Electrons in Atoms

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• Electrons in Atoms

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Light

• Light is a kind of electromagnetic radiation.
• All forms of electromagnetic radiation move at
  3.00 x 108 m/s.

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Parts of a Wave
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Parts of a Wave

• Origin the baseline of a wave
• Crest - high point on a wave
• Trough - low point on a wave
• The amplitude of a wave is the waves height from
  the origin to a crest, or from the origin to a
  trough.

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Wavelength

• Wavelength (represented by ?, the Greek letter
  lambda) is the shortest distance between
  equivalent points on a continuous wave.
• Wavelength - distance from crest to crest or
  trough to trough
• Wavelength is usually expressed in meters (m).

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Frequency

• Frequency (represented by f ) is the number of
  waves that pass a given point per second.
• Units are cycles/sec or hertz (Hz)

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Relationship Between Frequency and Wavelength

• c f l

• c the speed of light

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Frequency and Wavelength

• They are inversely related, which means that as
  one goes up the other goes down.

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Frequency and Wavelength

• Different frequencies of light correspond to
  different colors of light.

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The Electromagnetic Spectrum
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X-Rays
Radiowaves
Microwaves
Infrared .
Ultra-violet
GammaRays
Long Wavelength
Short Wavelength
Visible Light
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The Quantum Concept

• In 1900, the German physicist Max Planck began
  searching for an explanation as he studied the
  light emitted from heated objects.

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The Quantum Concept

• Matter can gain or lose energy only in small,
  specific amounts called quanta.
• That is, a quantum is the minimum amount of
  energy that can be gained or lost by an atom.

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• That is, while a beam of light has many wavelike
  characteristics, it also can be thought of as a
  stream of tiny particles, or bundles of energy,
  called photons.
• Thus, a photon is a particle of electromagnetic
  radiation with no mass that carries a quantum of
  energy.

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• Planck went further and demonstrated
  mathematically that the energy of a quantum is
  directly related to the frequency of the emitted
  radiation.

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Energy and Frequency

• E h f
• E energy of the photon (J Joules)
• f frequency (Hz)
• h is Plancks constant
• h 6.63 x 10-34 Joules.sec

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Energy and Frequency

• Looking at the equation, you can see that the
  energy of radiation increases as the radiations
  frequency, f, increases.

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The Photoelectric Effect

• Scientists knew that the wave model of light
  could not explain a phenomenon called the
  photoelectric effect.

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• In the photoelectric effect, electrons, called
  photoelectrons, are emitted from a metals
  surface when light of a certain frequency shines
  on the surface.

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• Einstein proposed that for the photoelectric
  effect to occur, a photon must possess, at a
  minimum, the energy required to free an electron
  from an atom of the metal.

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STOP HERE
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The Bohr Model of the Atom

• Niels Bohr, a young Danish physicist working in
  Rutherfords laboratory in 1913, suggested that
  the single electron in a hydrogen atom moves
  around the nucleus in only certain allowed
  circular orbits.

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The Bohr Model of the Atom

• The atom looked like a miniature solar system.
  The nucleus is represented by the sun, and the
  electrons act like the planets.

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Bohrs Model

• The orbits are circular and are at different
  levels.
• Amounts of energy separate one level from another.

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Modern View

• The atom has two regions and is 3- dimensional.
• The nucleus is at the center and contains the
  protons and neutrons.

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Modern View

• The electron cloud is the region where you might
  find an electron and most of the volume of an
  atom.

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Lets look at the Bohr Model in greater detail
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Bohrs Model

• Bohr proposed that electrons must have enough
  energy to keep them in constant motion around the
  nucleus.
• Electrons have energy of motion that enables them
  to overcome the attraction of the positive
  nucleus.

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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.
• Electrons reside in energy levels.

Fourth
Third
Increasing energy
Second
First
Nucleus
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The QuantumMechanical Model

• Building on Plancks and Einsteins concepts of
  quantized energy (quantized means that only
  certain values are allowed), Bohr proposed that
  the hydrogen atom has only certain allowable
  energy states.

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The QuantumMechanical Model

• The lowest allowable energy state of an atom is
  called its ground state.
• When an atom gains energy, it is said to be in an
  excited state.

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The QuantumMechanical Model

• When the atom is in an excited state, the
  electron can drop from the higher-energy orbit to
  a lower-energy orbit.
• As a result of this transition, the atom emits a
  photon corresponding to the difference between
  the energy levels associated with the two orbits.

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Atomic Spectrum

• What Color Tells Us About Atoms

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Atomic Spectrum

• By heating a gas of a given element with
  electricity, we can get it to give off colors.

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Atomic Spectrum

• Each element gives off its own characteristic
  colors.
• The spectrum can be used to identify the atom.

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• These are called line spectra.
• Each is unique to an element.

