Electrons in Atoms

1
Chapter 13

• Electrons in Atoms

2
Greek Idea

• Democritus and Leucippus
• Matter is made up of indivisible particles
• Dalton - one type of atom for each element

3
Thomsons Model

• Discovered electrons
• Atoms were made of positive stuff
• Negative electron floating around
• Plum-Pudding model

4
Rutherfords Model

• Discovered dense positive piece at the center of
  the atom
• Nucleus
• Electrons moved around
• Mostly empty space

5
Bohrs Model

• Why dont the electrons fall into the nucleus?
• Move like planets around the sun.
• In circular orbits at different levels.
• Amounts of energy separate one level from another.

6
Bohrs Model
Nucleus
Electron
Orbit
Energy Levels
7
Bohrs Model

• Further away from the nucleus means more energy.
• There is no in between energy
• Energy Levels

Fifth
Fourth
Third
Increasing energy
Second
First
Nucleus
8
The Quantum Mechanical Model

• Energy is quantized. It comes in chunks.
• A quanta 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.
• Schrodinger derived an equation that described
  the energy and position of the electrons in an
  atom

9
The Quantum Mechanical Model

• Things that are very small behave differently
  from things big enough to see.
• The quantum mechanical model is a mathematical
  solution
• It is not like anything you can see.

10
The Quantum Mechanical Model

• Has energy levels for electrons.
• Orbits are not circular.
• It can only tell us the probability of finding
  an electron a certain distance from the
  nucleus.

11
The Quantum Mechanical Model

• The atom is found inside a blurry electron
  cloud
• A area where there is a good chance of finding an
  electron.
• Draw a line at 90 probability

12
Quantum Numbers

• Principal Quantum Number (n) the energy level
  of the electron.

13
Atomic Orbitals

• Within each energy level the complex math of
  Schrodingers equation describes several shapes.
• These are called
• Atomic Orbitals Regions where there is a high
  probability of finding an electron.

14
S sublevel

• one s orbital for every energy level
• Spherical shaped

• Each s orbital can hold 2 electrons
• Called the 1s, 2s, 3s, etc.. orbitals.

15
P sublevel

• Start at the second energy level
• 3 different directions
• 3 different shapes (orbitals)
• Each can hold 2 electrons

16
P sublevel
17
D sublevel

• Start at the third energy level
• 5 different orbitals
• Each can hold 2 electrons

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F sublevel

• Start at the fourth energy level
• Have seven different shapes
• 2 electrons per shape

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F sublevel
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Summary
of orbitals
Max electrons
Starts at energy level
s
1
2
1
p
3
6
2
5
10
3
d
7
14
4
f
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By Energy Level

• First Energy Level
• only s orbital
• 2 electrons max. per orbital
• only 2 electrons
• 1s2

• Second Energy Level
• s and p orbitals are available
• 2 in s, 6 in p
• 2s22p6
• 8 total electrons

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By Energy Level

• Third energy level
• s, p, and d orbitals
• 2 in s, 6 in p, and 10 in d
• 3s23p63d10
• 18 total electrons

• Fourth energy level
• s,p,d, and f orbitals
• 2 in s, 6 in p, 10 in d, ahd 14 in f
• 4s24p64d104f14
• 32 total electrons

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By Energy Level

• Any more than the fourth and not all the orbitals
  will fill up.
• You simply run out of electrons

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Electron Configurations

• Notation to show the way electrons are arranged
  in atoms.
• Aufbau principle- electrons enter the lowest
  energy first.
• Pauli Exclusion Principle- at most 2 electrons
  per orbital - different spins
• Hunds Rule- When electrons occupy orbitals of
  equal energy they dont pair up until they have
  to .

25
One Problem!!!

• The orbitals do not fill up in a neat order.
• The energy levels overlap
• Lowest energy fill first.

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Electron Configuration

• Lets determine the electron configuration for
  Phosphorus
• Need to account for 15 electrons

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• The first two electrons go into the 1s orbital
• Notice the opposite spins
• only 13 more

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• The next electrons go into the 2s orbital
• only 11 more

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• The next electrons go into the 2p orbital
• only 5 more

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• The next electrons go into the 3s orbital
• only 3 more

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• The last three electrons go into the 3p orbitals.
• They each go into seperate shapes
• 3 upaired electrons
• 1s22s22p63s23p3

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Practice

• Try some of these in the chart on your own!
• Well check your work together

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The easy way to remember

• 1s2

• 2 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2

• 4 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2

• 12 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2

• 20 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

• 38 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2

• 56 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
  4f14 5d10 6p6 7s2

• 88 electrons

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
  4f14 5d10 6p6 7s2 5f14 6d10 7p6

• 108 electrons

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Exceptions to Electron Configuration
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Orbitals fill in order

• Lowest energy to higher energy.
• Adding electrons can change the energy of the
  large orbitals.
• Half filled orbitals have a lower energy.
• Makes them more stable.
• Changes the filling order for some d sublevels.

44
Write these electron configurations

• Chromium - 24 electrons
• 1s22s22p63s23p64s23d4 is expected
• But this is wrong!!

