+: 1s2 2s2 2p6 3s2 3p6
   Sulfide anion, S2–: 1s2 2s2 2p6 3s2 3p6

   Notice that the superscripts add up to the total number of
electrons on the atom or ion! Each of these examples
represents the “ground state” of the atom/ion because
the electrons are in the lowest possible energy levels.

Slide 48
Orbital Diagrams

Slide 49
Noble Gas “shorthand” Notation
   The core electrons can be indicated by a short
form: use the last noble gas to indicate filled
shells.
   Find the last noble gas before the element under
consideration. Start the electron configuration
with the symbol for this noble gas element and
just tack on the extra electrons:
   Ca: 1s2 2s2 2p6 3s2 3p6 4s2
   Ca: [Ar] 4s2

Ar: 1s2 2s2 2p6 3s2 3p6
Slide 50
There Have to be EXCEPTIONS
   As atoms get larger, their      Cr: [Ar] 4s1 3d5
outermost electrons             Cu: [Ar] 4s1 3d10
become closer in energy
   Large atoms have                    In both cases, one of the 4s
“exceptions” to the                  electrons is promoted to the
Aufbau Principle in the              “3d” subshell.
locations of their valence
electrons. We will
   Fe2+: [Ar] 3d6
memorize two: Cr and Cu
   Cu2+: [Ar] 3d9
   Transition metal cations
are also strange in their
electron configurations             Notice that the 4s electrons
are lost before the 3d
electrons!

Slide 51
Paramagnetism & Diamagnetism

   An atom is said to be            An atom is diamagnetic if
paramagnetic if it                all its electrons are paired –
possesses unpaired                it will NOT possess a net
electrons.                        magnetic field because
   These electrons will result       each electron’s field will be
in the atom possessing a          cancelled
net magnetic field               These atoms are NOT
   It will then be attracted         attracted into an external
into an external magnetic         magnetic field
field
   What is “ferromagnetism”?

Slide 52
Testing for Paramagnetism
See Kotz & Treichel page 290

Slide 53
Electron Configuration
Causes Periodic Trends
   “Periodicity” refers to similarities in behavior and
reactivity due to similar outer shell electron
configurations.
   All the Alkali Metals have one unpaired valence
electron; all the Noble Gases have completely
filled subshells
   We will examine periodic trends in atomic radius,
ionization energy, electronegativity, and electron
affinity

Slide 54
“Valence Electrons”
   Valence electrons are those electrons in the
highest principal energy level within an atom - the s
and p electrons in the outermost shell of an atom in
its ground state.
   These are the electrons involved in forming bonds
with other atoms during chemical reactions.
   Most common oxidation states (ion charge) for the
element can be derived from valence electrons
   Completely filled, half filled and empty sub-shells
have special stability (not sure why!?)

Slide 55
Valence Configurations
   The “valence configuration” is that part of an
electron configuration that describes the valence
electrons.
   For example, sodium’s valence configuration is
just “3s1”.
   The valence configuration of bromine is “4s24d5”.
We leave out the “3d10” electrons because they
are not in the outermost principal energy level.

Slide 56
Periodic Trends: Atomic Size
   Properties of elements and their ions are
determined by electron configurations
   Individual atoms cannot be accurately
measured.
   Atomic sizes are determined from bond
lengths of compounds with various atoms

Br-Br               Br-C
2.28 A              1.91 A
r = 1.14 A          rcarbon = 1.91-1.14
= 0.77 A

Slide 57
Atomic Size, cont.
   Atomic radii increase within a group (column) as
the principal quantum number of outermost shell
increases
   Electrons of outermost shells have greater
probability of being farther from the nucleus
   Atomic size decreases across a row (period)
from Left to Right, because the effective nuclear
charge increases.
   There are more and more protons in the nuclei
attracting more and more electrons!

Slide 58

Atoms get larger
as you move
down a group

Atoms get
smaller as you
go across a
period

Slide 59
Slide 60
Ionization Energies
   First Ionization energy: energy required to
remove an electron from the ground state
atom in the gaseous state to make a cation
   1st ionization of Mg:
Mg  Mg+         I.E. = 738 J/mol
[Ne]3s2  [Ne]3s1

   2nd ionization of Mg
Mg +  Mg2+          1450 J/mol
[Ne]3s1  [Ne]
Slide 61
Trend in First Ionization Energies

   First ionization energies decrease down a group
as radii increase:     Na > K > Cs
   First ionization energies increase across a row
as radii decrease: Li < F < Ne
   For similar reasons that explained trends in
   For a given element, IE2 always > IE1 because of
the same # protons in nucleus holding onto fewer
electrons.

