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Electrons in Atoms
Chapter 5
Chem X
“Indeed, to be useful, a model must
be wrong, in some respects – else
it would be the thing itself.”
 Models
are never fixed, but
change to explain new
observations.
The Nuclear Atom and
Unanswered Questions
 Rutherford’s
1.
2.
 In
model lacked details about…
how e- occupy space around the nucleus
Why the negatively charged electrons are
not pulled into the atom’s positive nucleus
early 1900’s…scientists observed that
certain elements emitted visible light when
heated in a flame
The Nuclear Atom and
Unanswered Questions
 In
early 1900’s…scientists observed that
certain elements emitted visible light when
heated in a flame
Analysis showed that an element’s chemical
behavior is related to the arrangement of electrons in
its atoms. More later…. We must first discuss light.
Wave Nature of Light
Radiation – a form of
energy that exhibits wavelike behavior as
it travels through space.
 Electromagnetic
 Travels
at the speed of light through a
vacuum.
Wave Nature of Light
Wave Nature of Light
 Electromagnetic
wavelength –
frequency Amplitude -
Radiation –
Wave Nature of Light
spectrum – all the colors of
the visible spectrum
 Continuous
Wave Nature of Light
 Ex.
5-1 Calculating Wavelength of an EM
Wave – Microwaves are used to transmit
information. What is the wavelength of a
microwave having a frequency of 2.44 x
109 Hz?
8.72 x 10-2 m
Particle Nature of Light
 The
1.
2.
wave model of light cannot explain…
Why heated objects emit only certain
frequencies of llight at a given temperature
Or why some metals emit electrons when
colored light of a specific frequency shines
on them
The Particle Nature of Light
The quantum concept –
waves don’t explain the emission of different
wavelengths of light at different
temperatures
Max Plank (1858-1947) – concluded matter
can gain or lose energy only in small,
specific amounts called quanta
Quantum – the minimum amount of energy
that can be gained or lost by an atom.
The Particle Nature of Light
Prior experience had lead scientists to
believe that energy could be absorbed and
emitted in continual varying quantities, with
no minimum limit to the amount.
Actually, the water’s temperature changes
by infinitesimal steps as its molecules
absorb quanta of energy
E = hn
Particle Nature of Light
effect – electrons, called
photoelectrons, are emitted from a metal’s
surface.
 Metals only ejected electrons above a
certain minimum frequency.
 Photoelectric
Particle Nature of Light
Photon – a particle of electromagnetic radiation
with no mass that carries a quantum
of energy
Einstein calculated
Ephoton = hn
KEY: Einstein was able to explain the photoelectric
affect by…giving electromagnetic radiation
particlelike properties.
Atomic Emission Spectra – the set of
frequencies of the electromagnetic waves
emitted by atoms of the element.
 It is always several individual lines of
color, NOT a continuous range of colors.
 Each element’s atomic emission spectrum
is unique.
Continuous Spectrum (top)
Various elements’ emission spectra
Atomic Emission Spectra
 The
fact that only certain colors appear in
an elements atomic emission spectrum
means that… only certain specific
frequencies of light are emitted.
 i.e. Only photons having certain specific
energies are emitted.
5.2 Quantum Theory and the
Atom
Bohr Model of the Atom
 Answered the question why…atomic
emission spectra are made up of only
certain frequencies
 Correctly predicted… the frequencies
of the lines in hydrogen’s atomic
emission spectrum
Energy States of hydrogen
 Proposed



that…
Ground State – lowest allowable energy
state of an atom
Excited state – when an atom gain energy
e- jumps to higher E level
Ionized state – (in class)
Bohr’s Model of the Atom
He proposed that the atom has only
certain allowable energy states.
ground state – lowest possible
excited state – atom gain energy
ionized state - atom gains
enough E to lose
an e-


