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• Why is the sky blue?
Chemistry Chapter 11
Arrangement of Electrons in Atoms
The 1998 Nobel Prize in
Physics was awarded "for
the discovery of a new
form of quantum fluid with
fractionally charged
excitations." At the left
is a computer graphic of
this kind of state.
The Puzzle of the Atom
Protons and electrons are attracted to each
other because of opposite charges
Electrically charged particles moving in a
curved path give off energy
Despite these facts, atoms don’t collapse
(Rutherford’s Model did not explain why! We need
a new model!)
What is light?
Light is made up of fluctuating electric
and magnetic fields that do not require
the existence of matter.
Visible light is a kind of electromagnetic
radiation (form of E with wavelike
behavior as travels through space.)
All form of light travel at a speed of
3.0 x 108 m/s
Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the
electron as a particle.
His son, George Thomson won the Nobel prize for
describing the wave-like nature of the electron.
The
electron is
a particle!
The
electron is
an energy
wave!
Confused??? You’ve Got Company!
“No familiar conceptions can be
woven around the electron;
something unknown is doing we
don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
Electromagnetic radiation propagates
through space as a wave moving at the
speed of light. (ALL FORMS!)
c = 
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
Types of electromagnetic radiation:
Light as Particles
• 1900’s German physisist Max Planck studied
emission of light by hot objects.
• Suggested objects emitted in small specific
amounts called quanta.
• Quantum: the minimum quantity of E that can be
lost or gained by an atom.
• Proposed E = hv
– h is “Planck’s Constant” 6.626 x 10-34 Js
The energy (E ) of electromagnetic
radiation is directly proportional to the
frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x 10-34 J·s)
 = frequency, in units of hertz (hz, sec-1)
Photoelectric Effect (early 1900’s)
• Could NOT be explained by the wave
theory of light.
• Refers to the emission of e’ from a
metal when light shines on it.
– For a given metal, no e’ were
emitted if the light was below a
certain v
– Wave theory said any v could supply
enough E to eject any e’
What are wave-like and particle like
properties???
Wave Properties
Particle Properties
Measurable frequency
and wavelength
Ability to interfere and
diffract
Absorbed and emitted by
matter (photoelectric
effect)
Emission of light by hot
objects
Significant feature is
repetitive nature
Creation of line-emission
spectra of elements
Albert Einstein (1905)
• Suggested dual wave-particle nature
• Each particle carried a quantum of E.
– Einstein called these particles “photons”
– E photon = hv
• For each e’ to be removed, must be struck by
single photon with minimum E required to knock e’
loose.
• e’ of different metals require different minimum v
to exhibit photoelectric effect.
Hydrogen-Atom Line Emission
Spectrum
• When current is passed through a gas at low P, the
potential E of some of the gas atoms increases.
– Ground state: lowest E state of an atom
– Excited state: atom has higher potential E
than the ground state.
• When at atom returns to ground state, it gives off
E. (neon lights)
• Results in Hydrogen’s Line Emission Spectrum.
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum
Electron transitions
involve jumps of
definite amounts of
energy.
This produces bands
of light with definite
wavelengths.
The Bohr Model of the Atom
I pictured
electrons orbiting
the nucleus much
like planets
orbiting the sun.
Neils Bohr
But I was
wrong! They’re
more like bees
around a hive.
WRONG!!!
What does a line emission spectrum
tell you?
• Proves that e’ are constantly moving from one E
level to another.
• Proves that the distance between the two levels is
always the same.
• Use
– Identification of unknown samples
Bohr’s Model of H Atom
• Linked an atom’s e’ to photon emission
• Said that e’ can circle the nucleus only is
specific paths or orbits.
• Orbits are separated from the nucleus by
large empty spaces (nodes) where the ecannot exist. (Rungs on a ladder)
• Remember, E increases as e’ are further
from the nucleus.
Bohr’s Model of H Atom
• While an e’ is in orbit, it cannot gain nor
lose E
• It can, move to a higher E level IF it gains
the amount of E = to the difference
between the higher E orbit and the lower
energy orbit.
• When falls, it emits a PHOTON = to the
difference.
• Absorption gains E/Emission gives off E
Take the good with the bad (Bohr’s
Model)
• Good:
– Led scientists to believe a similar model could
be applied to all atoms.
• Bad:
– Yikes!!! Did not explain the spectra of atoms
with more than one e’ or chemical behavior of
atoms!
The Wave-like Electron
The electron propagates
through space as an energy
wave. To understand the
atom, one must understand
the behavior of
electromagnetic waves.
Louis deBroglie
Schrodinger Wave Equation

