Where do the electrons go

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Where do the electrons go?
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What do we know about electrons?

• They are small (1/1840 of an amu)
• They carry a negative charge
• They can get excited
• They move around the nucleus

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Excited Electrons emit light
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Whats light got to do with it?

• When studying elements, scientist found that when
  certain elements were heated in a flame those
  elements emitted varying colors of light.
• Light carries energy
• Something in the atom must be carrying energy in
  order to emit light

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Electromagnetic Spectrum

• Visible light is part of the electromagnetic
  spectrum.
• Light travels as a wave

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Light as a Wave

• Wave characteristics include wavelength (nm),
  amplitude, frequency (1/s)

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Light as a Wave

• The speed of light can be calculated.
• Speed of light wavelength x frequency
• ALL electromagnetic waves travel at the same
  speed (this includes visible light) which is
  3.00 x 108 m/s
• 3.00 x 108 m/s wavelength x frequency

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Light as a Wave

• Wavelength and frequency are inversely
  proportional (as one increases the other
  decreases)

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Electrons emitting light
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Electrons emitting light

• Iron when heated first glows RED, then ORANGE,
  then at even hotter temperatures it glows BLUE.
• Remember that temperature (heat) is the measure
  of the average kinetic energy of particles.
• What do you think is moving?
• Electrons, the change in color of glowing iron is
  the direct result of an increase temperature

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Electrons as a particle of light

• Excited electrons can absorb only a certain
  amount of energy.
• The energy absorbed or emitted by electrons is
  called a quantum of energy
• AKA small packets of energy
• These quanta of energy are being absorbed by the
  iron. The more energy absorbed by the electron
  produces the color change

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The light is a direct result of exited electrons

• Max Planck created an equation that quantized the
  amount of energy in a quanta
• Equantum hv
• E is energy
• H is Plancks constant
• V is frequency (from the speed of light equation
  Cwavelength x frequency)

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Connecting Light to Electrons

• Atomic Emission Spectra- The colors (wavelengths)
  of light that are emitted by an element when it
  is heated.
• The light is emitted as a photon (small packet of
  light)
• Each element has a different atomic emission
  spectra because each element contains electrons
  with different energies.

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Atomic Emission Spectrum
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Quantum Theory

• If the atom has not absorbed any energy it is
  said to be in the GROUND state. If has absorbed
  energy it is said to be in an EXCITED state.

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Quantum Theory

• The amount of energy an atom could absorb is
  directly related to the atoms electrons.
• Thus the energy level is BOHRn
• The place where an electron can be found is
  called an orbit (Bohr described it as a circular
  ring around the nucleus)

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Bohr Models of the Atom

• The number of rings around the nucleus is equal
  to the period the element is located in.
• Ex Na is in period 3 and will have 3 orbits
  (rings) around it.
• Orbit 1 can hold 2 e-
• Orbit 2 can hold 8 e-
• Orbit 3 can hold 18 e-
• Orbit 4 can hold 32 e-

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Drawing Bohr Models

• Chlorine
• Boron
• Ca
• He

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Lets Talk Orbits

• Orbits represent energy amounts that the
  electrons have
• Electrons can be excited and move up an orbit
  (which means it gains energy)
• The orbits are NOT equal distances apart. Some
  are closer than others.

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Now that I am really smart

• Can I tell you exactly where an electron is at
  any given moment as it spins around the nucleus
  of an atom?
• NO. There is this thing called the Heisenberg
  Uncertainty Principle. IT IS IMPOSSIBLE to know
  precisely both the momentum and location of an
  electron at the same time.
• I can get real close though.

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Where is the electron?

• The probably location of an electron is a code of
  4 quantum numbers (aka- letters) that creates a
  3-D picture of the atom
• All the moving electron around the nucleus of an
  atom creates the electron cloud

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Quantum Numbers

• Principal Quantum Number (n) represents the
  energy level
• Principle Quantum Number (n)-indicates the size
  of the electron cloud and is commonly referred to
  as the energy level.
• Each energy level has a maximum number of
  electrons. (2n2)

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Principle Quantum Number
Energy Level Maximum of
Electrons 1 2 2 8 3
18 4 32 5 50 6
72
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Subsidiary Quantum Number

• Energy Sublevels or subshells- each energy level
  is subdivided into types of orbitals based on
  shape (s, p, d, f)

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Subsidiary Quantum Number
Energy Level Sublevel Maximum of
electrons 1 s
2 2 s, p 2,
6 3 s, p, d 2, 6, 10 4 s, p, d, f
2, 6, 10, 14
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Magnetic Quantum Number
Each sublevel contains a certain number of
orbitals ( of orbitals n2). Orbitals can hold
only 2 e-. The number of electrons can be
determined by 2n2.
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Magnetic Quantum Number
Sublevel of orbitals s 1
p 3 d 5
f 7
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Spin Quantum Number

• Each electron has either a or spin.
• The electron either rotates clockwise or counter
  clockwise.

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Electron Configurations

• Describe the probably location of electrons of an
  atom in its ground state.
• Three rules govern how electrons fill energy
  levels and orbitals

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Aufbau principle

• Electrons occupy the lowest energy orbital
  available.
• The filling of the orbitals is not always in
  order. You must follow the Aufbau diagram to
  determine order.

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The Pauli Exclusion Principle

• ONLY TWO electron can fill any orbital and their
  spins must be opposite to offset their repulsion

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Hunds Rule

• Fill orbitals of the same energy amount with one
  electron each then fill with the second electron.

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Ways to write Electron Configurations

• Orbital diagrams use arrows that represent each
  electron. (most detailed notation)
• Electron Notation uses superscripts to represent
  the electrons in each sublevel
• Noble Gas Notation- uses the closest Noble gas
  prior to (before) the element being configured.
  The Noble gas is placed in to indicate that
  all is the same to this point as the Noble Gas
  then electron notation is used

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Exceptions as always

• Half filled orbitals are more stable
• Chromium has a full 4s and 4 electron in its 3d,
  this is unstable. To stabilize the atom an
  electron is moved from the 4s to the 3d so that
  both sublevels are not half filled
• Copper has 9 electrons in its 3d. It takes one
  electron from the 4s and places it in the 3d.
  Now 4s is half-filled and stable while 3d is full
  and stable.

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Valence Electrons

• Valence electrons are electrons that are in the
  outermost orbitals which will always be the
  highest energy level orbital.
• Valence electrons are found only in the s or p
  sublevel.
• The number of valence electrons will never exceed
  8
• How many valence electrons do the following atoms
  have?
• Na, O, Kr, B, and Mg

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Valence Electrons

• Electron Dot Structures are used to denote only
  valence electrons
• Dots are placed around the four sides of the
  elements symbol. Place one dot on each side
  before placing a second dot on each side.
• Remember there will never be more than 8 dots.