Electron Configurations Chapter 5 Chemistry-CP

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Electron ConfigurationsChapter 5Chemistry-CP
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Radiant Energy

• The understanding of how electrons behave comes
  from studies of how light interacts with matter.
• Light carries energy through space in the form of
  waves and also in the form of extremely tiny,
  fast moving particles.
• Light has the properties of waves particles.

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Light as Waves

• Light waves are a form of electromagnetic
  radiation
• Electromagnetic radiation are energy waves

IMPORTANT ASPECTS OF WAVES Crest The top of a
wave Trough The bottom of a wave Amplitude
(A) The height of a waves crest (from origin
to crest) Wavelength (?) The distance between 2
consecutive crests or two consecutive
troughs Frequency (?) The number of waves that
pass a given point per second
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Waves

• Light travels at the speed of light.
• The speed of light is constant, which means it is
  always the same value 3.00 108 m/s
• Because light moves at a constant speed,
  wavelength frequency are inversely proportional
  as per the following equation.
• c ? ?
• c speed of light
• ? wavelength (lambda)
• ? frequency (measured in 1/s or s-1 or
  Hertz (Hz))

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c ? ?

• What is the frequency of a wave having a
  wavelength of 8.12 x 102 m?
• A helium neon laser produces red light whose
  wavelength is 633 nm. What is the frequency of
  this radiation?
• Calculate the wavelength of a radio wave with a
  frequency of 9.31 106 s-1.

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THE ELECTROMAGNETIC SPECTRUM
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Gamma Rays

• Generated by radioactive atoms, nuclear
  explosions and supernova explosions
• Can kill living cellsused for cancer treatment
• Used to sterilize medical equipment

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X-Rays

• Discovered by accident in 1895, when W.C.
  Roentgen shielded a cathode ray tube with black
  paper and found that a fluorescent light could be
  seen on a screen a few feet from the tube (first
  bone x-ray was of his wifes hand!)
• Electrons shot at an element (such as tungsten or
  molybdenum) with high energy can knock an
  electron out of that atom, producing x-rays
• Used for radiography, crystallography, astronomy,
  airport security

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Ultraviolet Radiation

• Gets its name from the fact that it consists of
  waves with frequencies higher than what humans
  associate with violet light
• Emitted from the sun, from black lights
• UV-B produces Vitamin D, too much DNA damage
  collagen fibers, can cause sunburn, may lead to
  cataracts
• Some animals, insects, birds and reptiles can see
  the near ultraviolet making certain flowers, etc.
  brighter to them.

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VISIBLE LIGHT

• Portion of the electromagnetic spectrum that is
  visible to the human eye
• ROYGBIVViolet has the highest frequency

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Infrared

• Below red
• Heat radiation
• Emitted from humans at normal body temperature
• Military use (surveillance, night vision, homing)
• Short ranged wireless communication, weather
  forecasting, remote temperature sensing
• Purple white light get on cheaper digital cameras
  (poor infrared filters)

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Microwaves

• Wireless LAN Bluetooth
• Radar Detectors, Air Traffic Control
• GPS
• The frequency of the waves used in microwave
  ovens, 2500 megahertz, targets water, sugar fat
  molecules
• Thin, sharp metals can not handle the electric
  current passing through them and may spark or
  start a fire
• Has never been conclusively shown that microwaves
  have biological effects
• http//www.youtube.com/watch?vUg8hSqkFUXY
• http//www.youtube.com/watch?vPIrd4172Czw

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Radio Waves

• Transport information through the atmosphere or
  space without wires
• AM FM Radio, TV transmission, mobile phones,
  military communications, wireless computer
  networks

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Visible Spectrum

• Part of the electromagnetic spectrum

• Continuous Spectrum One color fades gradually
  into the next.
• Different colors have different wavelengths.
• The color of visible light with the largest
  wavelength and lowest frequency is
• The color of visible light with the shortest
  wavelength and highest frequency is
• The brightness of visible light is determined by

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• Radiation with the largest wavelength and lowest
  frequency is
• Radiation with the shortest wavelength and
  highest frequency is
• Radiation with frequencies greater than visible
  light can pose health hazards because
• Radiation with frequencies lower than visible
  light are less harmful because

Aircraft Shipping Bands, Radio Waves
Gamma Rays
Have high enough energy to be capable of damaging
organisms
Do not have enough energy and pose no health
hazards
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What puzzled scientists about electromagnetic
radiation?