Bohr suggested that.. The single electron in a
hydrogen atom moves around the nucleus in only
certain allowed circular orbits.
The smaller the orbit, the lower the atom’s energy
state.
Bohr’s Model of the Atom
In order to go from one energy level to another,
electrons can only gain/lose energy in certain
amounts.
This is the idea of “quantized” energy.
Absorption
Emission
An explanation of hydrogen’s
line spectrum
n=1
ΔE = Efinal - Einitial
= Ephoton
= hn
See next slide for
bigger figures…
An explanation of hydrogen’s
line spectrum
An explanation of hydrogen’s
line spectrum
n=1
n=5
n=4
n=3
ΔE = Efinal - Einitial
= Ephoton
= hn
n=2
n=1
Atomic Emission Spectra
Atomic Emission -
Compound
Flame
Color
LiCl
NaCl
KCl
CaCl2
SrCl2
CuCl2
Electrons as Waves
Louis De Broglie ( 1892-1987 ) - 1924
l = h__
mv
Key: electrons have wave properties.
Quantum Mechanical Model of
the Atom
Heisenberg Uncertainty Principle – it is
fundamentally impossible to know precisely both
the velocity and position of a particle at the same
time
The Schrodinger wave equation
(1926)
Quantum mechanical model of the atom – model
where electrons are treated as waves
- limits the electron’s energy to certain values
-
Each solution is a wave function and is related to
the probability of finding the electron within a
particular volume of space around the nucleus
Atomic orbital – 3 dimensional region around the
nucleus that describes the electron’s probable
location. (See fig. 5-13 – e- density)
Joke…






Heisenberg, Schrodinger and Ohm are in a car. They
get pulled over. Heisenberg is driving and the cop
asks him "Do you know how fast you were going?"
"No, but I know exactly where I am" Heisenberg
replies.
The cop says "You were doing 55 in a 35." Heisenberg
throws up his hands and shouts "Great! Now I'm lost!"
The cop thinks this is suspicious and orders him to
pop open the trunk. He checks it out and says "Do you
know you have a dead cat back here?"
"We do now!" shouts Schrodinger.
The cop moves to arrest them. Ohm resists.
Hydrogen’s Atomic Orbitals
Orbital – contain 90% of e’s total probability
Principal quantum numbers ( n )
Principal energy levels – specified by n
Energy sublevels – as n increases, the
number of sublevels increases
Hydrogen’s Atomic Orbitals
 See
fig. 5-14 (Stadium seating)
Hydrogen’s Atomic Orbitals
Principal
quantum
number (n)
1
2
3
4
Sublevels
(types of
orbitals)
present
Total number of
Number of
orbitals related
orbitals related
to principal
to sublevel
energy level (n2)
5.3 Electron Configurations
 Ground
State Electron Configuration:
configuration – the arrangement
of electrons in an atom
 Electron
5.3 Electron Configurations
 The
1.
2.
3.
4.
aufbau principle –
All orbitals related to an energy sublevel are of
equal energy
In a multi-electron atom, the energy sublevels
within a principal energy level have different
energies
In order of increasing energy, the sequence of
energy sublevels within a principal energy level is
s, p, d, and f
Orbitals related to energy sublevels within one
principal energy level can overlap orbitals related
to energy sublevels within another principal level.
5.3 Electron Configurations
 Pauli
Exclusion Principle–
 A maximum of two electrons may occupy a
single atomic orbital, but only if the electrons
have opposite spins.
5.3 Electron Configurations
 Hund’s
Rule–
 single electrons with the same - spin must
occupy each equal energy orbital before
additional electrons with opposite spins can
occupy the same orbitals.
Quantum Mechanical Model of
the Atom
Heisenberg Uncertainty Principle –
Schrodinger Wave Equation –

waves the
Electrons are treated as ________,
equation can tell you the coordinates most
probable to find an electron of a given energy
atomic orbital –
Joke…






