Chemistry

1
Chemistry

• Chapter 5

2
Wave Nature of Light

• Visible light is a type of electromagnetic
  radiation
• It is a form of energy which exhibits wavelike
  behavior as it travels through space
• Other examples of electromagnetic radiation
  include microwaves, x-rays and radio waves

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Electromagnetic Radiation

• Made of photons
• A photon is a discrete packet of electromagnetic
  energy
• The Energy can be calculated as either
• Ehv or
• Ehc/?

4
Characteristics of Waves

• Wavelength (?) Lambda is the shortest distance
  between equivalent points on a continuous wave,
  it is measured from crest to crest and is usually
  expressed in meters
• Frequency (v) Nu is the number of waves that
  pass a given point per second. One Hertz (Hz
  The SI unit of frequency) equals one wave per
  second.

5
Wave Characteristics

• Amplitude is the waves height from origin to
  crest
• Wavelength and frequency do not affect the
  amplitude
• Wavelength and frequency are inversely
  proportional (as one increases the other
  decreases)

6
Electromagnetic Wave Relationship

• c ?v
• c the speed of light in a vacuum
• ? the wavelength
• v the frequency
• All electromagnetic waves, including visible
  light, travel at a speed of 3.00 x 108
  m/s in a vacuum

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Relationships

• c ??
• v c/?
• ? c/v
• MHz 106Hz
• Hz s-1
• nm 10-9m
• c 3.00 x 108m/s

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Wavelength
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Electromagnetic Waves
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• The visible light spectrum is from about 350nm
  (violet) to 800nm (red)
• Blue is around 450nm
• Below 350nm is ultraviolet above 800nm is
  infra-red

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Calculations

• A radio station broadcasts at 122.0 MHz.
  Calculate the wavelength of this frequency.
• Calculate the frequency of a radiation which is
  614 nm in length.
• What color is this radiation?

13
Practice Problems

• Page 140 1,2 4
• Page 145 14 (Figure 5.5 is on Page 139)
• Page 166 58 (nm 10-9m)

14
Particle Nature of Light

• In addition to exhibiting wavelike behaviors,
  light also behaves as a particle
• As a result, when objects are heated they will
  only emit electrons when light of a specific
  frequency shines on them
• Example Iron is gray at room temperature, glow
  red when heated, then orange and finally blue at
  excessively high temperature

15
Particle Nature

• The different frequencies of the colors cannot be
  explained by the wave nature of light, they are
  explained by the gain or loss of energy
• This energy is in the form of a Quantum
• A Photon is a particle which carries a quantum of
  energy
• The energy of the photon depends on the frequency

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Atomic Emission Spectra

• Neon lights are explained by the wave model of
  light
• The light is produced by passing electricity
  through the gas
• If the light neon emits is passed through a prism
  we dont get the full range of colors like we do
  with visible light
• Instead we observe discrete lines which
  correspond to the radiation emitted by neon

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Separation of light by a prism according to
wavelength
18
Continuous, emission, and absorption spectra
19
Spectra of Gas Discharges

• http//astro.u-strasbg.fr/koppen/discharge/
• http//chemistry.bd.psu.edu/jircitano/periodic4.ht
  ml

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Quick Check

• Page 145 10
• Page 166 34, 35, 38, 41

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Quantum Theory and the Atom

• Bohrs Model of the Atom proposed a quantum
  model of the hydrogen atom
• Predicted the frequencies of the lines in the
  atomic emission spectrum

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Ground State

• The lowest allowable energy state
• When electrons move into a higher energy level,
  they are said to be excited

23
Quantum Number

• The "primary quantum number," which is given the
  symbol n, corresponds to the energy level
• n 1, 2, 3, and so on
• This describes the size of the orbital
• The distance of an electron from an orbital is
  directly proportional to the energy of the
  electron

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Energy of a Hydrogen Atom

• The energy levels are similar to the rungs of a
  ladder
• As n increases the rungs become closer
• This means the energy differences between the
  levels is not constant
• Bohrs model only holds true for hydrogen

25
Quantum Mechanical Model

• 1925 Louis de Broglie proposed that particles of
  matter, including electrons, behaved as waves
• If an electron has wavelike motion, only certain
  frequencies, wavelengths and energies are
  possible
• He quantized them with the deBroglie equation
• ? h/mv

26
Heisenberg Uncertainly Principle

• Its impossible to know the velocity and the
  position of a particle at the same time
• Therefore, it is impossible to assign fixed paths
  to the orbits of the electrons
• We can only know the probability that an electron
  will occupy a certain region around the nucleus

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Schrodinger Equation

• Derived an equation which treated hydrogen's
  electron as a wave
• Created a model of the hydrogen atom which could
  be more readily expanded to explain other atoms
• The solution to this equation is known as a wave
  function it is related to the probability of
  finding an electron within a particular area
  around the nucleus

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Atomic Orbitals

• The wave function predicts a three dimensional
  area around the nucleus where an electron of a
  specific energy is likely to be found
• The orbital does not have a defined size, but it
  does have a particular shape