d
h

V 
8  m dx
2
2
2
2
 E
Equation for probability of a
single electron being found
along a single axis (x-axis)
Erwin Schrodinger
The electron as a standing wave:
Standing waves do not propagate through space
Standing waves are fixed at both ends
An orbital is a region within an atom where there
is a probability of finding an electron. This is a
probability diagram for the s orbital in the first
energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.
Heisenberg Uncertainty Principle
“One cannot simultaneously
determine both the position
and momentum of an electron.”
.
…
Werner
Heisenberg
Sizes of s orbitals
Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.
The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape
P orbital shape
There are three dumbbell-shaped p orbitals in
each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.
Things get a bit more
complicated with the five d
d orbital shapes
orbitals that are found in
the d sublevels beginning
with n = 3. To remember
the shapes, think of “double
dumbells”
…and a “dumbell
with a donut”!
Shape of f orbitals
Orbital filling table
Electron configuration of the
elements of the first three series
Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half
fill its 3d sublevel
Copper steals a 4s electron to FILL
its 3d sublevel
Pauli Exclusion Principle
No two electrons in an atom
can have the same four
quantum numbers.
Wolfgang
Pauli
Quantum Numbers
Each electron in an atom has a unique set of 4
quantum numbers which describe it.




Principal quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number
Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Number of electrons
that can fit in a shell:
2n2
Angular Momentum Quantum
Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell) in
which the electron is located.
Magnetic Quantum Number
The magnetic quantum number, generally
symbolized by m, denotes the orientation of the
electron’s orbital with respect to the three axes in
space.
Spin Quantum Number
Spin quantum number denotes the behavior
(direction of spin) of an electron within a magnetic
field.
Possibilities for electron spin:
1

