Download Bonding. A. Ionic bonds form when anions and cations arise

Bonding.
A.
Ionic bonds form when anions and cations arise from a transfer of electrons.
1. The transfer of electrons can occur if the two atoms have complementary octet rules.
Consider the case of common table salt, sodium chloride.
a) The octet rule for sodium is to lose one electron.
b) The octet rule for chlorine is to gain one electron.
2.
B.
C.
D.
These two octet rules are complementary. When the electron clouds of these two
atoms overlap, each can neatly fulfill its octet rule.
Covalent bonds form when it is not possible for electrons to be transferred and so must be
shared between atoms.
1. Generally this is the case when two nonmetals bond. The tendency of nonmetals is to
gain electrons according to the octet rule, a very easy thing to accomplish when they
bond with metals. But, when bonding to one another, this can't be accomplished.
2.
In this case we say that the octet rules are conflicting. Consider the case of water.
a) The octet rule for hydrogen is to gain one electron.
b) The octet rule for oxygen is to gain two electrons.
3.
These two octet rules are conflicting. When the electron clouds of these two atoms
overlap, each will claim the electrons in the overlapping region. The ensuing tug-of-war
is the covalent bond.
Using formulas to represent compounds.
1. In a chemical formula, the most electropositive element is listed first.
a) In an ionic compound, this is the cation that actually has a positive charge.
b) In a covalent compound, this is the more electropositive element that will have no
more than a partial positive charge.
2.
The most negative element is listed last.
a) In an ionic compound, this is the anion that actually has a negative charge.
b) In a covalent compound, this is the more electronegative element that will have no
more than a partial negative charge.
3.
Subscripts are used to indicate the number of atoms (in the covalent case) or ions (in
the ionic case) present.
Predicting the bonding type.
1. Use the octet rules.
a) Complementary - ionic bonding.
b) Conflicting - covalent bonding.
2.
Use the “stair-step” line.
a) Opposite side (metal - nonmetal) - ionic bonding.
b) Same side (nonmetal - nonmetal) - covalent bonding.
3.
Use electronegativity difference.
a) Tables list the electronegativities of the representative elements.
b) Electronegativity is a measure of the willingness of an atom to accept an additional
electron.
(1) Fluorine is the most electronegative element.
(2) Cesium and francium are the least electronegative.
c)
The difference in electronegativities ranges from a low of zero (100 % covalent
character) to a high of 3.3 (100 % ionic character).
d)
An electronegativity difference of 0.0 - 0.4 indicates true covalent bonding.
e)
An electronegativity difference of 0.5 - 2.0 indicates polar covalent bonding.
f)
An electronegativity difference of 2.1 - 3.3 indicates true ionic bonding.
E.
Because of the electronegativity differences between atoms, it is not always possible for the
octet rules to be followed rigorously. Oxidation numbers offer a summary of the octet rule
each atom followed in the bonding process. Follow these rules to determine the oxidation
number of any atom:
1. The oxidation number of any element in its free (or uncombined) state is 0.
2. The oxidation number of any ion is the charge of the ion.
3. The oxidation number of any family IA metal is +1.
4. The oxidation number of any family IIA metal is +2.
5. The oxidation number of hydrogen is +1 (except as the hydride ion when it is -1).
6. The oxidation number of oxygen is -2 (except as the peroxide ion when it is -1).
7. The sum total of oxidation numbers in a neutral compound is zero.
8. The sum total of oxidation numbers in a radical (or polyatomic ion) is equal to the
charge of the radical.
F.
Oxidation number examples.
1. H2SO4
2. Na2S
-2
3. Cr2O 7
+1 +6 -2
+1
-2
+6
-2
+2 +6 -8
+2
-2
+12 -14