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• The spectrum of light released from excited atoms
  of an element is called the emission spectrum of
  that element.

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Changing the Energy

• As the electrons fall from the excited state,
  they release energy in the form of light.

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Ultraviolet
Visible
Infrared

• The further they fall, the greater the energy.
• This results in a higher frequency.

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Question

• Use the Chemistry Reference Tables to answer the
  following
• An electron falls from energy level 5 to energy
  level 3. What is the wavelength of the light
  emitted?

(1282 nm)
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Question

1. An electron falls from energy level 6 to energy
   level 2. What is the wavelength of the light
   emitted?

(410 nm)
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Question

1. An electron falls from energy level 3 to energy
   level 1. What type of electromagnetic radiation
   is emitted (infrared, visible or ultraviolet)?

(UV)
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Question

1. An electron falls from energy level 4 to energy
   level 2. What type of electromagnetic radiation
   is emitted (infrared, visible or ultraviolet)?

(visible)
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Question

1. An electron falls from energy level 5 to energy
   level 2. What color of visible light is emitted?

(blue)
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Question

1. An electron falls from energy level 3 to energy
   level 2. What color of visible light is emitted?

(red)
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The QuantumMechanical Model

• Like Bohrs model, Schrodingers quantum
  mechanical model limits an electrons energy to
  certain values.

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• The space around the nucleus of an atom where the
  atoms electrons are found is called the electron
  cloud.

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• A three-dimensional region around the nucleus
  called an atomic orbital describes the electrons
  probable location.

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Energy Levels

• In general, electrons reside in principal energy
  levels.

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Energy Levels

• As the energy level number increases, the orbital
  becomes larger, the electron spends more time
  farther from the nucleus, and the atoms energy
  level increases.

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Sublevels

• Principal energy levels contain energy sublevels.
  
• Principal energy level 1 consists of a single
  sublevel, principal energy level 2 consists of
  two sublevels, principal energy level 3 consists
  of three sublevels, and so on.

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Sublevels

• Sublevels are labeled s, p, d, or f.
• The s sublevel can hold 2 electrons, the p
  sublevel can hold 6 electrons, the d sublevel can
  hold 10 electrons, and the f sublevel can hold 14
  electrons.

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s block
p block
d block
f block
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Orbitals

• Sublevels contain orbitals.
• Each orbital may contain at most two electrons.

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s orbitals

• One s orbital
• for every
• energy level
• Spherical
• shaped
• Called the 1s, 2s, 3s, etc orbitals

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p orbitals

• Start at the second energy level
• 3 different directions
• 3 different dumbbell shapes

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p Orbitals
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d orbitals

• Start at the third energy level
• 5 different shapes

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f orbitals

• Start at the fourth energy level
• Have seven different shapes

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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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s block
p block
d block
f block
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s- block
s2
s1
s2

• Really have to include helium.
• Helium has the properties of the noble gases.

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Transition Metals - d block
d4
d9
d1
d2
d3
d5
d6
d7
d8
d10
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The p-block
p1
p2
p6
p3
p4
p5
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f - block

• inner transition elements

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1 2 3 4 5 6 7

• Each row (or period) is the energy level for s
  and p orbitals.

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• d orbitals fill up after previous energy level so
  first d is 3d even though its in row 4.

1 2 3 4 5 6 7
3d
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1 2 3 4 5 6 7
4f 5f

• f orbitals start filling at 4f

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Electron Configurations

• Electron configurations represent the way
  electrons are arranged in atoms.
• Aufbau principle - Electrons enter the lowest
  energy first.

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Electron Configurations

• This causes difficulties because of the overlap
  of orbitals of different energies.
• At most there can be only 2 electrons per
  orbital, and they must have opposite spins.

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Electron Configuration

• Hunds Rule - When electrons occupy orbitals of
  equal energy, they dont pair up with an electron
  of opposite spin until they have to.

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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 total

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2

• 12 electrons total

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2

• 20 electrons total

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Fill from the bottom up following the arrows

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

• 38 electrons total

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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 total

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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 total

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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 total

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Electron Configuration

• Lets determine the electron configuration for
  phosphorus.

• The atomic number of phosphorus is 15, so we need
  to account for 15 electrons.

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Electron Configuration

• 1s2

2s2
2p6 3s2
3p3

• 2 electrons

• 4 electrons

• 12 electrons

• 15 electrons

• The 4s sublevel does not need to be used to get
  the 15 electrons.

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Electron Configuration

• Lets determine the electron configuration for
  chromium.

• The atomic number of chromium is 24, so we need
  to account for 24 electrons.

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Electron Configuration

• 1s2

2s2
2p6 3s2
3p6 4s2
3d4

• 2 electrons

• 4 electrons

• 12 electrons

• 20 electrons

• 24 electrons

• The 4p and 5s sublevels do not need to be used to
  get the 24 electrons.

