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Chapter 12
Chemical Bonding
Ionic and Covalent Bonds – Who does what
Ionic Bonds – occur between a metal
cation and a non-metal anion, a metal
gives electron(s) to a non-metal.
Example: Na – Cl Fe – O
Covalent Bonds – occur between 2 non-
metals, electrons are shared by 2 nuclei.
Example: CH4 O = O
Figure 12.1: The formation of a bond between two atoms.
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•In an H2 molecule, the electrons reside primarily in the space between the 2 nuclei.
•Each H atom in the H2 molecule pulls the electrons equally based on the electronegativity of H.
•We call this type of bond “nonpolar covalent”
Figure 12.2: Probability representations of the electron
sharing in HF.
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•When we look at two different atoms, we see that the electrons are not shared equally, such as between H and F.
•The fluorine has a stronger attraction than the hydrogen for the shared
electrons.
•We call this bond “polar covalent”.
Dipole Moments
Any diatomic (2 atom) molecule that has a polar bond
has a dipole moment. When a molecule has a center of
positive charge and a center of negative charge, we call
this a dipole moment
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Figure 12.5: Charge distribution
in the water molecule.
Figure 12.6:
Polar water molecules are strongly attracted to positive ions
by their negative ends.
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Figure 12.6: Polar water molecules are strongly attracted to
negative ions by their positive ends.
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Figure 12.4:
The three possible
types of bonds. a) nonpolar covelant
(equal sharing of electrons
b) polar covelant
(unequal sharing of
electrons
c) ionic
(giving away and taking of
electrons)
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8
Table 12.2
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When we look at the electron configuration of atoms becoming ions, we see that they are trying to achieve a noble gas configuration.
General Rules Ionic bonds:
◦ Metals will become positive ions, they will lose electrons
to achieve the electron configuration of the previous
noble gas before them.
◦ Non-metals will become negative ions, they will gain
electrons to achieve the electron configuration of the next
noble gas.
Covalent bonds:
◦ Atoms in covalent bonds are sharing electrons, but all the
same, they want to have noble gas configurations in their
outer valence shells so they share electrons to achieve
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Practice
Determine the electron configuration for the
following elements, and then determine who will
gain and who will lose electrons in an ionic bond.
S and Mg
Al and O
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Ion Size vs Atom Size When atoms become ions,
the size of each changes.
A cation(+) is always smaller than the parent atom.
An anion (-) is always larger than the parent atom.
To explain this, think about who has lost electrons and who has gained electrons.
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Figure 12.9:
Relative sizes of some ions and their
parent atoms.
Lewis Structures
We must look at how many valence electrons an atom has to determine the Lewis structure of an atom.
Example: Fluorine
How many valence electrons?
Draw 1 dot for each valence electron around the symbol for the element. We draw the electrons in pairs on each of the 4 sides of the element symbol.
F
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Practice
Draw the Lewis structures for the following
elements:
He Al
H Mg
O K
C
N
S
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Lewis Dot Diagrams and Bonding
Ionic Bonds – let’s look at KBr
How many valence electrons do K and Br have?
Who is going to gain and lose electrons in this bonding interaction?
Once these 2 atoms have bonded in an ionic bond, each has gained or lost electrons, we would write the Lewis dot diagram for these as following
K+ [::Br::]-
Now you try, write the Lewis dot diagram for NaCl
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Lewis Dot Diagrams and Bonding
Covalent bonding – atoms have to share electrons to make them “happy”
The most important requirement for the formation of a stable compound is that the atoms achieve noble gas electron configurations.
Let’s look at OCl2 How many valence e-?
That’s not enough to gain or steal, so what if they share 2 pairs of electrons.
Bond Types
Single bond involves two atoms sharing one electron
pair. Example: F-F
Double bond involves two atoms sharing two pairs of
electrons. Example: O=O
Triple bond involves two atoms sharing 3 pairs of
electrons. Example: N≡N
Remember in each case, each atom wants a noble gas
configuration with 8 outer valence electrons.
Not all the electrons involved maybe bonded, you may
find you have a lone pair of electrons, a pair of electrons
that are not bonded.
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Practice
Use Lewis dot diagrams for the following atoms
to determine the bonds:
CH4
HF
NH3
N2
O2
SO2
HINT: Some may have multiple bonds.
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