Chemical Bonding and Molecular Structure
What Are Chemical Bonds, and Why Do They Form?
A chemical bond
is the result of an attraction
between atoms or ions. The types of bonds that a molecule contains
will determine its physical properties, such as melting point, hardness,
electrical and thermal conductivity, and solubility. How do chemical
bonds occur? As we mentioned before, only the outermost, or valence
electrons of an atom are involved in chemical bonds. Let’s begin
our discussion by looking at the simplest element, hydrogen. When
two hydrogen atoms approach each other, electron-electron repulsion
and proton-proton repulsion both act to try to keep the atoms apart.
However, proton-electron attraction can counterbalance this, pulling
the two hydrogen atoms together so that a bond
Look at the energy diagram below for the formation of an H–H bond.
As you’ll see throughout our discussion, atoms will often
gain, lose, or share electrons in order to possess the same number
of electrons as the noble gas that’s nearest them on the periodic
table. All of the noble gases have eight valence electrons (s2p6)
and are very chemically stable, so this phenomenon is known as the octet
rule. There are, however, certain exceptions to the octet
rule. One group of exceptions is atoms with fewer than eight electrons—hydrogen
(H) has just one electron. In BeH2, there
are only four valence electrons around Be: Beryllium contributes
two electrons and each hydrogen contributes one. The second exception
to the octet rule is seen in elements in periods 4 and higher. Atoms
of these elements can be surrounded by more than four valence pairs
in certain compounds.
Types of Chemical Bonds
You’ll need to be familiar with three types of chemical
bonds for the SAT II Chemistry exam: ionic bonds, covalent bonds,
and metallic bonds.
Ionic bonds are the result of an electrostatic
attraction between ions that have opposite charges; in other words,
cations and anions. Ionic bonds usually form between metals and nonmetals;
elements that participate in ionic bonds are often from opposite
ends of the periodic table and have an electronegativity difference
greater than 1.67. Ionic bonds are very strong, so compounds that
contain these types of bonds have high melting points and exist
in a solid state under standard conditions. Finally, remember that
in an ionic bond, an electron is actually transferred from
the less electronegative atom to the more electronegative element.
One example of a molecule that contains an ionic bond is table salt,
Covalent bonds form when electrons are shared between
atoms rather than transferred from one atom to another. However,
this sharing rarely occurs equally because of course no two atoms
have the same electronegativity value. (The obvious exception is
in a bond between two atoms of the same element.) We say that covalent
bonds are nonpolar if the electronegativity difference
between the two atoms involved falls between 0 and 0.4. We say they
are polar if the electronegativity difference falls
between 0.4 and 1.67. In both nonpolar and polar covalent bonds,
the element with the higher electronegativity attracts the electron
pair more strongly. The two bonds in a molecule of carbon dioxide,
CO2, are covalent bonds.
Covalent bonds can be single, double, or triple. If only
one pair of electrons is shared, a single bond is formed.
This single bond is a sigma bond (s),
in which the electron density is concentrated along the line that
represents the bond joining the two atoms.
However, double and triple bonds occur frequently (especially
among carbon, nitrogen, oxygen, phosphorus, and sulfur atoms) and
come about when atoms can achieve a complete octet by sharing more
than one pair of electrons between them. If two electron pairs are
shared between the two atoms, a double bond forms,
where one of the bonds is a sigma bond, and the other is a pi
bond (p). A pi bond is a bond in which the electron
density is concentrated above and below the line that represents
the bond joining the two atoms. If three electron pairs are shared
between the two nuclei, a triple bond forms. In a triple
bond, the first bond to form is a single, sigma bond and the next
two to form are both pi.
Multiple bonds increase electron density between two nuclei:
they decrease nuclear repulsion while enhancing the nucleus-to-electron
density attractions. The nuclei move closer together, which means
that double bonds are shorter than single bonds and triple bonds
are shortest of all.
Metallic bonds exist only in metals, such
as aluminum, gold, copper, and iron. In metals, each atom is bonded
to several other metal atoms, and their electrons are free to move throughout
the metal structure. This special situation is responsible for the
unique properties of metals, such as their high conductivity.
