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 is formed. 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, NaCl.
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
1. 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.
2. Determine
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.
3. Place
one pair of electrons between each pair of bonded atoms and subtract the number
of electrons used for each bond (2) from your total.
4. Place
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.
5. If
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!
Example
Which one of the following
molecules contains a triple bond: PF3, NF3, C2H2, H2CO, or HOF?
Explanation
The answer is C2H2, 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 C2H2, 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.
Example
How many sigma (s) bonds
and how many pi (p) bonds does the molecule ethene, C2H4, contain?
Explanation
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 Structures
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 ion CO32-:
Notice that resonance
structures differ only in electron pair positions, not atom positions!
Example
Draw the Lewis structures
for the following molecules: HF, N2, NH3, CH4, CF4, and NO+.
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