• Why do objects at different temperatures give off
  different color light?
• Why do different elements emit different colors
  when heated?

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Plancks Theory

• Suggested that the energy emitted or absorbed by
  an object is restricted to pieces of particular
  sizes called quanta.
• Substances can emit or absorb only certain
  amounts of energy (so only certain wavelengths)
• Showed that frequency and energy are directly
  proportional

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Plancks Theory
Plancks Theory

• E h ?
• h Plancks constant 6.626 10-34 J?s
• Joule (J) S.I. Unit for Energy

What is the approximate energy of Ultraviolet
Light? What is the energy of radiation with a
wavelength of 290 nm?
How much energy is contained in a wave with a
frequency of 2 x 1016 Hz?
What is the frequency of a wave with an energy of
2.90 x 1022 J?
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Why cant you see quantum effects in the everyday
world around you?

• Plancks constant is very small, therefore, each
  quantum of energy is very small
• Quanta are too small to see in the everyday world
• Atoms are very small, so in relation to the atom,
  quanta are significant

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The Photoelectric Effect

• Proposed by Albert Einstein
• In the photoelectric effect, electrons are
  ejected from the surface of a metal when the
  metal absorbs photons
• Photon Quantum of light (a tiny particle of
  light)

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The Photoelectric Effect

• When a photon strikes the surface of a metal, it
  transfers its energy to an electron in a metal
  atom
• If the energy of the photon is too small for the
  electron, the electron stays put
• If the photon has enough energy, the electron
  will escape the surface of the metal.

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The Photoelectric Effect

• Why is it easier for violet light (vs. red light)
  to cause the photoelectric effect?

Violet light has a higher frequency and,
therefore, more energy than red light.
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Light has a Dual Nature

• A photon behaves like a particle but always
  travels at the speed of light and has an
  associated frequency and wavelength
• In 1923, Arthur Compton showed that a photon
  could collide with an electron
• Light possesses the properties of both particles
  and waves

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How can atoms gain or lose energy?

• Atoms can only gain or lose energy in a quantum
• Take a look through your spectral tube at the
  emission tube at the front of the room.
• How does what youre looking at demonstrate the
  idea above?

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Line Spectrum

• A spectrum that contains only certain colors, or
  wavelengths

• When heat or electricity is passed through an
  atom, the atom absorbs the energy and then gives
  off that energy in the form of light
• The emitted light is unique for every element
• Atomic Emission Spectrum An atomic fingerprint
  showing the emission line spectrum of that atom
• Useful in identifying an element

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NIELS BOHR1913PLANETARY MODEL OF THE ATOM
Electrons move in defined orbits around the
nucleusjust as the planets move around the sun.
Orbit Region outside the nucleus where
electrons are found
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Bohrs Postulate

• Was applicable only to hydrogen
• Able to show that electrons move to higher energy
  levels (excited states) when they absorb
  radiation.
• Electrons will immediately return to the lower
  energy levels (ground state) by emitting energy
  of a specific wavelength

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Light has a Dual Nature

• When light travels through space, it acts like a
  wave
• When light interacts with matter, it acts like a
  particle
• De Broglie predicted matter waves--that matter
  should behave like waves and exhibit a wavelength
• Clinton Davisson Lester Germer proved that
  electrons (believed to be particles) were
  reflected from a matter like waves
• Mass of an object must be very small in order to
  observe its wavelength

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Heisenberg Uncertainty Principle