45
Chromium is actually

• 1s22s22p63s23p64s13d5
• Why?
• This gives us two half filled orbitals.
• Slightly lower in energy.
• The same principle applies to copper.

46
Coppers electron configuration

• Copper has 29 electrons so we expect
• 1s22s22p63s23p64s23d9
• But the actual configuration is
• 1s22s22p63s23p64s13d10
• This gives one filled orbital and one half filled
  orbital.
• Remember these exceptions

47
Light

• The study of light led to the development of the
  quantum mechanical model.
• Light is a kind of electromagnetic radiation.
• Electromagnetic radiation includes many kinds of
  waves
• All move at 3.00 x 108 m/s ( c)

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Parts of a wave
Orgin
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Parts of Wave

• Orgin - the base line of the energy.
• Crest - high point on a wave
• Trough - Low point on a wave
• Amplitude - distance from origin to crest
• Wavelength - distance from crest to crest
• Wavelength - is abbreviated l Greek letter lambda.

50
Frequency

• The number of waves that pass a given point per
  second.
• Units are cycles/sec or hertz (hz)
• Abbreviated n the Greek letter nu
• c ln

51
Frequency and wavelength

• Are inversely related
• As one goes up the other goes down.
• Different frequencies of light is different
  colors of light.
• There is a wide variety of frequencies
• The whole range is called a spectrum

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Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Long Wavelength
Short Wavelength
Visible Light
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Atomic Spectrum

• How color tells us about atoms

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Prism

• White light is made up of all the colors of the
  visible spectrum.
• Passing it through a prism separates it.

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If the light is not white

• By heating a gas with electricity we can get it
  to give off colors.
• Passing this light through a prism does something
  different.

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Atomic Spectrum

• Each element gives off its own characteristic
  colors.
• Can be used to identify the atom.
• How we know what stars are made of.

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• These are called discontinuous spectra
• Or line spectra
• unique to each element.
• These are emission spectra
• The light is emitted given off.

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Light is a Particle

• Energy is quantized.
• Light is energy
• Light must be quantized
• These smallest pieces of light are called
  photons.
• Energy and frequency are directly related.

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Energy and frequency

• E h x n
• E is the energy of the photon
• n is the frequency
• h is Plancks constant
• h 6.6262 x 10 -34 Joules sec.
• joule is the metric unit of Energy

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The Math in Chapter 13

• Only 2 equations
• c ln
• E hn
• Plug and chug.

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Examples

• What is the wavelength of blue light with a
  frequency of 8.3 x 1015 hz?
• What is the frequency of red light with a
  wavelength of 4.2 x 10-5 m?
• What is the energy of a photon of each of the
  above?

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An explanation of Atomic Spectra
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Where the electron starts

• When we write electron configurations we are
  writing the lowest energy
• The energy level and electron starts from is
  called its ground state.

Why does hamburger have lower energy than
steak? Because it is in the ground state.
64
Changing the energy

• Lets look at a hydrogen atom

65
Changing the energy

• Heat or electricity or light can move the
  electron up energy levels

66
Changing the energy

• As the electron falls back to ground state it
  gives the energy back as light

67
Changing the energy

• May fall down in steps
• Each with a different energy

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Ultraviolet
Visible
Infrared

• Further they fall, more energy, higher frequency.
• This is simplified
• the orbitals also have different energies inside
  energy levels
• All the electrons can move around.
• Line Spectra

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What is light

• Light is a particle - it comes in chunks.
• Light is a wave- we can measure its wave length
  and it behaves as a wave
• If we combine Emc2 , cln, E 1/2 mv2 and E
  hn
• We can get l h/mv
• The wavelength of a particle.

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Matter is a Wave

• Does not apply to large objects (anything bigger
  than an atom)
• A baseball has a wavelength of about 10-32 m when
  moving 30 m/s
• An electron at the same speed has a wavelength of
  10-3 cm
• Big enough to measure.

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The physics of the very small

• Quantum mechanics explains how the very small
  behaves.
• Classic physics is what you get when you add up
  the effects of millions of packages.
• Quantum mechanics is based on probability because

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Heisenberg Uncertainty Principle

• It is impossible to know exactly the speed and
  velocity of a particle.
• The better we know one, the less we know the
  other.
• The act of measuring changes the properties.

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More obvious with the very small

• To measure where a electron is, we use light.
• But the light moves the electron
• And hitting the electron changes the frequency of
  the light.

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After
Before
Photon changes wavelength
Photon
Electron Changes velocity
Moving Electron
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Valence Electrons

• Electrons in outermost energy level
• Look for highest principle quantum number in
  electron configuration.
• 1s22s22p63s23p3 highest energy level is 3rd,
  which contains a total of 5 electrons, so 5
  valence electrons
• 1s22s22p63s23p64s23d104p1 How many
  valence electrons?