Slide 62
Slide 63
Trends in First
Ionization Energies

Slide 64
First Ionization Energy (cont’d)
   Lowest IE1 are in the             Oxygen has slightly lower
lower left corner of               IE1 than nitrogen because
periodic table (Fr and Cs)         it takes more energy to
   Irregularities in IE1 trends       disrupt N’s half-filled
arise from special                 orbital configuration
electron configurations.          Boron has slightly lower
   Half filled sub-shells are         IE1 than Beryllium
especially stable: every           because its first electron
orbital has only one               is removed from the p-
electron.                          subshell further from the
nucleus than the first
electron removed in Be

Slide 65
Electron Affinity
   Energy change when a gas phase atom gains an
electron to become an anion.
Cl (g) + e-  Cl– (g) EA = – 349 kJ/mol

   The more negative an EA, the greater the affinity for
electrons!
   Electron affinities “tend” to become more negative
across a period and less negative moving down a group
   EA least negative for alkaline earth and noble gas
elements: they have full or half full s- or p-orbitals and it
would take a lot of energy to add another electron

Slide 66
Slide 67
Slide 68
Slide 69
Ion Sizes
   Cations: radius is always smaller than the
neutral atom.
   Loss of electrons (to make cation) leaves fewer
electrons attracted to same # of protons
   OR….. loss of outermost shell altogether as in
group 1A [Ne]3s1  [Ne] for sodium
   Anions: radius is always larger than neutral
atom.
   Same # protons pulling on more electrons;
also electrons repel each other.

Slide 70
Slide 71
Isoelectronic Atoms and Ions
   Isoelectronic atoms or ions have same electron
configurations – the same number of electrons
in same orbitals
   Na+ , Ne and F- are isoelectronic: 1s2 2s2 2p6
   For isoelectronic atoms and ions, the species
with the most protons will be smallest
   It has greater nuclear attraction for the same
number of electrons, so the electron shells will
be pulled closer to nucleus.

Slide 72
Bonding in Molecules

Ionic Bonds
Polar Covalent Bonds
Non-Polar Covalent Bonds

73
Definitions

   ELECTRONEGATIVITY is the attraction an atom
has for electrons in a chemical bond.
   A POLAR BOND is one between atoms with
different electronegativities - one atom attracts
electrons more strongly than the other. As a
result, one atom acquires a partial positive
charge and the other a partial negative charge.

Slide 74
More Definitions

   A DIPOLE is a separation of charge. A polar
bond is an example of a dipole.
   A DIPOLE MOMENT is the magnitude of a
dipole - it depends on the size of the charges
involved.
   A POLAR MOLECULE is one that possesses an
overall dipole moment.

Slide 75
Classifying Chemical Bonds

   IONIC bonds are characterized as…
   bonds between metals and non-metals
   bonds between cations and anions
   bonds involving a transfer of electrons (from a
metal to a non-metal)

   The best classification of an ionic bond uses the
concept of electronegativity.

Slide 76
Classifying Chemical Bonds

   An ionic bond forms between two atoms with a
large difference in electronegativity - usually
stated as being greater than about 1.7 Paulings.

   e.g. NaCl: difference is 2.0 Paulings
   e.g. CuO: difference is 1.7 Paulings

Slide 77
Classifying Chemical Bonds

   COVALENT bonds are usually described as…
   bonds between two non-metal atoms
   bonds involving a sharing of electrons

   Using the concept of electronegativity, we can
see that there are two types of covalent bonds.

Slide 78
Classifying Chemical Bonds

   POLAR COVALENT BONDS result from
unequal sharing of electrons by atoms that
creates a small dipole.
   This type of bond occurs when there is a small
but significant difference in electronegativity -
between 0.4 and 1.7 Paulings.
   e.g. N-H: diff is 0.9 Paulings
   e.g. Cu-S: diff is 0.7 Paulings

Slide 79
Classifying Chemical Bonds

   NON-POLAR COVALENT BONDS result from
an equal sharing of electrons.
   There is no dipole created since the electrons
are shared equally.
   This bond forms between atoms with similar (or
the same) electronegativities - differences
between 0 and 0.4 Paulings.

Slide 80
Exercise

   Classify the following chemical bonds. Show
   Li-F
   Br-Br
   Fe-O
   Al-Cl
   C-S
   C-H

Slide 81
Polarity of Molecules

   A diatomic molecule (e.g. HF) will be polar if its
bond is polar.
   Larger molecules are polar only when the
following criteria are met:
   they possess at least one polar bond
   their geometries do not result in the dipoles
cancelling each other out! (e.g. CO2)

Slide 82

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