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Chapter 4 Arrangement of Electrons in Atoms 4.1 The

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Title: Chapter 4 Arrangement of Electrons in Atoms 4.1 The


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

2
Properties of Light
  • Electromagnetic Radiation
  • EM radiation are forms of energy which move
    through space as waves
  • There are many different types of EM waves
  • visible light
  • x-rays
  • ultraviolet light
  • infrared light
  • radio waves

Waves Not Just For The Beach Anymore!
3
EM Waves
  • Move at speed of light 3.00 x 108 m/s
  • Speed is equal to the frequency times the
    wavelength c v?
  • Frequency (v) is the number of waves passing a
    given point in one second
  • Wavelength (?) is the distance between peaks of
    adjacent waves
  • Speed of light is a constant, so v? is also a
    constant v and ? must be inversely proportional

4
(No Transcript)
5
Light and Energy The Photoelectric Effect
  • Electrons are emitted from a metal when light
    shines on the metal
  • Incoming EM radiation from the left ejects
    electrons, depicted as flying off to the right,
    from a substance.
  • Radiant energy is transferred in units (or
    quanta) of energy called photons
  • (Max Planck)

6
Photoelectric Effect
energy
absorption spectrum
p
ground state e-
no
e-
When a specific or quantized amount of energy is
exposed to the atom, the electron jumps from its
ground or original state to an excited state
7
When the excited electron returns to lower
energy levels, it releases energy in the form of
light
Photoelectric Effect
energy photon
emission spectrum!
p
excited state e-
no
e-
travels at the speed of light (3.00 x 108 m/s)
8
  • A photon is a particle of energy having a rest
    mass of zero and carrying a quantum of energy
  • A quantum is the minimum amount of energy that
    can be lost or gained by an atom
  • Energy of a photon is directly proportional to
    the frequency of radiation
  • E hv (h is Plancks constant,
  • 6.62554 x 10 -24 J sec)

9
Electromagnetic Spectrum
  • Wavelength increases?
  • Frequency decreases?
  • Energy decreases?

10
Electromagnetic Spectrum
11
Wave-Particle Duality
  • Energy travels through space as waves, but can be
    thought of as a stream of particles (Einstein)
  • Each particle has 1 quantum of energy.

12
Line Spectrums
  • Ground State The lowest energy state of an atom
  • Excited State A state in which an atom has a
    higher potential energy than in its ground state
  • example Neon lights

13
Emissions Spectrum
  • Bright line spectrum Light is given off by
    excited atoms as they return to lower energy
    states
  • Light is given off in very definite wavelengths
  • A spectroscope reveals lines of particular
    colors- light passed through a prism specific
    frequencies given off.

14
The Hydrogen Line Spectrum
  • Definite frequency
  • Definite wavelength

http//student.fizika.org/nnctc/spectra.htm
15
Bohr Model
e-
Niels Bohr
p
no
Electrons circle around the nucleus on their
energy level
Energy levels
16
The Bohr Model of the Atom
  • Electron Orbits, or Energy Levels
  • Electrons can circle the nucleus only in allowed
    paths or orbits
  • The energy of the electron is greater when it is
    in orbits farther from the nucleus
  • The atom achieves the ground state when atoms
    occupy the closest possible positions around the
    nucleus
  • Electromagnetic radiation is emitted when
    electrons move closer to the nucleus.

17
The Bohr Atomic Model
18
Energy transitions
  • Energies of atoms are fixed and definite
    quantities
  • Energy transitions occur in jumps of discrete
    amounts of energy
  • Electrons only lose energy when they move to a
    lower energy state

19
Shortcomings of the Bohr Model
  • Doesn't work for atoms larger than hydrogen
  • (more than one electron)
  • Doesn't explain chemical behavior

20
Chapter 4 Arrangement of Electrons in Atoms
  • 4.2 The Quantum Model of the Atom

21
Electrons as Waves and Particles
  • Louis deBroglie (1924)
  • Electrons have wavelike properties
  • Consider the electron as a wave confined to a
    space that can have only certain frequencies

22
The Heisenberg Uncertainty Principle
"It is impossible to determine simultaneously
both the position and velocity of an electron or
any other particle.
Werner Heisenberg- 1927
  • Electrons are located by their interactions with
    photons
  • Electrons and photons have similar energies
  • Interaction between a photon and an electron
    knocks the electron off of its course

23
The SchrØdinger Wave Equation
  • Proved quantization of electron energies and is
    the basis for Quantum Theory
  • Quantum theory describes mathematically the wave
    properties
  • of electrons and other very small particles
  • Electrons do not move around the nucleus in
    "planetary orbits"

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24
The SchrØdinger Wave Equation
  • Electrons exist in regions called orbitals
  • An orbital is a three-dimensional region around
    the nucleus that indicates the probable location
    of an electron
  • SchrØdinger equation for probability of a single
    electron being found along a single axis (x-axis)