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Question

• Write the electron configuration for aluminum
  (Al).

(1s2 2s2 2p6 3s2 3p1)
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Question

• Write the electron configuration for neon (Ne).

(1s2 2s2 2p6)
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Question

• Write the electron configuration for calcium (Ca).

(1s2 2s2 2p6 3s2 3p6 4s2)
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Question

• Write the electron configuration for iron (Fe).

(1s2 2s2 2p6 3s2 3p6 4s2 3d6)
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Question

• Write the electron configuration for bromine (Br).

(1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5)
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Example

• Identify the element with the following electron
  configuration 1s2 2s2 2p6 3s1

• Add the superscript numbers together and find the
  element with that atomic number.

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Example

• 1s2 2s2 2p6 3s1
• 2 2 6 1 11

• Element 11 is sodium (Na).

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Question

• Identify the element with the following electron
  configuration 1s2 2s2 2p6 3s2 3p4

(sulfur - S)
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Question

• Identify the element with the following electron
  configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d9

(copper - Cu)
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Question

• Identify the element with the following electron
  configuration 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2

(germanium - Ge)
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Electron Configuration Using a Noble Gas
Abbreviation

• In order to write this type of configuration,
  find the noble gas (from Group 8A) that comes
  before the element in question.

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Electron Configuration Using a Noble Gas
Abbreviation

• Put the symbol for the noble gas in brackets and
  then write the part of the configuration that
  follows to reach the desired element.

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Example

• Write the electron configuration using a noble
  gas abbreviation for magnesium (Mg).

• Neon is the noble gas that proceeds magnesium.

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Example

• Put neons symbol in brackets.

• Ne

• Now use the periodic table to determine the rest
  of the configuration.

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s block
p block
1
2
d block
3
4
5
6
7
f block
Neon is in light blue magnesium is in bright
yellow.
The noble gas electron configuration for
magnesium is Ne 3s2
The additional configuration is 3s2.
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Example

• Write the electron configuration using a noble
  gas abbreviation for nickel (Ni).

• Argon is the noble gas that proceeds nickel.

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Example

• Put argons symbol in brackets.

• Ar

• Now use the periodic table to determine the rest
  of the configuration.

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p block
s block
1
2
d block
3
4
5
6
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The noble gas electron configuration for nickel
is Ar 4s2 3d8
f block
Argon is in light blue nickel is in bright
yellow.
The additional configuration is 4s2 3d8.
(Remember you subtract 1 from the d sublevel row
number.)
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Question

• Write the electron configuration using a noble
  gas abbreviation for fluorine (F).

(He 2s2 2p5)
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Question

• Write the electron configuration using a noble
  gas abbreviation for silicon (Si).

(Ne 3s2 3p2)
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Question

• Write the electron configuration using a noble
  gas abbreviation for zirconium (Zr).

(Kr 5s2 4d2)
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Valence Electrons

• The electrons in the outermost energy level are
  called valence electrons.

• You can also use the periodic table as a tool to
  predict the number of valence electrons in any
  atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18.

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Valence Electrons

• All atoms in Group 1, like hydrogen, have one
  valence electron. Likewise, atoms in Group 2 have
  two valence electrons.

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Valence Electrons

• All atoms in Group 13 have three valence
  electrons.
• All atoms in Group 14 have four valence
  electrons.
• All atoms in Group 15 have five valence electrons.

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Valence Electrons

• All atoms in Group 16 have six valence electrons.
• All atoms in Group 17 have seven valence
  electrons.
• All atoms in Group 18 have eight valence
  electrons, except helium which only has two.

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Valence Electrons

• All atoms in sublevels d and f have 2 valence
  electrons.

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Question

• How many valence electrons does each of the
  following elements have?
• carbon (C)

(4)
b) bromine (Br)
(7)
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Question

• c) iron (Fe)

(2)
d) potassium (K)
(1)
e) aluminum (Al)
(3)
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Lewis Dot Diagrams

• Because valence electrons are so important to the
  behavior of an atom, it is useful to represent
  them with symbols.

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Lewis Dot Diagrams

• A Lewis dot diagram illustrates valence electrons
  as dots (or other small symbols) around the
  chemical symbol of an element.

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• Each dot represents one valence electron.

• In the dot diagram, the elements symbol
  represents the core of the atomthe nucleus plus
  all the inner electrons.

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Electron (Lewis) Dot Diagrams

• Write the symbol.

X

• Put one dot for each valence electron.

• Dont pair electrons up until you have to.

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Question
Write a Lewis dot diagram for chlorine.
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Question
Write a Lewis dot diagram for calcium.
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Question
Write a Lewis dot diagram for potassium.