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Ionic Bonding
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Electron Transfer
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Ionic eh? Suppose ions are involved?
You guessed it! Ionic bonding is the result of a transfer of electrons
between two (or more) atoms.
The result of the electron transfer are ions (charged atoms).
Recall that a cation is an ions with a positive charge; while an anion is an
atom with a negative charge.
When is bonding ionic?
The general rule of thumb is that a metal and a non-metal make an ionic
compound.
This is true more often than not, but to be sure we should use the
electronegativity difference (EN).
To calculate EN we take the greater electronegativity value (atom 1) and
subtract the lesser (atom 2).
EN
Bonding Character
> 1.7
0.4 – 1.7
< 0.4
Ionic
Polar Covalent
Non-Polar Covalent
0
Purely Covalent
Our focus for
today.
Testing EN
Will sodium chloride (NaCl) be ionic or covalent?
Metal and non-metal should be ionic.
EN = 3.16 (Cl) – 0.93 (Na) = 2.23
EN > 1.7, therefore NaCl is ionic!
Will beryllium (BeCl2) be ionic or covalent?
Metal and non-metal should be ionic.
EN = 3.16 (Cl) – 1.57 (Be) = 1.59
0.4 < EN < 1.7, therefore BeCl2 is polar covalent!!!
From now on, check the EN to be sure of the type of bonding!
So how does it work?
Ionic bonding arises from transfer of electrons from metals to non-metals
in an attempt by each species to gain noble gas configuration.
HUH?!?!?
“Every atom wants 8 electrons!” (the octet rule)
Atoms from Groups 1 – 13 will give up their electrons to make cations
and achieve noble gas configuration.
Atoms from Groups 15 – 17 will gain electrons and achieve noble gas
configuration by making anions.
So how does it work?
The ionic bonding 3 step process (consider NaCl):
1) Na loses an electron (after I.E. provided) to make Na+ + 1e2) Cl picks up the electron (and releases E.A.) from sodium to make Cl3) Na+ and Cl- come together to make NaCl (and release lattice energy_
Stable octets!!!
(noble gas configurations)
An Animated Look
The ionic bonding 3 step process
1)2Na + IE  2Na+ + 2e2)Cl2 + 2e-  2Cl- + EA
3)2Na + Cl2  2Na+Cl- (2NaCl) + LE
An Electronic Look
Consider the formation of NaCl
Na + IE  Na+ + 1e- 3s1  2p6
Cl + 1e-  Cl- + EA (3p5  3p6)
The Ionic Lattice
Ionic compounds typically:
 have very high melting points
 conduct electricity in the molten state and when dissolved in water
 are brittle (fracture) solids
 have low volatility
The strength of the bonding
interaction is the result of the sum of
all bonding interactions, every Na+
affects every Cl-.
While the individual bonds between
2 atoms are quite weak the overall
effect of the lattice makes the
interactions very strong.
Lattice Energy
When an ionic lattice forms it releases a large amount of energy which we
call the lattice energy. (Need to put this in to break it apart)
The smaller the ions, the
tighter the fit an higher the
lattice energy.
LiF DHlatt = -1036 kJ/mol
NaCl DHlatt = -787.3 kJ/mol
Lattice energy also becomes MUCH
stronger as the charge on the ions
becomes larger.
Na+ O2- (Na2O) = -2481 kJ/mol
Al3+ O2- (Al2O3) = - 15916 kJ/mol
A look a physical properties
Like dissolves like.
A polar substance, such as water will easily dissolve an ionic compound.
The electronegative oxygen atoms pickup Na+ while the electropositive
hydrogens pickup the Cl-.
As the lattice energy
terms gets larger in
magnitude the solids
are much harder to
break up and dissolve
in water.
Ions and Atomic Size
Atomic radius largely depends on the effective nuclear charge.
Lets consider what happens to the size of an atom when it loses electrons
(effective nuclear charge goes up  atom gets smaller)
When an atom gains electrons the effective nuclear charge decreases, the
electrons get farther away  it gets bigger
Li+
Be2+
B3+ C4+
N3-
O2-
F-
A cation is smaller than the original atom while an anion
is larger.
Ionic Compounds
1.
Check the ionization energy to see if each molecule will have an ionic bond.
2.
Draw the Lewis structures for each element
3.
Give the 3 step solutions for the formation of the following compounds:
a)
Aluminum nitride
b)
Rubidium sulfide
c)
Magnesium oxide
d)
Potassium bromide
e)
Sodium oxide
f)
Lithium chloride
g)
Potassium sulfide
h)
Francium fluoride
i)
Strontium iodide
j)
Cesium phosphide