20 Chemistry
Chapter 9 Notes
Molecular Compounds and Covalent Bonds
We have already learned how ionic compounds form. A metal looses 1 or more electrons
to obtain a stable electron configuration and a non-metal gains one or more electrons to
obtain a stable electron configuration. The positively charged metal ion is
electrostatically attracted to the negative non-metal ion and the strong ionic bond is
formed.
Molecular compounds do not involve metals. As such there are no metals to give up
their electrons and become positively charged. About the only thing that molecular
compounds have in common is the fact that the atoms that make them up acquire a stable
electron configuration. But without electrons being lost and gained how is this
possible?
With molecular compounds, the atoms that make up the molecule share electrons. More
specifically, one electron from each of two atoms (that will bond), pair up. This shared
pair of electrons is called a covalent bond.
H + H + O
H O
H
The two electrons between the hydrogen and oxygen atoms are the shared pairs. Recall
the Lewis dot representation shows only the outer electrons. If you count the electrons
around each hydrogen you get 2 (the stable configuration for hydrogen). If you count the
number of electrons around the Oxygen symbol, you get eight (again a stable
configuration).
Other examples
H
H N H
H
Try p. 244 # 1 - 5
and
H C H
H
Another way of representing these covalent bonds is with a line as indicated below.
H
H N H
and
H C H
H
H
Multiple Covalent Bonds
In many molecules, more than one pair of shared electrons are required to attain a stable
electron configuration. Examples
A Double Covalent Bond:
O
+
O
O
O
+
N
N
N
A Triple Covalent Bond:
N
Bond Dissociation Energy
Because molecules are stable, energy is released when bonds form. Conversely, energy
is required to break bonds. The amount of energy required to break a specific covalent
bond is called bond dissociation energy. Consider these three molecules:
F
O
F
O
N N
BDE equals 159 kJ/mol
BDE equals 498 kJ/mol
Decreasing
Bond
Length
BDE equals 945 kJ/mol
Can you deduce the relationship between the number of covalent bonds between two
atoms and the BDE?
Consider the following organic molecule:
H
H
C
H
C
H
The sum of the bond dissociation energy values for all the bonds in a compound is used
to determine the chemical potential energy available in a molecule of that compound.
See p. 781 for structural formulas for sugars and p. 699 for other fuels
In chemical reactions involving molecules, the bonds of the reactant molecules require
energy to be broken and energy is released when the product molecule bonds form. If the
energy required by the reactants is greater than the energy released by the products, the
reaction is endothermic. If however the energy required by the reactants is less than the
energy released by the products, the reaction is exothermic.
Naming Molecular Compounds
A binary molecular compound is one that contains two different nonmetal elements. For
binary molecular compounds (when you know the chemical formula):



The first word in the name is the same as the elements name
The second word has an “ide” ending
Use prefixes (see table 9-1 p. 248) to indicate the number of atoms of each
element
* - note that we omit the mono prefix on the first element
* - also note that we drop the final letter in the prefix if the element starts with a vowel
example - P2O5 is called diphosphorus pentoxide
Try p. 249 # 13 – 17
Naming Acids
Many acids are molecular compounds that are dissolved in water to liberate hydrogen
ions. We learn more about acids later. How we name acids depends on whether the acid
contains oxygen (oxyacid) or not.
1) Acids that do not contain oxygen – usually contain hydrogen and some other element.
To name these acids, we start with the prefix hydro followed by the root of the second
element with an ic ending. examples
HCl – hydrochloric acid
HBr – hydrobromic acid
HCN – hydrocyanic acid
2) Oxyacids – are those acids that have an oxyanion (polyatomic ion containing oxygen)
like NO3- , or SO42- . Here’s a couple of things to keep in mind when naming oxyacids:



i.e.
The acid name derives from the the anion (there is no hydrogen in the name)
the anion ends in ic if the oxyanion ends in ate
the anion ends in ous if the oxyanion ends in ite
an acid derived from the polyatomic SO42- ion forms sulfuric acid – H2SO4
an acid derived from the polyatomic SO32- ion forms sulfurous acid – H2SO3
Try p. 250 # 18 – 22
Try p. 247 # 6 – 12 a - d
Drawing Lewis Structures







