Bohr model and electron configuration

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Bohr model and electron configuration

• Mrs. A. Kay
• Chem 11

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Bohrs Model

• Why dont the electrons fall into the nucleus?
• Move like planets around the sun.
• In circular orbits at different levels.
• Amounts of energy separate one level from
  another.

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Bohrs Model
Nucleus
Nucleus
Electron
Electron
Orbit
Orbit
Energy Levels
Energy Levels
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Bohr postulated that

• Fixed energy related to the orbit
• Electrons cannot exist between orbits
• The higher the energy level, the further it is
  away from the nucleus
• An atom with maximum number of electrons in the
  outermost orbital energy level is stable
  (unreactive)

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How did he develop his theory?

• He used mathematics to explain the visible
  spectrum of hydrogen gas
• http//www.mhhe.com/physsci/chemistry/essentialche
  mistry/flash/linesp16.swf

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Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Long Wavelength
Short Wavelength
Visible Light
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The line spectrum

• electricity passed through a gaseous element
  emits light at a certain wavelength
• Can be seen when passed through a prism
• Every gas has a unique pattern (color)

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Line spectrum of various elements
Helium
Carbon
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Bohrs Triumph

• His theory helped to explain periodic law
• Halogens are so reactive because it has one e-
  less than a full outer orbital
• Alkali metals are also reactive because they have
  only one e- in outer orbital

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Drawback

• Bohrs theory did not explain or show the shape
  or the path traveled by the electrons.
• His theory could only explain hydrogen and not
  the more complex atoms

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• Further away from the nucleus means more energy.
• There is no in between energy
• Energy Levels

Fifth
Fourth
Third
Increasing energy
Second
First
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The Quantum Mechanical Model

• Energy is quantized. It comes in chunks.
• A quanta is the amount of energy needed to move
  from one energy level to another.
• Since the energy of an atom is never in between
  there must be a quantum leap in energy.
• Schrödinger derived an equation that described
  the energy and position of the electrons in an
  atom

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Atomic Orbitals

• Principal Quantum Number (n) the energy level
  of the electron.
• Within each energy level the complex math of
  Schrödinger's equation describes several shapes.
• These are called atomic orbitals
• Regions where there is a high probability of
  finding an electron

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S orbitals

• 1 s orbital for
• every energy level
• 1s 2s 3s
• Spherical shaped
• Each s orbital can hold 2 electrons
• Called the 1s, 2s, 3s, etc.. orbitals

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P orbitals

• Start at the second energy level
• 3 different directions
• 3 different shapes
• Each orbital can hold 2 electrons

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• The p Sublevel has 3 p orbitals

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The D sublevel contains 5 D orbitals

• The D sublevel starts in the 3rd energy level
• 5 different shapes (orbitals)
• Each orbital can hold 2 electrons

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The F sublevel has 7 F orbitals

• The F sublevel starts in the fourth energy level
• The F sublevel has seven different shapes
  (orbitals)
• 2 electrons per orbital

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Summary
Starts at energy level
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Electron Configurations

• The way electrons are arranged in atoms.
• Aufbau principle- electrons enter the lowest
  energy first.
• This causes difficulties because of the overlap
  of orbitals of different energies.
• Pauli Exclusion Principle- at most 2 electrons
  per orbital - different spins

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Electron Configurations

• First Energy Level
• only s sublevel (1 s orbital)
• only 2 electrons
• 1s2
• Second Energy Level
• s and p sublevels (s and p orbitals are
  available)
• 2 in s, 6 in p
• 2s22p6
• 8 total electrons

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• Third energy level
• s, p, and d orbitals
• 2 in s, 6 in p, and 10 in d
• 3s23p63d10
• 18 total electrons
• Fourth energy level
• s,p,d, and f orbitals
• 2 in s, 6 in p, 10 in d, and 14 in f
• 4s24p64d104f14
• 32 total electrons

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Electron Configuration

• Hunds Rule- When electrons occupy orbitals of
  equal energy they dont pair up until they have
  to .

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• The first to electrons go into the 1s orbital
• Notice the opposite spins
• only 13 more

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• The next electrons go into the 2s orbital
• only 11 more

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• The next electrons go into the 2p orbital
• only 5 more

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• The next electrons go into the 3s orbital
• only 3 more

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• The last three electrons go into the 3p orbitals.
• They each go into separate shapes
• 3 unpaired electrons
• 1s22s22p63s23p3

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Orbitals fill in order

• Lowest energy to higher energy.
• Adding electrons can change the energy of the
  orbital.
• Half filled orbitals have a lower energy.
• Makes them more stable.
• Changes the filling order

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Write these electron configurations

• Titanium - 22 electrons
• 1s22s22p63s23p64s23d2
• Vanadium - 23 electrons 1s22s22p63s23p64s23d3
• Chromium - 24 electrons
• 1s22s22p63s23p64s23d4 is expected
• But this is wrong!!

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Chromium is actually

• 1s22s22p63s23p64s13d5
• Why?
• This gives us two half filled orbitals.
• Slightly lower in energy.
• The same principal applies to copper.

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Coppers electron configuration

• Copper has 29 electrons so we expect
• 1s22s22p63s23p64s23d9
• But the actual configuration is
• 1s22s22p63s23p64s13d10
• This gives one filled orbital and one half filled
  orbital.
• Remember these exceptions

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Great site to practice and instantly see results
for electron configuration.
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Practice

• Time to practice on your own filling up electron
  configurations.
• Do electron configurations for the first 20
  elements on the periodic table.