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Energy Sublevels

• Each principal Energy level as denoted by the
  Principal Quantum number, n, contains sublevels
• Principal level 1 has 1 sublevel
• Principal level 2 has 2 sublevels
• Principal level 3 has 3 sublevels

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Shapes of Orbitals

• Sublevels are labeled s, p, d or f according to
  the shape of the atoms orbitals
• S orbitals are spherical (there is 1)
• P orbitals are dumbell shaped (there are 3)
• D orbitals (there are 5) 4 of them are shaped
  like 4 leaf clovers, the 5th is a dumbell with a
  ring around the center
• F orbitals are complex multilobed shapes

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Hydrogen's electron - the 1s orbital
In the hydrogen case, the electron can be found
anywhere within a spherical space surrounding
the nucleus. The diagram shows a cross-section
through this spherical space.
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Each orbital has a name

• The orbital occupied by the hydrogen electron is
  called a 1s orbital. The "1" represents the fact
  that the orbital is in the energy level closest
  to the nucleus. The "s" tells you about the shape
  of the orbital. s orbitals are spherically
  symmetric around the nucleus

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2s Orbital
The orbital on the left is a 2s orbital. This is
similar to a 1s orbital except that the region
where there is the greatest chance of finding the
electron is further from the nucleus - this is an
orbital at the second energy level.
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3s, 4s (etc) orbitals get progressively further
from the nucleus

• 2s (and 3s, 4s, etc) electrons spend some of
  their time closer to the nucleus than you might
  expect. The effect of this is to slightly reduce
  the energy of electrons in s orbitals. The nearer
  the nucleus the electrons get, the lower their
  energy.

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p orbitals
A p orbital is rather like 2 identical balloons
tied together at the nucleus. The diagram on
the right is a cross-section through that
3-dimensional region of space.
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px, py and pz

• Unlike an s orbital, a p orbital points in a
  particular direction
• All levels except for the first level have p
  orbitals

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d and f orbitals

• For the moment, you need to be aware that there
  are sets of five d orbitals at levels from the
  third level upwards, but you probably won't be
  expected to draw them or name them. Apart from a
  passing reference, you won't come across f
  orbitals at all

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Pauli Exclusion Principle

• States no two electrons can occupy the same
  quantum state
• The two electrons that occupy an energy level
  must have opposite spins
• The direction of the spin is designated by s (-s
  or s)
• The 1s orbital with its two electons of
  opposite spin is illustrated

39
Hund's Rule

• Electrons fill low energy orbitals (closer to the
  nucleus) before they fill higher energy ones
• Where there is a choice between orbitals of equal
  energy, they fill the orbitals singly as far as
  possible

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The diagram (not to scale) summarizes the
energies of the orbitals up to the 4p level.
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Three Rules Regarding Electron Configurations

• Aufbau Principle-Electrons enter orbitals of
  lowest energy first
• Pauli Exclusion Principle-An atomic orbital may
  describe a maximum of two electrons
• Hunds Rule-When electrons occupy orbitals of
  equal energy one electron enters each orbital
  with spins parallel

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Filling the Orbitals

• The real oddity is the position of the 3d
  orbitals. They are at a slightly higher level
  than the 4s - and so it is the 4s orbital which
  will fill first, followed by all the 3d orbitals
  and then the 4p orbitals. Similar confusion
  occurs at higher levels, with so much overlap
  between the energy levels that the 4f orbitals
  don't fill until after the 6s, for example.
• http//intro.chem.okstate.edu/WorkshopFolder/Elect
  ronconfnew.html

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The Order for Configurations
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Writing Electron Configurations

• strategy start with hydrogen, and build the
  configuration one electron at a time (the Aufbau
  principle )
• fill subshells in order by counting across
  periods, from hydrogen up to the element of
  interest

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Quick Check

• 1) How many p orbitals are there
• 2) How many electrons can each p orbital hold?
• 3) How many electrons are in an element with the
  configuration 1s2 2s2 2p6
• 4) What does the Pauli Exclusion Principle state?
• 5) How many d orbitals are there

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Valence Electrons

• The electrons in the outermost orbital of the
  element
• Elements have certain charecteristics, or trends,
  based upon the number of valence electrons
• For the Group A elements, with the exception of
  Helium (He), the group number tells you the
  number of valence electrons

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Lewis Structure

• The Lewis Structure, or electron dot diagram, is
  a way of showing the number of valence electrons
  in an element
• Lewis structures help us to keep track of
  electrons when elements form ions and participate
  in reactions to form compounds

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Lewis Structure
50
Quick Check

• 1) What is a Lewis Structure?
• 2) How does the group number relate to the
  electrons in an element?
• 3) Silicon is in Group 4A and contains 14
  electrons. Please draw the electron dot diagram
  of silicon.
• 4) Draw a p orbital
• 5) How many electrons can a p orbital hold?

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• Page 160
• 21-25

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