77
Keeping Track of Electrons

• Atoms in the same column
• Have the same outer electron configuration.
• Have the same valence electrons.
• Easily found by looking up the group number on
  the periodic table.
• Group 2A - Be, Mg, Ca, etc.-
• 2 valence electrons

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Electron Dot Structures

• Way of keeping track of the number of valence
  electrons in an atom or ion.
• Shows dots to represent valence electrons.
• Two kinds Ground State and Bonding

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Ground State Dot Diagrams

• Write the symbol.
• Put one dot for each valence electron
• Pair up first two to indicate s orbital
  electrons, then the other three sides dont
  double up until they have to.

X
80
Bonding Electron Dot diagrams

• Write the symbol.
• Put one dot for each valence electron
• Dont pair up until they have to

X
81
Ground State Dot Diagram for Aluminum

• Aluminum has 3 valence electrons.
• First we write the symbol.

Al

• Then add 2 electron to one side.

• Then one at a time to other sides until they are
  forced to pair up.

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Ground State Dot Diagram for Oxygen

• Oxygen has 6 valence electrons.
• First we write the symbol.

O

• Then add 2 electrons to one side.

• Then one at a time to other sides until they are
  forced to pair up.

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The Bonding Dot diagram for Nitrogen

• Nitrogen has 5 valence electrons.
• First we write the symbol.

N

• Then add 1 electron at a time to each side.

• Until they are forced to pair up.

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Write the bonding electron dot diagram for

• Na
• Mg
• C
• O
• F
• Ne
• He

85
Stable Electron Configurations

• All atoms react to achieve noble gas
  configuration.
• Noble gases have 2 s and 6 p electrons.
• 8 valence electrons .
• Also called the octet rule.

Ar
86
Electron Configurations for Cations

• Metals lose electrons to attain noble gas
  configuration.
• They make positive ions.
• If we look at electron configuration it makes
  sense.
• Na 1s22s22p63s1 - 1 valence electron
• Na 1s22s22p6 -noble gas configuration

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Electron Dots For Cations

• Metals will have few valence electrons

Ca
88
Electron Dots For Cations

• Metals will have few valence electrons
• These will come off

Ca
89
Electron Dots For Cations

• Metals will have few valence electrons
• These will come off
• Forming positive ions

Ca2
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Electron Configurations for Anions

• Nonmetals gain electrons to attain noble gas
  configuration.
• They make negative ions.
• If we look at electron configuration it makes
  sense.
• S 1s22s22p63s23p4 - 6 valence
  electrons
• S-2 1s22s22p63s23p6 -noble gas
  configuration.

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Electron Dots For Anions

• Nonmetals will have many valence electrons.
• They will gain electrons to fill outer shell.

P
P-3
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• 1s1
• 1s22s1
• 1s22s22p63s1
• 1s22s22p63s23p64s1
• 1s22s22p63s23p64s23d104p65s1
• 1s22s22p63s23p64s23d104p65s24d10 5p66s1
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67
  s1

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He

• 1s2
• 1s22s22p6
• 1s22s22p63s23p6
• 1s22s22p63s23p64s23d104p6
• 1s22s22p63s23p64s23d104p65s24d105p6
• 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6

2
Ne
10
Ar
18
Kr
36
Xe
54
Rn
86
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S- block
s1
s2
He
H

• Alkali metals all end in s1
• Alkaline earth metals all end in s2
• really have to include He but it fits better
  later.
• He (helium) has the properties of the noble gases.

Li
Be
Na
Mg
K
Ca
Rb
Sr
Cs
Ba
Fr
Ra
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The P-block
p1
p2
p6
p3
p4
p5
B
C
N
O
F
Ne
Al
Si
P
S
Cl
Ar
Ga
As
Ge
Br
Se
Kr
I
Te
Sb
Sn
In
Xe
Tl
Po
Bi
Pb
Rn
At
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Transition Metals -d block
s1 d5
s1 d10
d1
d2
d3
d5
d6
d7
d8
d10
99
F - block

• inner transition elements

100
1 2 3 4 5 6 7

• Each row (or period) is the energy level for s
  and p orbitals.

101

• D orbitals fill up after previous energy level so
  first d is 3d even though its in row 4.

1 2 3 4 5 6 7
3d
102
1 2 3 4 5 6 7
4f 5f

• f orbitals start filling at 4f

103
Writing Electron configurations the easy way

• Yes there is a shorthand

104
Electron Configurations repeat

• The shape of the periodic table is a
  representation of this repetition.
• When we get to the end of the column the
  outermost energy level is full.
• This is the basis for our shorthand.

105
The Shorthand

• Write the symbol of the noble gas before the
  element.
• Then the rest of the electrons.
• Aluminum - full configuration.
• 1s22s22p63s23p1
• Ne is 1s22s22p6
• so Al is Ne 3s23p1

106
More examples

• Ge 1s22s22p63s23p64s23d104p2
• Ge Ar 4s23d104p2
• Hf1s22s22p63s23p64s23d104p65s2 4d105p66s24f145d2
• HfXe6s24f145d2

107
The Shorthand Again
Sn- 50 electrons
The noble gas before it is Kr
Takes care of 36
Next 5s2
Then 4d10
Finally 5p2
Kr
5s2
4d10
5p2