2.3 Electron Arrangement

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2.3 Electron Arrangement

• 2.3.1 Describe the electromagnetic spectrum
• 2.3.2 Distinguish between a continuous spectrum
  and a line spectrum
• 2.3.3 Explain how the lines in the emission
  spectrum of hydrogen are related to electron
  energy levels
• 2.3.4 Deduce the electron arrangement for atoms
  and ions up to Z20

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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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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)
• Think of Noble gases

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Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Long Wavelength
Short Wavelength
Visible Light
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Wavelength and frequency
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How did he develop his theory?

• He used mathematics to explain the visible
  spectrum of hydrogen gas
• Lines are associated with the fall of an excited
  electron back down to its ground state energy
  level.
• http//www.mhhe.com/physsci/chemistry/essentialche
  mistry/flash/linesp16.swf

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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
Helium
Carbon
Continuous line spectrum
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Those who are not shocked when they first come
across quantum theory cannot possibly have
understood it. (Niels Bohr on Quantum Physics)
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Wavelengths and energy

• Understand that different wavelengths of
  electromagnetic radiation have different
  energies.
• cv?
• cvelocity of wave (2.998 x 108 m/s)
• v(nu) frequency of wave
• ?(lambda) wavelength

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• Bohr also postulated that an atom would not emit
  radiation while it was in one of its stable
  states but rather only when it made a transition
  between states.
• The frequency of the radiation emitted would be
  equal to the difference in energy between those
  states divided by Planck's constant.

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• Ehigh-Elow hv hc/?
• h3.983 x 10-13 Jsmol-1 Planks constant
• E energy of the emitted light (photon)
• v frequency of the photon of light
• ? is usually stated in nm, but for calculations
  use m.
• This results in a unique emission spectra for
  each element, like a fingerprint.
• electron could "jump" from one allowed energy
  state to another by absorbing/emitting photons of
  radiant energy of certain specific frequencies.

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• Energy must then be absorbed in order to "jump"
  to another energy state, and similarly, energy
  must be emitted to "jump" to a lower state.
• The frequency, v, of this radiant energy
  corresponds exactly to the energy difference
  between the two states.
• In order for the emitted energy to be seen as
  light the wavelength of the energy must be in
  between 380 nm to 750 nm

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For Hydrogen only!

• En -R/n2, where R is -1312 kJ/mol and n is
  principle quantum number (energy level)
• Example Calculate the energy required to ionize
  a mole of electrons from the 4th to the 2nd
  energy level in a hydrogen atom?
• E4 -1312 / 42 - 82 kJ
• E2 -1312 / 22 - 328 kJ
• E4 E2 - 82 kJ (- 328 kJ) 246 kJ

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• What is the wavelength of light emitted when
  electrons go from n4 to n2 ? Is it visible to
  our eyes?
• E hc/?, therefore ? hc/E
• ? (3.983 x 10-13 kJsmol-1)(2.998 x 108
  ms-1)/(246 kJmol-1)
• 4.85 x 10-7 m
• Convert to nm and see if its visible! (1 nm 1 x
  10-9 m)
• (4.85 x 10-7 m)( 1nm) 485 nm (Its probably the
  green line)
• 1 x 10-9 m

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Bohrs Triumph

• His theory helped to explain periodic law (the
  trends from the periodic table)
• Halogens (gp.17) are so reactive because it has
  one e- less than a full outer orbital
• Alkali metals (gp. 1) 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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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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Energy level populations

• Electrons found per energy level of the atom.
• The first energy level holds 2 electrons
• The second energy level holds 8 electrons (2 in s
  and 6 in p)
• The third energy level holds 18 electrons (2 in
  s, 6 in p and 10 in d) There is overlapping here,
  so when we do the populations there will be some
  changes.
• That is as far as this course requires us to go!

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Examples for group 1

• Li 2.1
• Na 2.8.1
• K 2.8.8.1

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A good sitehttp//www.chemguide.co.uk/basicorg/b
onding/orbitals.html
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Electron ConfigurationHL only

• 12.1.3 State the relative energies of s, p, d,
  and f orbitals in a single energy level
• 12.1.4 State the maximum number of orbitals in
  a given energy level.
• 12.1.5 Draw the shape of an s orbital and the
  shapes of px, py and pz orbitals
• 12.1.6 Apply the Aufbau principle, Hunds rule
  and the Pauli exclusion principle to write
  electron configurations for atoms and ions up to
  Z54.

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

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• Phosphorous, 15 e- to place
• 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.