Drawing Lewis Structures
Here are some rules to follow when drawing Lewis structures—you
should follow these simple steps for every Lewis structure you draw,
and soon enough you’ll find that you’ve memorized them. While you
will not specifically be asked to draw Lewis structures on the test,
you will be asked to predict molecular shapes, and in order to do
this you need to be able to draw the Lewis structure—so memorize
these rules! To predict arrangement of atoms within the molecule
Find the total number of valence electrons by adding
up group numbers of the elements. For anions, add the appropriate
number of electrons, and for cations, subtract the appropriate number
of electrons. Divide by 2 to get the number of electron pairs.
which is the central atom—in situations where the central atom has
a group of other atoms bonded to it, the central atom is usually
written first. For example, in CCl4, the
carbon atom is the central atom. You should also note that the central
atom is usually less electronegative than the ones that surround
it, so you can use this fact to determine which is the central atom
in cases that seem more ambiguous.
one pair of electrons between each pair of bonded atoms and subtract the
number of electrons used for each bond (2) from your total.
lone pairs about each terminal atom (except H, which can only have two
electrons) to satisfy the octet rule. Leftover pairs should be assigned
to the central atom. If the central atom is from the third or higher
period, it can accommodate more than four electron pairs since it
has d orbitals in which to place them.
the central atom is not yet surrounded by four electron pairs, convert
one or more terminal atom lone pairs to double bonds. Remember that
not all elements form double bonds: only C, N, O, P, and S!
Which one of the following molecules contains a triple
bond: PF3, NF3, C2H2,
H2CO, or HOF?
The answer is C2
which is also known as ethyne. When drawing this structure, remember
the rules. Find the total number of valence electrons in the molecule
by adding the group numbers of its constituent atoms. So for C2
this would mean C = 4
2 (since there
are two carbons) = 8. Add to this the group number of H, which is
1, times 2 because there are two hydrogens = a total of 10 valence
electrons. Next, the carbons are clearly acting as the central atoms
since hydrogen can only have two electrons and thus can’t form more
than one bond. So your molecule looks like this: H—C—C—H. So far
you’ve used up six electrons in three bonds. Hydrogen can’t support
any more electrons, though: both H’s have their maximum number!
So your first thought might be to add the remaining electrons to
the central carbons—but there is no way of spreading out the remaining
four electrons to satisfy the octets of both carbon atoms except
to draw a triple bond between the two carbons.
For practice, try drawing the structures of the other
four compounds listed.
How many sigma (s) bonds and how many pi (p) bonds does
the molecule ethene, C2H4, contain?
First draw the Lewis structure for this compound, and
you’ll see that it contains one double bond (between the two carbons)
and four single bonds. Each single bond is a sigma bond, and the
double bond is made up of one sigma bond and one pi bond, so there
are five sigma bonds and one pi bond.
Exceptions to Regular Lewis Structures—Resonance
Sometimes you’ll come across a structure that can’t be
determined by following the Lewis dot structure rules. For example,
ozone (O3) contains two bonds of equal bond
length, which seems to indicate that there are an equal number of
bonding pairs on each side of the central O atom. But try drawing
the Lewis structure for ozone, and this is what you get:
We have drawn the molecule with one double bond and one
single bond, but since we know that the bond lengths in the molecule
are equal, ozone can’t have one double and one single bond—the double
bond would be much shorter than the single one. Think about it again,
though—we could also draw the structure as below, with the double
bond on the other side:
Together, our two drawings of ozone are resonance structures
for the molecule. Resonance structures
are two or
more Lewis structures that describe a molecule: their composite
represents a true structure for the molecule. We use the double-directional
arrows to indicate resonance and also bracket the structures or
simply draw a single, composite picture.
Let’s look at another example of resonance, in the carbonate
Notice that resonance structures differ only in electron
pair positions, not atom positions!
Draw the Lewis structures for the following molecules:
HF, N2, NH3, CH4,
CF4, and NO+.