Chapter 4 Arrangement of Electrons in Atoms

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Chapter 4 Arrangement of Electrons in Atoms

• Section 4.1
• The Development of a New Atomic Model

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Light Is it a Wave?

• 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 (X-rays, UV rays, microwaves, radio waves).
  
• They all move at the same speed, 3.00 x 108 m/s,
  which is called symbol (c)

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Parts of a Wave
Origin
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Parts of a Wave

• Origin - 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.

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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

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Frequency and Wavelength

• They 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, which is
  called a spectrum.

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Radio waves
Microwaves
Infrared .
Ultraviolet
X-Rays
Gamma Rays
Long Wavelength
Short Wavelength
Visible Light
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Light Is it a Particle?

• Studied by Max Planck in 1900
• He wondered why a hot object does not emit (give
  off) electromagnetic radiation all the time (it
  should if light were just a wave)
• He proposed that a hot object emits energy in
  small packets (chunks) called quanta.
• A quantum is the minimum quantity of energy that
  can be lost or gained by an atom.

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Light Is it a Particle?

• Planck devised a formula E h?, where v is
  frequency and h 6.626 x 10-34 Js (Plancks
  constant).
• This was expanded on by Einstein in 1905 and
  refers to the duality of light (its both a
  wave a particle).
• Einstein called the particles of light photons.
• A photon has zero mass at rest and its energy
  depends on the frequency.
• Einstein explained the photoelectric effect
  (emission of electrons by a metal when light
  strikes it) by referring to these photons.

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An Explanation of Atomic Spectra

• Just Watch!

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Where the electron starts

• The lowest energy level of an electron is called
  its ground state.
• When an electron has higher energy than in its
  ground state, its called an excited state.

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Changing the Energy

• Lets look at a hydrogen atom

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Changing the Energy

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

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Changing the Energy

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

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Changing the Energy

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

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Visible Ultraviolet Ultraviolet Ultraviolet

Infrared
Visible
Ultraviolet


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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.

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Bohrs Model

• Why dont the electrons fall into the nucleus?
• Move like planets around the sun.
• In circular orbits (orbitals) at different
  levels.
• Amounts of energy separate one level from
  another.
• A good model, but did not explain everything.

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Bohrs Model
Nucleus
Electron
Orbit
Energy Levels
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Bohrs Model

Fifth

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

Fourth
Third
Increasing energy
Second
First
Nucleus
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Chapter 4 Arrangement of Electrons in Atoms

• Section 4.2
• The Quantum Model of the Atom

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Electrons as Waves

• If light could behave as a particle and wave,
  what about electrons?
• Louis de Broglie in 1924 did this and found
  electrons could be diffracted (bent) and can
  interfere with each other.
• This lead to the question, if electrons are both
  waves and particles, then where are they in the
  atom?

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Heisenberg Uncertainty Principle

• Studied by Werner Heisenburg in 1927
• He said, that 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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After
Before
Photon
Photon changes wavelength
Moving electron
Electron changes velocity
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The Quantum Mechanical Model

• Developed by Erwin Schrodinger derived an
  equation that described the energy and position
  of the electrons in an atom
• 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.

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The Quantum Mechanical Model Cont

• 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.

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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.

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The Quantum Mechanical Model Just watch

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

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Atomic Orbitals Quantum Numbers

• 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.
• They specify the properties of the orbitals and
  the electrons in each orbital.
• It is like the chemical address or seating chart
  for an electron.

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Principal Quantum Numbers

• Principal Quantum Number (n) is the energy level
  of the electron.
• The maximum energy level for an atom is the same
  as the period it is on in the Periodic Table.
• The total number of electrons per energy level is
  found by 2n2.

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Angular Momentum Quantum Number

• Indicates shape of orbital.
• Symbolized by l
• For every n, the number of orbital shapes l is 0
  and all positive integers ? n-1

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S orbitals

• 1 s orbital for every energy level
• Spherical shaped
• Each s orbital can hold 2 electrons
• Called the 1s, 2s, 3s, etc.. orbitals.
• Their l 0

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P orbitals

• Start at the second energy level
• 3 different directions
• 3 different shapes
• Holds total of 6 electrons
• Their l 1

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P Orbitals
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D orbitals

• Start at the third energy level
• 5 different shapes
• Holds total of 10 electrons
• Their l 2

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F orbitals

• Start at the fourth energy level
• Have seven different shapes
• Holds a total of 14 electrons
• Their l 3

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F orbitals
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Summary
of shapes
Max of 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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Magnetic Quantum Number

• Indicates orientation of orbital.
• Symbolized by m
• It is found by all integers from l to l
• For example, if l 1, then m -1, 0, and 1

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Spin Quantum Number

• Indicates spin of orbital.
• Since every orbital holds 2 electrons, there are
  only 2 values for spin, 1/2 and 1/2

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Chapter 4 Arrangement of Electrons in Atoms

• Section 4.3
• Electron Configurations

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Electron Configurations

• The way electrons are arranged in atoms.
• Aufbau principle- electrons enter the lowest
  energy first.
• This causes difficulties because of the overlap
  of orbitals of different energies.
• Pauli Exclusion Principle- at most 2 electrons
  per orbital - different spins (one 1/2, other
  1/2)

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Electron Configurations

• Hunds Rule- When electrons occupy orbitals of
  equal energy they dont pair up until they have
  to .
• Lets determine the electron configuration for
  Phosporus
• 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 separate shapes
• 3 unpaired electrons
• 1s22s22p63s23p3

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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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This is the Diagonal Rule!
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Orbitals fill in order

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

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Write these electron configurations

• Titanium - 22 electrons
• 1s22s22p63s23p64s23d2
• Vanadium - 23 electrons 1s22s22p63s23p64s23d3
• Chromium - 24 electrons
• 1s22s22p63s23p64s23d4 is expected
• But this is wrong!!

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Chromium is actually

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

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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