Electronic structure in atoms

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Electronic structure in atoms

• AH Chemistry, Unit 1(a)

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• Bohrs theory cannot explain spectra for atoms
  more complex than hydrogen, which show sub-lines.
• In 1926 Erwin Schrodinger devised a theory to
  describe more complicated atoms quantum
  mechanics.
• This classifies regions of atoms into SHELLS,
  SUB-SHELLS and ORBITALS.

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Quantum numbers

• Each electron in an atom is described by 4
  different quantum numbers
• n, l, ml, ms
• The first three (n, l and ml) describe an atomic
  orbital where an electron is found
• 3 numbers because three dimensions
• A fourth quantum number describes the spin of an
  electron

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Quantum Numbers 1 and 2
SHELLS

• 1. PRINCIPAL QUANTUM NUMBER (n) one on
  which energy of an electron principally depends
  orbitals with same n are in same shell also
  determines size of an orbital values 1, 2, 3,

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Quantum Numbers 1 and 2
SUB-SHELLS
2. ANGULAR MOMENTUM QUANTUM NUMBER (l) energy of
an atom also depends to a small extent on l
distinguishes orbitals of same n by giving them
different shapes any integer value from 0 to
n-1 orbitals of same n but different l are in
different sub-shells
s p d f g
0 1 2 3
4 Example 2p indicates shell 2 (energy),
sub-level p (shape)
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• Werner Heisenberg stated in 1927 that it is
  impossible to know with precision both the
  position and the momentum of an electron.
• The observer affects the observed.
• Only noticeable on the sub-atomic scale (e.g. an
  electron, not a baseball, nor dust).
• It is not possible to define a point in space
  where an electron will be found.
• However, we can obtain the probability of finding
  an electron at a certain point.

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An area where there is a greater than 90 chance
of finding an electron atomic orbital
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Quantum Number 3
ORBITALS

• 3. MAGNETIC QUANTUM NUMBER (ml)
  distinguishes orbitals in same shell and subshell
  (i.e. energy and shape) by giving them a
  different orientation in space values -l to l

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s orbitals
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p orbitals
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d orbitals
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SHELL
SUB-SHELL
ORBITAL
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Pauli Exclusion Principle

• No two electrons in any one atom can have the
  same four quantum numbers
• An orbital can hold at most two electrons, and
  then only if the two electrons in that orbital
  have opposite spins

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Quantum number 4

• 4. SPIN QUANTUM NUMBER (ms) each orbital can
  hold 2 electrons, and each has a different
  direction of spin, -½ and ½

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Hunds Rule

• Electrons fill orbitals singly and with parallel
  spins before pairing occurs.

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Electronic configurations

• Orbital notation

Spectroscopic notation
Outer electrons, control chemical properties
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Practise

• Write electronic configurations (obital-box and
  spectroscopic) for
• Helium
• Beryllium
• Oxygen
• Aluminium
• Calcium
• Bromine

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Aufbau principle

• The Aufbau principle is a building up principle,
  which helps you to write electronic
  configurations
• When electrons are placed in orbitals, the
  energy levels are filled up in order of
  increasing energy

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Filling orbitals
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s,
4f, 5d, 6p, 7s, 5f
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Energy depends on n and l Different orbitals in
the same sub-shell have the same energy Orbitals
with the same energy are called
degenerate, There is interaction among different
sub-shells in higher energy levels, which lowers
their energy and explains the order of filling.
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