Bohr

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Bohrs Model of the Atom

• Chem 11

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Bohrs Model

• Why dont the electrons fall into the nucleus?
• e- move like planets around the sun.
• They move 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

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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
• The colours can be seen when passed through a
  prism
• Every gas has a unique pattern (colour)

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

Fifth
Fourth
Third
Increasing energy
Second
First
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Line spectrum of various elements
Helium
Carbon
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Line spectrum of various elements
Carbon
Helium
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• Electrons orbiting closest to the nucleus are
  said to be in the lowest energy state called the
  ground state
• Atoms can absorb an amount of energy
• This promotes an electron to a higher energy
  level called the excited state
• This energy level is unstable and so the electron
  will fall back to its ground state
• When it does this, the excess energy will be
  emitted as light

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• n2 is the ground state
• Electron is promoted to n3 the excited state
• Electron falls back down and light is given off

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• When the e- falls from one energy level to
  another, an amount of energy is emitted as light
• This light emitted at specific wavelengths, which
  corresponds to our atomic spectra
• Each atom will have different electron jumps
  therefore emitting different amounts of energy as
  light
• This creates different line spectra for various
  elements

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Lets watch this video

• http//www.mhhe.com/physsci/chemistry/chang7/esp/f
  older_structure/pe/m1/s3/

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• http//www.mhhe.com/physsci/astronomy/applets/Bohr
  /applet_files/Bohr.html
• http//highered.mcgraw-hill.com/sites/0072482621/s
  tudent_view0/interactives.html

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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 e-.
• His theory could only explain hydrogen and not
  the more complex atoms

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The Quantum Mechanical Model

• Energy is quantized meaning 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.
• An equation has been developed 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
  equation describes several shapes.
• These are called atomic orbitals
• Orbitals are regions where there is a high
  probability of finding an e-

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S sublevel

• 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 sublevel

• 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- e- enter the lowest energy
  first.
• This causes difficulties because of the overlap
  of orbitals of different energies.
• Pauli Exclusion Principle- at most 2 e-per
  orbital with 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 e- occupy orbitals of equal
  energy they dont pair up until they have to .
• Lets determine the e- configuration for
  Phosphorus
• Need to account for 15 electrons

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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.

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• http//en.wikipedia.org/wiki/Periodic_table