• An electron is located by striking that electron
  with a photon which bounces back to a detection
  device
• The electron is so small in mass that the
  electron is moved by the collision
• Proved a problem with Bohrs model You cannot
  think of electrons moving in defined paths
  because there is no way to prove the electrons
  follow defined paths

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MODELS OVER TIME
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Quantum Mechanical Modela.k.a Wave Model

• Explains the properties of atoms by treating the
  electron as a wave that has quantized its energy
• Does not describe exact positions of the
  electrons instead describes the probability that
  electrons will be found in certain locations
  around the nucleus

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Electron Cloud
An illustration that uses a blurry cloud to
illustrate the probability of finding an electron
in various locations around the nucleus.
(Determined by wave functions electron density
charts)
Areas of high electron density are the most
probable locations of the electrons.
Areas of low electron density are the least
probable locations of the electrons.
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Atomic Orbitals

• Region of space where the electron is located
• Have characteristic shapes, sizes and energies
• Do not describe how the electron actually moves
• The orbital occupied is determined by the amount
  of energy of an electron

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s-Orbital
The s-orbital consists of 1 orbital on all 3
axes 1 orbital has a maximum of 2 electrons
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p-Orbital
The p-orbital can exist on 3 different axes (x, y
and z). Therefore there are 3 p orbitals. The
p-sublevels 3 orbitals can hold a maximum of 6
electrons (2 on each of the 3 orbitals).
A p-orbital has a dumbbell shape
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d-Orbital
There are 5 different orientations of a
d-orbital. The d-sublevels 5 orbitals can
hold a maximum of 10 electrons (2 electrons on
each orbital).
A d-orbital has a cloverleaf shape
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f-Orbital
An f-orbital has a complex shape
There are 7 different orientations of the
f-orbital. The f-sublevel can hold a maximum
of 14 electrons (2 for each orbital).
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Energy Orbitals

• Energy of electrons are quantized (exact)
• Principal Energy Levels or Principal Quantum
  Number designates the distance of the electron
  from the nucleus

Principal energy levels are divided into sublevels
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Sublevels
Sublevels of the atom are designated s, p, d f
The number of the energy level tells you how many
sublevels are present within that sublevel.
Another words Energy Level 1 has __________
Sublevel Energy Level 2 has __________
Sublevels Energy Level 3 has __________
Sublevels Energy Level 4 has __________
Sublevels The electrons address consists of its
principal energy level, its sublevel, and its
electrons within that sublevel
1
2
3
4
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SUBLEVEL s
Orbital Shape Max of electrons Region on Periodic Table Orbital Models
1 orbital s Sphere 2 Groups 1 2 (1st tower)
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SUBLEVEL p
Orbital Shape Max of electrons Region on Periodic Table Orbital Models
3 orbitals px py pz dumbbell 6 Groups 13-18 (2nd tower)
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SUBLEVEL d
Orbital Shape Max of electrons Region on Periodic Table Orbital Models
5 orbitals dxy dxz dyz dx2-y2 dz2 cloverleaf 10 Groups 3-12 (transition metals)
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SUBLEVEL f
Orbital Shape Max of electrons Region on Periodic Table Orbital Models
7 orbitals complex 14 Bottom 2 rows (inner-transition metals)
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Some Atomic Models
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More Models
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Example
Beryllium ______ protons, ______ electrons E-
Configuration 1s22s2
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Example
Oxygen ______ protons, ______ electrons E-
Configuration 1s22s22p4
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PRACTICE PROBLEMS

• Electron configurations for 3 different elements
  are given below. Draw the atomic model of each
  element and then identify the element.
• Examples 1s22s1 1s22s22p3 1s22s22p63s23p4
• 1) 1s22s22p1 2) 1s2 3) 1s22s22p63s1

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Example
Boron ______ protons, ______ electrons E-
Configuration 1s22s22p1
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Examples