2
1

2
Assigning the Numbers
1. Principle: n = shell (period)
2. Angular Momentum: l = shape (s=0, p=1,
d=2)
3. Magnetic: ml = +l to -l
4. Spin: +1/2 or – 1/2
Chlorine
• Write the orbital notation
• Write the electron configuration
• So, can we now write the quantum #’s for each e’?
e’
Electron Conf. and Quant #’s
n
principle
l
shape
Ml
spin
orientation
e’ config.
1,2
1
0
0
+/- 1/2
-
1s2
3,4
2
0
0
+/- 1/2
2s2
5-10
2
2
2
1
1
1
-1
0
+1
+/- 1/2
2p6
11,12
3
0
0
+/- 1/2
3s2
13-17
3
3
3
1
1
1
-1
0
+1
+/- 1/2
3p5
Key Terms
• Ground State Electron Configurations
(1s22s2…/Orbital Notations __ __
__ __ __
• Aufbau Principle (lowest first/periodic guide)
• Hund’s Rule: Single e’ before pairing begins
• Pauli Exclusion (up/down—no 2 e’ same set of
quant. #’s)
• Highest Occupied Level
• Inner-Shell e’
• Noble Gas Notation
Orbitals in outer energy levels DO penetrate into
lower energy levels. Penetration #1
This is a probability
Distribution for a
3s orbital.
What parts of the
diagram correspond
to “nodes” – regions
of zero probability?
Which of the orbital types in the 3rd energy level
Does not seem to have a “node”?
WHY NOT?
Penetration #2
Principle, angular momentum, and magnetic quantum numbers: n,
l, and ml
Chemistry Chapter 5
The Periodic Law
Law of Mendeleev:
– Properties of the elements recur in
regular cycles (periodically) when the
elements are arranged in order of
increasing atomic mass.
Mendeleev’s Periodic Table
Dmitri Mendeleev
Missing?
14
Si
28.09
??
50
Sn
118.71
Named missing
element
“Ekasilicon”
From base word
“eka” meanging
next in order
“Ekasilicon”
Predicted
Properties
Atomic Mass
72 amu
Density
5.5 g/cm3
Melting Point
825° C
Observed
Properties
“Ekasilicon”
Predicted
Properties
Observed
Properties
Atomic Mass
72 amu
72.61 amu
Density
5.5 g/cm3
5.32 g/cm3
Melting Point
825° C
938° C
“Ekasilicon”
Predicted
Properties
Observed
Properties
Oxide Formula
XO2
GeO2
Chloride Formula
XCl4
GeCl4
Modern Russian Table
Orbital filling table
Periodic Table with Group Names
s-block elements:
• Called main group elements
– 1: ns1 Alkali Metals
– 2: ns2 Alkaline Earth Metals
p-block elements:
• Called main group elements
– 13: ns2np1 Boron Family
– 14: ns2np2 Carbon Family
– 15: ns2np3 Nitrogen Family
– 16: ns2np4 Oxygen Family
– 17: ns2np5 Fluorine Family(Halogens)
– 18: ns2np6 Noble Gases
Metalloids:
•
•
•
•
Mostly brittle solids
Harder and denser than s-block
Softer and less dense than d-block
With exception of Bismuth, found in nature only as
compound
• Once obtained as free metals, that are stable in
the presence of air
d-block elements:
• Called transition metals
– metals with typical metallic properties
– Less reactive than group 1 or 2
– Some are so unreactive that they do not
easily form compounds (e.g. palladium,
platinum, and gold)
f block elements:
• Between Group 3 & 4
• Shiny metals
• Similar in reactivity to Group 2
– Lanthanides (period 6):
• rare earth metals
– Actinides (period 7):
• All radioactive
• First 4 have been found naturally, rest are
lab made
Determination of Atomic Radius:
Half of the distance between nuclei in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels
Table of
Atomic
Radii
Ionization Energy - the energy required to remove an
electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
Tends to decrease down a group
Outer electrons are further from the
nucleus
Ionization of Magnesium
Mg + 738 kJ  Mg+ + eMg+ + 1451 kJ  Mg2+ + eMg2+ + 7733 kJ  Mg3+ + e-
Table of 1st Ionization Energies
Another Way to Look at Ionization Energy
Electron Affinity - the energy change associated
with the addition of an electron
Affinity tends to increase across a period
Affinity tends to decrease as you go down
in a group
Electrons further from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals
Table of Electron Affinities
Ionic Radii
Cations
Anions
Positively charged ions
Smaller than the corresponding
atom
Negatively charged ions
Larger than the corresponding
atom
Table of Ion Sizes
Valence Electrons
or
“Outer Shell Electrons”
• The electrons available to be lost, gained, or
shared in the formation of chemical compounds.
• Valence Electrons hold atoms together in chemical
compounds.
Electronegativity
A measure of the ability of an atom in a chemical
compound to attract electrons.
(same trends as electron affinity)
Electronegativities tend to increase across
a period
Electronegativities tend to decrease down a
group or remain the same
•F most electronegative: 4.0 (Pauling scale)
•Nitrogen, Oxygen, and Halogens most
•Alkali and Alkali Earth metals least (1.0-0.7)
•Values used in chemical BONDING
Periodic Table of Electronegativities
Summation of Periodic Trends
Long
Wavelength
=
Low Frequency
=
Low ENERGY
Short
Wavelength
=
High Frequency
=
High ENERGY
Wavelength Table