25
Atomic Orbitals Quantum Numbers
  • Quantum Numbers specify the properties of atomic
    orbitals and the properties of the electrons in
    orbitals
  • Principal Quantum Number (n)
  • Angular Momentum Quantum Number (l)
  • Magnetic Quantum Number (m)
  • Spin Quantum Number

26
Principal Quantum Number (n)
  • Indicates the main energy levels occupied by the
    electron
  • Values of n are positive integers
  • n 1 is closest to the nucleus, and lowest in
    energy
  • The number of orbitals possible per energy level
    (or "shell") is equal to n2

27
Quantum Theory
Energy levels, n
e-
n 1
p
n 2
no
n 3
n 4
28
Angular Momentum Quantum Number (l)
  • Indicates the shape of the orbital
  • Number of orbital shapes n
  • Possible values are l 0, 1, 2, or 3
  • Shapes are designated s, p, d, f
  • S shape is spherical
  • P shape is a dumbbell, or figure 8

Click Here!
29
Magnetic Quantum Number (m)
  • The orientation of the orbital around the nucleus
  • s orbitals have only one possible orientation,
  • m 0
  • p orbitals have three possible,
  • m 1, 0 or -1
  • d orbitals have five possible,
  • m -2, -1, 0, 1, or 2
  • f orbitals have 7 possible orientations

30
Spin Quantum Number
  • Indicates the fundamental spin states of an
    electron in an orbital
  • Two possible values for spin, 1/2, -1/2
  • A single orbital can contain only two electrons,
    which must have opposite spins

31
Diagrams of the Orbitals
32
Summary of the Quantum Numbers
33
Chapter 4 Arrangement of Electrons in Atoms
  • 4.3 Electron Configurations

34
Energy Levels
35
Writing Electron Configurations
Aufbau Principle An electron occupies the
lowest-energy orbital that can receive it
Pauli Exclusion Principle No two electrons in the
same atom can have the same set of four quantum
numbers
  • Rules

Hund's Rule Orbitals of equal energy are each
occupied by one electron before any orbital is
occupied by a second electron, and all electrons
in singly occupied orbitals must have the same
spin
36
Orbital Notation (1 of 3)
  • Unoccupied orbitals are represented by a line,
    _____
  • Lines are labeled with the principal quantum
    number and the sublevel letter
  • Arrows are used to represent electrons
  • Arrows pointing up and down indicate opposite
    spins

Pauli Exclusion Principle No two electrons in the
same atom can have the same set of four quantum
numbers (occupy the same space _at_ the same time)
37
Writing Electron Configurations
Hund's Rule Orbitals of equal energy are each
occupied by one electron before any orbital is
occupied by a second electron, and all electrons
in singly occupied orbitals must have the same
spin
1st e-
2nd e-
3rd e-
4th e-
38
Configuration Notation (2 of 3)
  • The number of electrons in a sublevel is
    indicated by adding a superscript to the sublevel
    designation
  • Hydrogen 1s1
  • Helium 1s2
  • Lithium 1s2 2s1

Aufbau Principle An electron occupies the
lowest-energy orbital that can receive it
(Always start with n 1 and work your way up)
39
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
18
Sublevel Blocks on the Periodic Table
40
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41
Fill-In the Periodic Table using Energy Levels
Sublevels
42
Orbital Filling Order
START
1s
2s
2p
3p
3d
3s
4f
4p
4d
4s
5p
5d
5f
5s
6d
6s
6p
7s
7p
FINISH
43
Noble Gas Notation (3 of 3)
  • The configuration begins with the preceding noble
    gass symbol in brackets and is followed by the
    rest of the configuration for the particular
    element.
  • Ne 3s23p5

44
Terms
  • Highest occupied energy level
  • The electron containing energy level with the
    highest principal quantum number
  • Inner shell electrons
  • Electrons that are not in the highest energy
    level
  • Octet Rule
  • Highest energy level s and p electrons are
    filled (8 electrons)

45
Octet Rule
  • Characteristic of noble gases, Group 18
  • Exceptions Hydrogen Helium
  • Noble gas configuration
  • Outer main energy level fully occupied, usually
    (except for He) by eight electrons
  • This configuration has extra stability

46
Survey of the Periodic Table
  • Elements of the Fourth Period
  • Irregularity of Chromium
  • Expected 1s22s22p63s23p64s23d4
  • Actual 1s22s22p63s23p64s13d5

Rule Sublevels are most stable when they are
either half or completely filled. Electrons will
shift to different energy levels to accommodate
this stability whenever possible.
47
Survey of the Periodic Table
  • Several transition and rare-earth elements borrow
    from smaller sublevels in order to half fill
    larger sublevels
  • i.e. d borrows 1 e- from s
  • This accounts for some of the unexpected electron
    configurations found within the transition
    elements.

48
Magnetic Fields Electrons
  • Paramagnetic
  • When an atom has unpaired electrons, it will be
    attracted into a magnetic field
  • Ex 1s22s22p2
  • Dimagnetic
  • When an atom has only paired electrons, it will
    be slightly repelled by a magnetic field
  • Ex 1s22s22p63s2
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