Heisenberg, Schrodinger and Ohm are in a car. They
get pulled over. Heisenberg is driving and the cop
asks him "Do you know how fast you were going?"
"No, but I know exactly where I am" Heisenberg
replies.
The cop says "You were doing 55 in a 35." Heisenberg
throws up his hands and shouts "Great! Now I'm lost!"
The cop thinks this is suspicious and orders him to
pop open the trunk. He checks it out and says "Do you
know you have a dead cat back here?"
"We do now!" shouts Schrodinger.
The cop moves to arrest them. Ohm resists.
Quantum Mechanical Model of
the Atom
Energy
Level
1
2
3
4
# electrons
# orbitals
needed
Types & #’s of
orbitals
Orbitals
s-orbitals
p-orbitals
d-orbitals
f-orbitals
Quantum Mechanical Model of
the Atom
___
4s
___ ___ ___
4p
___ ___ ___ ___ ___
4d
___
___ ___ ___
___ ___ ___ ___ ___
3s
3p
3d
___
___ ___ ___
2s
2p
___
1s
___ ___ ___ ___ ___ ___ ___
4f
Quantum Mechanical Model of
the Atom
___ ___ ___
6p
___
5s
5p
4d
___ ___ ___ ___ ___
3d
4s
___ ___ ___
___
3p
3s
___ ___ ___
___
2p
1s
___ ___ ___ ___ ___
___ ___ ___
___
___
4f
___ ___ ___
4p
2s
___ ___ ___ ___ ___ ___ ___
5d
___
6s
___ ___ ___ ___ ___
Link to Notebook Periodic table
Electron Configuration
bau principle – an electron will occupy
the _____________________ that can
receive it.
 auf
exclusion principle – (no two
electrons can have the same set of four
quantum numbers)
i.e. If two electrons are in the same
orbital, they must have _______________.
 Pauli
Electron Configuration
 Hund’s
Rule – the lowest energy
configuration for an atom is the one having
the maximum number of unpaired
electrons allowed by the Pauli principle in
a particular set of degenerate orbitals.
 i.e. as electrons fill into degenerate
orbitals one must go into each orbital will
___________ spins before you
_______________________.
Electron Configuration
 Noble
gas Short-hand - a shortened econfiguration that starts with the previous
noble gas
 Orbital Notation - uses lines to denote
each orbital with electrons written as
arrows
(the orbital types 4s etc. are sometime
written underneath the lines)
Electron Configuration
Hydrogen
Scandium
Helium
Zinc
Lithium
Barium
Carbon
Thallium
Argon
Radon
Calcium
Electron Configuration
Hydrogen 1s1
Helium
1s2
Lithium
1s22s1
Carbon
1s22s22p2
Argon
1s22s22p6 3s23p6
Calcium 1s22s22p6 3s23p64s2
Electron Configuration
Scandium 1s22s22p6 3s23p64s23d1
Zinc
1s22s22p6 3s23p64s23d10
Barium
1s22s22p6 3s23p64s23d104p65s24d105p66s2
Thallium
1s22s22p6 3s23p64s23d104p65s24d105p66s2
4f145d106p1
Radon
1s22s22p6 3s23p64s23d104p65s24d105p66s2
4f145d106p6
Electron Configuration
Electron configuration - shows how many
electrons are in each type of orbital
Fluorine 1s22s22p5
Sulfur
1s22s22p63s23p4
Germanium 1s22s22p63s23p64s23d104p2
Radon
1s22s22p63s23p64s23d104p65s24d105p6
6s24f145d106p6
Electron Configuration
Noble gas Short-hand - a shortened econfiguration that starts with the previous
noble gas
 Fluorine
[He]2s22p5
 Sulfur
[Ne] 3s23p4
 Germanium [Ar] 4s23d104p2
 Radon
[Xe] 6s24f145d106p6
Electron Configuration
Orbital Notation - uses lines to denote each orbital with
electrons written as arrows
(can use long or shorthand)
(the orbital types 4s etc. are sometime written underneath the lines)
[He]2s2p  
[Ne] 3s3p  
[Ar] 4s3d    4p  .
[Xe] 6s4f       5d    
6p  
Radon
[Xe] 6s24f145d106p  
It’s usually only necessary to write orbital notation for the highest
energy orbital (as done with radon).
Fluorine
Sulfur
Germanium
Radon
Electron Configuration Review
Write the electron configuration for each of
the following:
1.
1.
2.
3.
4.
Argon
Carbon
Osmium
Barium
Electron Configuration Review
1.
What is the electron configuration for
Xenon?
2.
What is the orbital notation for Xenon?
3.
What’s Xenon’s highest energy level?
What’ Xenon’s highest energy orbital?
4.
Electron Configuration Review
1.
How many valence electrons do
each of the following elements
have?
1.
2.
3.
4.
5.
Chlorine
Calcium
Aluminum
Carbon
Tin