Ch 11: Introduction to Chemical Bonding
Few atoms exist as independent particles. Most exist as combinations held together by
chemical bonds. The way in which they bond affects the properties of the substance.
Atoms form bonds by obtaining stable electron configurations (i.e., a full valance shell of
electrons), either by sharing electrons or by transferring them.
11.1 Types of Bonds
1. A chemical bond is a force that holds two or more atoms together and makes them
function as a unit.
a. Some bonds are easily broken others are very hard to break.
b. Bond Energy – energy required to break a bond.
2. Ionic Bonds
a. Result from the electrical attraction between a metal ion (cation) and a nonmetal ion
(anion).
b. An ionic compound is formed when an atom that loses electrons easily (a metal) reacts
with an atom that has an affinity for electrons (a nonmetal). One or more electrons are
transferred from the metal to the nonmetal.
c. There is generally a large electronegativity difference between the atoms (see section
11.2).
3. Covalent bonds
a. Result from the sharing of electron pairs between two nonmetal atoms.
b. One or both of the atoms can contribute the electrons to be shared.
c. Atoms share electrons because they have similar electronegativities (see figure 11-1
and section 11.2).
d. Shared electron pairs are considered to be localized between two atoms because this is
where they spend most of their time.
1
e. There are two kinds of covalent bonds:
(1)
Nonpolar Covalent (usually simply called “covalent”)
(a) Non-polar bonding results when two identical non-metals equally
share electrons between them. This type of bond is also formed between carbon and
hydrogen. The atoms have the same or almost identical electronegativities. (See fig 11.1)
(b)
(2)
Examples include H2, Cl2, N2, etc.
Polar covalent bonds
(a) Result from an unequal sharing of electrons. The atom with the higher
electronegativity attracts the electrons more strongly than the atom with less
electronegativity, resulting in an unequal sharing of electrons.
(b) Shared pairs of electrons are shifted from the center between the two
participating atoms, making one end of the molecule slightly positive and the other end
slightly negative. The bond is polarized.
(c)
Ex. Hydrogen fluoride (see fig 11-2)
(d)
Delta (δ) indicates a partial or fractional charge.
(e) It is important to remember that polar molecules (unless ions) are overall
neutral. The partial charges caused by the unequal sharing of electrons are not the same as
ionic charges which are caused by the actual transfer of electrons between atoms.
4. Hydrogen Bond.
a. A weak attraction between molecules that occurs when a
hydrogen atom is bonded to a small, highly electronegative
atom, such as N, O, or F.
b. The shared electrons are pulled closer to the more
electronegative atom, producing a partial positive charge on the
hydrogen and a partial negative charge on the more
electronegative atom.
c. The hydrogen end of one molecule has an attraction to
the more electronegative side of an adjacent molecule.
2
5. Metallic bonds
a. Bonding between metal atoms in pure metals or alloys (combinations of two or more
different metals).
b. Metal atoms are difficult to separate but can slide past each other fairly easily.
Bonding in metals is strong but nondirectional, meaning the bonds occur in any direction.
c. Because metal atoms are relatively large they can lose their outer electrons easily.
Large numbers of metal atoms share their valence electrons but in a manner different from
covalent bonding.
d. The metal atoms in a sample pool their valence electrons into an evenly distributed
“sea” of electrons that “flows” between and around the metal nuclei and core electrons. The
electrons are delocalized and move freely throughout the piece of metal.
e. The distinctive properties of metals (malleable, ductile, conductive, high melting
points, low ionization energy, etc.) are caused by the larger size of metal atoms, their
delocalized electrons, and the way in which metal atoms slide past each other but do not
easily separate.
f. In contrast, solid nonmetals are brittle/crumbly because their atoms are smaller with
electrons that are not easily lost, nor do they slide past each other.
3
11.2 Electronegativity
1. Electronegativity is the tendency of an atom to attract shared electrons to itself. (See fig.
11.3)
2. Group 1 and 2 elements lose electrons easily. They have low electronegativities.
3. Group 7 elements, however, can attract electrons easily. They have high
electronegativities.
4. Electronegativity generally increases going across the periodic table and decreases as you
go down the groups.
5. Using electronegativity values:
a. If the electronegativity difference, EN, between atoms is zero, the bond is covalent
(nonpolar).
b. If the EN is between 0.4 and 1.8, the bond is considered polar covalent. The more
electronegative atoms hold the shared electrons more strongly than the less electronegative.
This gives one end of the molecule a partial positive charge and one end a partial negative
charge.
c. Bonds between metals and nonmetals generally have a high EN. Electrons are
transferred, not shared, and the bond is ionic.
d. See Fig 11.4 and below:
4
d. Most bonds have some covalent and some ionic character. An ionic bond simply has
more ionic than covalent character, and a covalent bond has more covalent than ionic
character. There is no definite dividing line or cut-off between bond types. They exist along
a spectrum of ionic vs covalent character present in the bond. See Table 11.1 and below:
11.3 Bond Polarity and Dipole Moments
1. Molecules with polar bonds can have a dipole moment, defined as a property of a
molecule whereby the charge distribution can be represented by a center of positive and a
center of negative charge
2. The molecule behaves as though it has two centers of charge, one positive and one
negative.
3. See example for HF and H2O (Fig 11-5) in text. The dipole moment is generally depicted
as an arrow pointing away from the center of positive charge.
5
4. Water molecules are polar and have a dipole moment.
a. Since a water molecule has a bent shape its individual polar bonds result in a net
dipole moment. Notice that the center of positive charge lies between the hydrogen atoms.
The center of negative charge is by the oxygen. Since oxygen is more electronegative than
the hydrogen the bonding electrons are not shared equally.
b. Because water is polar it is attracted to other water molecules by forming hydrogen
bonds. This allows water to remain in the liquid state on the earth’s surface.
c. In contrast, carbon dioxide molecules, which
are three times more massive than water molecules,
remain in the gaseous state because they are
nonpolar and do not attract each other.
d. If water was not polar, it would only exist as
a gas at the normal range of temperatures on the
Earth’s surface. There would be no liquid water –
no oceans, lakes, etc.!! No life!
e. Water is able to dissolve most ionic compounds. They are attracted to positive ions
by their negative ends, and to negative ions by their positive ends. (See fig 11.6 and below)
f. Notice how anions are surrounded
by the positive ends of water molecules,
and cations are surrounded by the
negative ends of water molecules.
g. Ions are kept separated as long as
water is present. When the water is
removed the ions can reform the solid
crystal.
6
5. Molecules can have polar bonds yet be nonpolar.
a. CO2 is an example of a molecule that has polar bonds: O=C=O, but because the two
C=O dipole moments cancel each other out the molecule has no net polarity.
b. Carbon tetrachloride, CCl4, is another example of a molecule with polar bonds:
C―Cl, but the four resulting dipole moments cancel each other out resulting in a nonpolar
molecule with no net dipole moment.
7