Section 6.2: Drawing and Naming Compounds
Objectives:
1. Draw Lewis structures to show the arrangement of valence electrons among atoms in molecules and
polyatomic ions.
2. Explain the difference between single, double, and triple covalent bonds.
3. Draw resonance structures for simple molecules and polyatomic inons, and recognize when they are
required.
4. Name binary inorganic covalent compounds by using prefixes, roots, and suffixes.
Key Terms:
 valence electron
 Lewis structure
 unshared pair
 single bond



double bond
triple bond
resonance structure
Lewis Electron-Dot Structures
Chemical bonding involves the transfer or sharing of electrons in the outer energy shell of atoms. These
electrons are called valence electrons. The movement of valence electrons that results in bonding occurs so
that atoms can achieve a stable valence shell similar to outer shell structure of a noble gas.
Ionic bonding involves complete transfer of valence electrons. This creates charged particles called ions.
Positive ions, called cations, are generally made of metals. Negative ions, called anions, are generally made
from non-metals. Cations and anions come together to form compounds called salts.
In contrast to ionic bonding, covalent bonding involves the sharing of valence electrons to achieve stability.
A non-polar covalent bond occurs where the bonding electrons are equally attracted to both bonded atoms. A
polar covalent bond takes place when the shared pair of electrons is held more closely by the atom that is
more electronegative.
Lewis electron-dot structures provide a system used to represent and keep track of the valence electrons in a
stable atom. Recall that a valence electron is an electron that is found in the outermost energy shell of an
atom and that these electrons determine the chemical properties of the atom. In a Lewis dot structure, the
structural formula of these electrons is represented by dots. A pair of dots or dashes between two atomic
symbols represents covalent bonds. Lewis dot structures help you keep track of the valence electrons.
Lewis structures are used to model covalently bonded molecules. A Lewis structure shows only the valence
electrons in the atom or molecule. A molecule consists of two or more atoms that are covalently bonded. The
nuclei and inner electrons of the atom are represented by the symbol of the element. Each side of the
chemical symbol can hold up to 2 electrons. Since there are 4 sides (left, right, top, bottom), the total number
of valence electrons is up to eight electrons. For example, chlorine has 7 valence electrons ([Ne]3s23p5) and
the following Lewis structure:
The following table shows the electron configuration and Lewis structures of the elements in the second row
of the periodic table as they appear before bonding. Note that as you go from element to element, you add a
dot to each side of the element’s symbol. Do not begin to pair dots until all four sides of the element’s
symbol has a dot (the electrons try to stay as far apart as possible until they have to pair).
Electron
Element
Configuration
Li
Be
B
C
N
O
F
Ne
1s22s1
1s22s2
1s22s22p1
1s22s22p2
1s22s22p3
1s22s22p4
1s22s22p5
1s22s22p6
Number of
Valence
Electrons
1
2
3
4
5
6
7
8
Lewis structure of atoms in second row before bonding:
When non-metals come together to form a covalent bond, they share electrons so that each element will have
8 electrons. This tendency of bonded atoms to have octets of valence electrons is called the octet rule. When
two chlorine atoms form a covalent bond, each atom contributes one electron to a shared pair.
In a Lewis structure, each chlorine atom in Cl2 has three pairs of electrons that are not part of the bond.
These pairs are called unshared pairs or lone pairs. The pair of dots that represents the shared pair of
electrons can also be shown by a long dash. Remember, each long dash represents two shared electrons.
The guidelines for drawing Lewis structures of molecules that contain multiple atoms are as follows:
1. Count the total number of valence electrons in the compound. If you are finding the structure of
anion, remember to add one electron for each negative charge and subtract one electron for each
positive charge.
2. Predict the arrangement of the atoms in the molecule, drawing a line to represent a single bond
between each pair of bonded atoms. Note that the first atom listed in the formula is usually the central
atom. Also remember that H can only bond to one other atom.
3. Calculate the number of valence electrons left over after forming the single bonds. Remember that
each single bond you drew counts for two electrons.
4. Place electrons around the outer atoms first until each is surrounded by eight electrons (the octet
rule). Remember that H does not follow the octet rule. It will only have two electrons.
5. Place any leftover electrons around the central atom. It is ok to have more than eight electrons
(violates the octet rule) around the central atom as long as it is an element from the third row or lower
on the periodic table. If there are not enough electrons to put eight around the central atom, go back
and try changing one or more of the single bonds to double or triple bonds. Double or triple bonds
are most commonly formed between C, N, O, and S.
Multiple Bonds
Atoms can share more than one pair of electrons in a covalent bond. Atoms can share 2 pairs (a double bond)
or 3 pairs (a triple bond). An example is the bond between the two oxygen atoms in oxygen gas, O2. Each
oxygen atom has 6 valence electrons. To make an octet, each oxygen atom needs two more electrons. In
order for each atom in the O2 molecule to have 8 electrons, they must share two electrons and form a double
bond.
An example of a triple bond is that formed between the two nitrogen atoms in nitrogen gas (N2). Because the
two N atoms share the electrons equally, the triple bond is a nonpolar covalent bond. Also, the triple bond in
nitrogen gas is strong making nitrogen gas fairly inert (or non-reacting).
Resonance Structures
Some molecules, such as ozone, O3, cannot be represented by a single Lewis structure. Ozone has two Lewis
structures, as shown below.
Each O atom follows the octet rule, but the two structures represent different arrangements of single and
double bonds. When a molecule has two or more possible Lewis structures, the two structures are called
resonance structures. The actual structure of O3 is an average between the two resonance structures. Another
molecule that has resonance is sulfur dioxide (SO2).
Naming Covalent Compounds
Covalent compounds are named using rules used to name ionic compounds. However, in covalent
compounds, a prefix is used to give the number of elements in the compound.
Prefix
monoditritetrapenta-
Number of atoms
1
2
3
4
5
Example
CO
SiO2
SO3
SCl4
SbCl5
Name
carbon monoxide
silicon dioxide
sulfur trioxide
sulfur tetrachloride
antimony pentachloride
Prefixes are added to the first element in the name only if the molecule contains more than one atom of that
element. If the molecule contains only one atom of the first element given in the formula, the prefix mono- is
left off. The vowels a and o are dropped from the prefix that is added to a word beginning with a vowel. For
example, CO is carbon monoxide, not carbon monoxide. Similarly, N2O4 is named dinitrogen tetroxide, not
dinitrogen tetraoxide.