Ionic bonds and main
group chemistry
Towards the noble gas configuration
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Noble gases are unreactive – they have filled
shells
Shells of reactive elements are unfilled
Achieve noble gas configuration by gaining or
losing electrons
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Metals lose electrons – form positive ions
Nonmetals gain electrons – form negative ions
Lewis dot model
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The nucleus and all of the core electrons are represented by the
element symbol
The valence electrons are represented by dots – one for each
Number of dots in Lewis model is equal to group number (in
1 – 8 system)
The Octet Rule
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All elements strive to
become a noble gas, at
least as far as the
electrons are concerned.
Filling the outer shell –
8 electrons
Achieve this by adding
electrons
Or taking them away
Predicting ion charges
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s and p block elements are easy:
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charge = group number for cations
charge = -(8 – group number) for anions
Less predictable for transition
metals
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Occurrence of variable ionic charge
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Cr2+, Cr3+, Cr4+, Cr6+ etc.
4s electrons are lost first and then the 3d
Desirable configurations coincide with empty,
half-filled or filled 3d orbitals
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Fe2+ ([Ar]3d6) is less stable than Fe3+ ([Ar]3d5)
Ionic size and charge
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Loss of electrons
increases the effective
nuclear charge – ion
shrinks
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Gain of electrons
decreases the effective
nuclear charge – ion
expands
Ionization energy
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Energy required to remove an electron from a
neutral gaseous atom
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Always positive
Follows periodic trend
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Increases across period
Decreases down group
Removal of electrons from filled or half-filled shells is not as favourable
[He]2s22p3
[He]2s22p4
[He]2s2
[He]2s22p1
Higher ionization energies
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Depend on group number
Much harder to remove electrons from a filled shell
Stepwise trend below illustrates this Partially filled –
Completely
filled – core
electrons
valence
electrons
Electron affinity
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Energy released on adding an electron to a neutral gaseous
atom
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Values are either
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negative – energy released, meaning negative ion formation is favourable
Or zero – meaning can’t be measured and negative ions are not formed
Addition of electrons to filled or half-filled shells is not favoured (e.g. He, N)
It is easier to add an electron to Na (3s1) than to Mg (3s2)
Ionic bonding
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Reaction between elements that form positive and
negative ions
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Metals (positive ions) and nonmetals (negative ions)
Neutral Na + Cl → ionic Na+Cl
[Ne]3s1 + [Ne]3s23p5 = [Ne]+ + [Ar]-
Stability of the ionic lattice
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Simply forming ions does not give an energy payout:
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Ei(Na) = 496 kJ/mol
Ea(Cl) = -349 kJ/mol
Net energy investment required
Formation of a crystal lattice releases energy
The lattice energy is the energy released on bringing
ions from the gas phase into the solid lattice
Depends on coulombic attraction between ions
-U = κz1z2/d (κ = 8.99x109 JmC-2
Born-Haber cycle for calculating
energy
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The lattice energy can be obtained using other
experimentally determined quantities and the energy
cycle
Lattice energies follow simple trends
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As ionic charge increases, U increases (U  z1z2)
As ion size decreases, U increases (U  1/d)
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U(LiF) > U(LiCl) > U(LiBr)
U(NaI) < U(MgI2) < U(AlI3)
The Octet Rule
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Main-group elements undergo reactions which
leave them with eight valence electrons
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Group 1 (ns1) M+
Group 2 (ns2) M2+
Group 6 (ns2np4) X2Group 7 (ns2np5) X-
Works very well for second row (Li – F)
Many violations in heavier p-block elements
(Pb2+, Tl+, Sb3+)