Lesson 8.3 Covalent Bonding
Suggested Reading
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Zumdahl Chapter 8 Section 8.1, 8.2, 8.3
Essential Question
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What are the basic characteristics of covalent bonds?
Learning Objective
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List and define the three types of bonding.
Predict relative electronegativity for the positions of atoms in the periodic table.
Predict relative bond polarities from the electronegativities of atoms.
Predict whether a compound of two elements would be ionic or covalent from the
position of the elements in the periodic table of from their electronegativity
values,
Introduction
In the previous lesson we looked at ionic compounds, which are typically solids with
high melting points. Many substances, however, are molecular, which are gasses,
liquids, or low melting point solids consisting of molecules. A molecule is a group of
atoms, frequently nonmetals, strongly linked by chemical bonds. Covalent bonds
generally can't be understood on the same basis as ionic bonds, which are bonded by
the attraction of oppositely charged ions. Consider the H2 molecule; the two atoms are
held together tightly although no ions are present. In 1916 G.N. Lewis proposed that the
strong attractive force between two atoms in a molecule results from a covalent bond, a
chemical bond formed by the sharing of a pair of electrons between atoms. A few years
later the covalent bond in hydrogen was quantitatively explained using the newly
discovered theory of quantum mechanics. This provided evidence in support Lewis'
hypothesis. Lets look at the findings.
Describing Covalent Bonding
Consider the formation of a covalent bond between to H atoms. As the atoms approach
one another, their 1s orbitals begin to overlap. Each electron can then occupy the space
around both atoms such that the two electrons can be shared by the atoms. The
electrons are attracted simultaneously by the positive charges of the two hydrogen
nuclei, which causes a covalent bond to form. Although ions do not exist, the force that
hold the atoms together can still be regarded as arising from the attraction of oppositely
charged particles, nuclei and electrons.
The potential-energy curve for H2 shows that the formation of the bond is energetically
favorable. The graph plots the potential energy against the distance between nuclei. If
the nuclei are too close the atoms repel (top left). Too far, and they do not attract one
another (middle right). The bond occurs when the H atoms achieve a lowest energy
state observed on the potential energy diagram (bottom middle). The distance between
nuclei at this minimum energy is called the bond length of H2. It is the normal distance
between nuclei in the molecule. The bond length of H2 is 74 pm. The bond energy is
the average enthalpy change for the breaking of the bond in H2. Bond length is equal to
the difference between the minimum potential energy and PE = 0.
Recall that Lewis dot structures can be used to illustrate covalent bonds as follows.
Coordinate Covalent Bonds
Often covalent bonds are formed between atoms that both donate electrons as is the
case of hydrogen shown above. However, it is possible for both electron to come from
the same atom. A coordinate covalent bond (dative bond) is formed when both
electrons of the bond are donated by one atom.
This type of covalent bond is not essentially different from other covalent bonds.
However, being aware of and having an understanding of this type of bond can enhance
your understanding of how molecular compounds form.
Multiple Bonds
In the molecules we have described so far, each of the bonds has been a single bond;
a covalent bond in which a single pair of electrons is shared by two atoms. But it is
possible for atoms to share two or more electron pairs. A double bond is a covalent
bond in which atoms share two pairs of electrons. A triple bond is a covalent bond in
which atoms share three pairs of electrons. A comparison of the Lewis structures for
methane, CH4, ethylene, C2H4, and acetylene, C2H2 shows these bonds. Note the octet
of electrons on each carbon atom. Double bonds are most often formed by C, N, O, and
S atoms. Triple bonds from mostly to C and N atoms.
Polar Covalent Bonds; Electronegativity
When the electronegativities of the atoms in a covalent bond are not equal, electrons
are not shared equally. In this case there may be partial transfer of electron density from
one atom to the other resulting in a bond or molecule with oppositely charged ends. The
water molecule provides a familiar example of this. The greater the electronegativity
difference, the more ionic character the bond has. Bonds that are partly ionic are
called polar covalent bonds.
Nonpolar covalent bonds occur when electrons are shared equally between atoms.
This type of bond arises when the electronegativities of the two atoms are equal such
as in the case of the H2 molecule where the difference in electronegativity is 0.
Recall the electronegativity is a measure of the ability of an atom in a molecule to
draw bonding electrons to itself. Electronegativity is an experimentally derived quantity
and several scales have been proposed. The most widely used scale was developed by
Linus Pauling. These values are sometimes called "Pauling's electronegativity values"
or more simply the "Pauling scale". Pauling's values are often represented on the
periodic table as show below.
The difference in electronegativity between atoms allows us to predict what type of bond
the atoms will form as shown below.
In general, if the electronegativity difference is between 0.0 and 0.4 the bond is
nonpolar covalent. If the difference is between 0.4 and 2.0 it is polar covalent while
differences greater than two generally result in an ionic bonds.
Example: Using Electronegativities to Obtain Relative Bond Polarities
Use electronegativity values to arrange the following bonds in order of increasing
polarity: P-H, H-O, and C-Cl.
Solution:
To solve this problem you would using the Pauling scale to determine the absolute
value of the electronegativity difference between the atoms in each bond. The bonds
are then placed in order of increasing polarity based on these differences. The greater
the difference the more polar the bond.
The differences are: P-H, 0.0; H-O, 1.4; C-Cl, 0.5. The order is therefore P-H, C-Cl, HO.
HOMEWORK: Practice exercises 8.1-8.4, book questions pg. 382 questions 23, 25,
26,29,30,31,39