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Chapter 4 Arrangement of Electrons in Atoms

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Title: Chapter 4 Arrangement of Electrons in Atoms


1
Chapter 4 Arrangement of Electrons in Atoms
  • Section 4.1
  • The Development of a New Atomic Model

2
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)

3
Parts of a Wave
Origin
4
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.

5
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

6
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.

7
Radio waves
Microwaves
Infrared .
Ultraviolet
X-Rays
Gamma Rays
Long Wavelength
Short Wavelength
Visible Light
8
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.

9
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.

10
An Explanation of Atomic Spectra
  • Just Watch!

11
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.

12
Changing the Energy
  • Lets look at a hydrogen atom

13
Changing the Energy
  • Heat or electricity or light can move the
    electron up energy levels

14
Changing the Energy
  • As the electron falls back to ground state it
    gives the energy back as light

15
Changing the Energy
  • May fall down in steps
  • Each with a different energy

16
Visible Ultraviolet Ultraviolet Ultraviolet

Infrared
Visible
Ultraviolet


17
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.

18
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.

19
Bohrs Model
Nucleus
Electron
Orbit
Energy Levels
20
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
21
Chapter 4 Arrangement of Electrons in Atoms
  • Section 4.2
  • The Quantum Model of the Atom

22
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?

23
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.

24
After
Before
Photon
Photon changes wavelength
Moving electron
Electron changes velocity
25
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.

26
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.

27
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.

28
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

29
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.

30
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.

31
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

32
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

33
P orbitals
  • Start at the second energy level
  • 3 different directions
  • 3 different shapes
  • Holds total of 6 electrons
  • Their l 1

34
P Orbitals
35
D orbitals
  • Start at the third energy level
  • 5 different shapes
  • Holds total of 10 electrons
  • Their l 2

36
F orbitals
  • Start at the fourth energy level
  • Have seven different shapes
  • Holds a total of 14 electrons
  • Their l 3

37
F orbitals
38
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
39
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

40
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

41
Chapter 4 Arrangement of Electrons in Atoms
  • Section 4.3
  • Electron Configurations

42
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)

43
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

44
  • The first two electrons go into the 1s orbital
  • Notice the opposite spins
  • only 13 more

45
  • The next electrons go into the 2s orbital
  • only 11 more

46
  • The next electrons go into the 2p orbital
  • only 5 more

47
  • The next electrons go into the 3s orbital
  • only 3 more

48
  • The last three electrons go into the 3p orbitals.
  • They each go into separate shapes
  • 3 unpaired electrons
  • 1s22s22p63s23p3

49
The easy way to remember
  • 1s2
  • 2 electrons

50
Fill from the bottom up following the arrows
  • 1s2 2s2
  • 4 electrons

51
Fill from the bottom up following the arrows
  • 1s2 2s2 2p6 3s2
  • 12 electrons

52
Fill from the bottom up following the arrows
  • 1s2 2s2 2p6 3s2 3p6 4s2
  • 20 electrons

53
Fill from the bottom up following the arrows
  • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
  • 38 electrons

54
Fill from the bottom up following the arrows
  • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
  • 56 electrons

55
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

56
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

57
This is the Diagonal Rule!
58
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

59
Write these electron configurations
  • Titanium - 22 electrons
  • 1s22s22p63s23p64s23d2
  • Vanadium - 23 electrons 1s22s22p63s23p64s23d3
  • Chromium - 24 electrons
  • 1s22s22p63s23p64s23d4 is expected
  • But this is wrong!!

60
Chromium is actually
  • 1s22s22p63s23p64s13d5
  • Why?
  • This gives us two half filled orbitals.
  • Slightly lower in energy.
  • The same principal applies to copper.

61
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
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