• Helium ______ protons, ______ electrons
• E- Configuration 1s2

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Examples

• Sodium ______ protons, ______ electrons
• E- Configuration 1s22s22p63s1

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Electron Spin

• Electrons spin either clockwise or
  counterclockwise
• The spinning creates a magnetic field
• Clockwise is like a magnet whose north pole is
  pointing up
• Counterclockwise behaves like a magnet whose
  north pole is pointing down
• Parallel Spins result in a net magnetic effect
• Opposite Spins cancel each other out

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Pauli Exclusion Principle
-1925, Austrian physicist-Wolfgang Pauli -States
that each orbital in an atom can hold at most 2
electrons and that these electrons must have
opposite spins (or be paired).
Sublevels Orbitals Max of e- s 1
2 p
3 6 d
5 10
f 7
14
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Electron Configuration

• The addresses of an atoms electrons
• Determined by distributing the atoms electrons
  among levels, sublevels and orbitals based on a
  set of principles
• Orbitals from lowest to highest energy
• s? p ? d ? f
• Ground State The electrons are in the lowest
  energy levels available

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How do electrons occupy energy levels?

• Aufbau Principle Electrons are added one at a
  time to the lowest energy orbitals available
  until all the electrons are accounted for
• Pauli Exclusion Principle An orbital can hold a
  maximum of 2 electrons that must spin in opposite
  directions
• Hunds Rule Electrons occupy equal-energy
  orbitals so that a maximum number of unpaired
  electrons results

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Orbital Diagrams
4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
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What happens when an element in its ground state
is supplied with electricity or heat?

• Electrons may move to the excited state.
• Excited State Energy level attained when an
  electron absorbs energy and jumps to a higher
  energy level

Ground State
Excited State
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For each pair of orbital diagrams below, which
represents the ground state and which represents
the excited state of that atom?
4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
Scandium
4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
Magnesium
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What happens to the excited electron?
http//www.meta-synthesis.com/webbook/11_five/five
04.jpg
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Exceptions to the Aufbau Rule

• A half-full or full d sublevel will increase an
  atoms stability
• An electron may be removed from the s sublevel to
  create a full or half full d sublevel

4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
4p ____ ____ ____ 3d ____ ____ ____ ____
____ 4s ____ 3p ____ ____ ____ 3s ____ 2p
____ ____ ____ 2s ____ 1s ____
Cr
Cu
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Groups (also called Families)

• The vertical columns on the periodic table
• There are 18 groups, labeled with the numbers
  1-18.

1
18
2
15
17
16
14
13
3
4
5
7
6
9
8
12
11
10
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Corresponding Regions on the Periodic Table
p-block Representative Elements Groups
13-18 Begins with Principal Energy Level
2 Contains 6 elements because each p sublevel can
hold up to 6 electrons
s-block Representative Elements Groups 1
2 Begins with Principal Energy Level 1 Contains 2
elements because each s sublevel can hold 2
electrons
f-block Inner Transition Metals lanthanides
actinides (bottom 2 rows) Begins with Principal
Energy Level 4 Contains 14 elements because each
f sublevel can hold up to 14 electrons
d-block Transition Metals Groups 3-12 Begins
with Principal Energy Level 3 Contains 10
elements because each d sublevel can hold up to
10 electrons
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He
N
Ti
I
Fr
Ce
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Noble Gas Configuration Uses the symbol of the
noble gas in brackets to represent the inner
level electrons of an atom.
1s
2s
2p
3s
3p
3d
4p
4s
4d
5p
5s
Cd
5d
6p
6s
Ba
7p
7s
6d
4f
U
5f
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VALENCE ELECTRONS

• The electrons in the outermost energy level.
• Remember, the number in front of the sublevel
  indicates the energy level
• 1s22s22p6
• Sofind the highest energy level and add up all
  the electrons in that level.

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EXAMPLES

• Calcium
• Aluminum
• Iodine
• Oxygen
• Iron

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ENERGY
Electrons with the most energy are located
farthest from the nucleus
Electrons with the lowest energies are located
close to the nucleus.
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Quantum
A quantum is the specific amount of energy needed
for an electron to move between energy levels.