guess the location of the individual atoms in this two dimensional representation
add up the total number of available electrons (valence electrons) of all the atoms
in the atom or polyatomic ion. For negative ions, the number of available
electrons is increased by the amount of charge on the ion. For positive ions, the
number of available electrons is decreased by the charge on the ion. Why?
divide this number by two to get the number of electron pairs
place one covalent bond (one pair) between each two atoms in the structure
place electrons (in pairs) around the outer atoms to satisfy the octet rule keeping
track of the number of pairs you have used up
any left over pairs must satisfy the octet rule for inner atoms until you have run
out of pairs
finally, ensure that there are 8 electrons around each atom (two for hydrogen) and
if not, you may have to shift a pair to make a double (or triple) covalent bond.
Example
NH3 and CH4
H
H N H
H
and
H C H
H
Try p. 255 # 30 – 34 (also try any or all of the polyatomic ions from the reference sheet)
Section 9-4
We are only responsible to know that different types of molecules have different shapes
and the shape depends on the number and type of covalent bonds and the number of
unshared pairs of electrons. The model used to determine molecular shape is called the
Valence Shell Electron Pair Repulsion model or VSEPR.
See table 9-3 p. 260
Section 9-5 - Polarity and Covalent Bonds
When chemical reactions take place, chemical compounds are “broken down” and new
bonds are formed dictating the products of the reaction. But what determines which kind
of bond (and thus which substance) will form in the reaction? The answer is the
electronegativities of the atoms present. Review electronegativities of elements in
chapter 6 and refer to figure 9-15 p. 263 (note that noble gases are excluded…why?)
Note that Fluorine has the highest electronegativity (3.98) and Francium has the lowest
(0.70).
The type of bond formed between two atoms in a chemical compound can be predicted
by the difference between their electronegativities.
If the two atoms bonding are identical, than the difference between their
electronegativities is zero and the bond is a non-polar covalent or pure covalent bond.
The pair of electrons that is shared is shared equally because each atom has an equal
attraction for the electrons in the bond.
For covalent bonds formed between two non-identical atoms, the pair of electrons will
not be shared equally because the two atoms have different electronegativities. This
unequal sharing results in a covalent bond that is polar thus the bond is called a polar
covalent bond. How polar the bond is depends on how different the electronegativities
of the atoms are.
As the difference in electronegativities becomes greater, the bond becomes less covalent
and more ionic. Any bond that has an electronegativity difference of 1.70 or less is
considered covalent and any bond that has an electronegativity difference of 1.70 or
more is considered ionic. But these are just labels. Most bonds are really neither one or
the other but rather have characteristics of both.
Back to Polar Covalent Bonds
The unequal sharing of electrons in a polar covalent bond results in the electrons
spending a greater percentage of their time closer to the atom with the higher
electronegativity. As a result the bond itself is polarized meaning one end has a different
charge than the other. The end of the bond where there is a greater likelihood of finding
electrons is negatively charged while the other end is positively charged. This charge is
indicated with the symbol δ- or δ+ as indicated in the example below;
δ+
H
δO
This polar covalent bond is often referred to as a dipole.
Molecules can be described as polar or non-polar as well. In order for a molecule to be
polar it must contain polar covalent bonds but not all molecules that have polar covalent
bonds are polar. As a general rule, symmetric molecules (refer to p. 260) tend to be nonpolar and asymmetric molecules tend to be polar. See the example that compares the
polarity of H2O and CCl4 molecules on p. 265.
Inter-Molecular Forces
The forces that attract the molecules of a covalent compound are called van der Waals
forces. There are three types of van der Waals forces
 non-polar compounds have dispersion forces or induced dipoles – they are weak
 polar compounds have stronger dipole-dipole forces
 hydrogen bonds which form between hydrogen and either fluorine, oxygen or
nitrogen – they are strongest of the three inter-molecular forces
In general, the inter-molecular forces between the molecules of a molecular compound
are weaker than the forces that hold an ionic substance together. Why?
Bonds and Physical Properties
Melting points - of molecular substances are lower i.e. sugar melts but salt doesn’t
Boiling points - of molecular substances are lower i.e. many molecular compounds exist
as gases at room temperature
Hardness
- of molecular substances are lower i.e wax is a molecular solid
Solubility
- polar compounds (molecular and ionic) compounds dissolve in polar
substances and non-polar compounds dissolve in non-polar substances