5 Steps to a 5 AP Chemistry 2019

Bonding ❮ 151

In the hydrogen molecule, the electrons are shared equally. Each hydrogen nucleus has

one proton equally attracting the bonding pair of electrons. A bond like this is called a non-

polar covalent bond. In cases where the two atoms involved in the covalent bond are not

the same, the attraction is not equal, and the bonding electrons are pulled toward the atom

with the greater attraction. The bond becomes a polar covalent bond, with the atom that has

the greater attraction taking on a partial negative charge and the other atom a partial positive

charge. Consider for example, HF(g). The fluorine has a greater attraction for the bonding

pair of electrons (greater electronegativity) and so takes on a partial negative charge. Many

times, instead of using a single line to indicate the covalent bond, an arrow is used with the

arrowhead pointing toward the atom that has the greater attraction for the electron pair:

−

+→

δδ+−HF

The electronegativity (EN) is a measure of the attractive force that an atom exerts on a

bonding pair of electrons. Electronegativity values are tabulated. In general, electronegativi-

ties increase from left to right on the periodic table, except for the noble gases, and decrease

going from top to bottom. This means that fluorine has the highest electronegativity of any

element. If the difference in the electronegativities of the two elements involved in the bond

is great (>1.7), the bond is considered to be mostly ionic in nature. If the difference is slight

(<0.4), it is mostly nonpolar covalent. Anything in between is polar covalent.

Many times the Lewis structure will be used to indicate the bonding pattern in a cova-

lent compound. In Lewis formulas, the valence electrons that are not involved in bonding

are shown as dots surrounding the element symbols, while a bonding pair of electrons is

represented as a dash. There are several ways of deriving the Lewis structure, but here is one

that works well for those compounds that obey the octet rule.

Draw the Lewis structural formula for CH 4 O.
First, write a general framework for the molecule. In this case, the carbon must be
bonded to the oxygen, because hydrogen can form only one bond. Hydrogen is never central.
Remember: Carbon forms four bonds.

H

HCOH

H

To determine where all the electrons are to be placed, apply the N - A = S rule where:

N = sum of valence electrons needed for each atom. The two allowed values are two for

hydrogen and eight for all other elements.

A = sum of all available valence electrons

S = # of electrons shared and S/2 = # bonds

For CH 4 O, we would have:

1 C 4 H 1 O

N 8 + 4 (2) = 8 + 8 = 24

A 4 + 4(1) = 4 + 6 = 14

S = N – A = 24 – 14 = 10 bonds = S/2 = 10/2 = 5

Place the electron pairs, as dashes, between the adjacent atoms in the framework

and then distribute the remaining available electrons so that each atom has its full octet,