Chemistry 103A; Sections 5, 6, 7, 8; Lecture 27, 25 Oct 00

Recall:

We were talking about Bond Order

Bond order is a measure of how many pairs of electrons are shared between two atoms to bond those two atoms together. Recall that a chemical bond is a sharing of electron pairs between two atoms.

For molecules without resonance the bond order is easy.

A single bond has a bond order of 1 because there is one pair of electrons being shared between the two atoms.

A double bond has a bond order of 2 because the two atoms share two pairs of electrons.

Likewise, in a triple bond the bond order is 3 because three pairs of electrons are shared between the two atoms bound together.

 Things are a little more complicated in molecules with resonance. To determine the bond order when there is resonance we must recognize that some of the electron pairs are split between more than one pair of atoms. For example, O3 is a resonance hybrid molecule.

The bond order of each OO is 1.5. The 1 from the single bond and 0.5 from the pair involved in the resonance.

Examples:

SO2,

SO3.

 In more complicated cases we can calculate the bond order by a formula.

If we let natom pairs = the number of atom pairs of a specific type involved in the resonance, and nelectron pairs = the number of pairs of electrons involved in the resonance then the bond order is nelectron pairs/natom pairs.

Examples:

NO3-

C6H6.
 
 

Bond Length

The bond length is defined as the distance between the nuclei of the two atoms involved in the bond.

There are three effects that govern bond lengths:

Bonds with higher bond orders tend to be shorter.
So, for example, a CºC bond is shorter than a C=C bond which is shorter than a C-C bond.

Likewise, a carbon-carbon bond in benzene (C6H6) is shorter than an unhybridized carbon-carbon bond, but longer than an unhybridized C=C bond.
 

 Bonds lengths tend to increase as you go down a column in the periodic table. That is, bonds between larger atoms tend to be longer.
  An S- H bond would be expected to be longer than an O- H bond because S is a larger atom than O.

Likewise, the H- I bond is longer than the H- Br bond which is longer than the H- Cl bond which in turn is longer than the H- F bond.
 

 Bond lengths tend to decrease as you go to the right in a row in the periodic table.
  C- H bonds are longer than N- H bonds are longer than O- H bonds which are longer than F- H bonds.  
(There are two tables of average bond lengths on page 401 of your text to help you check these rules. I will not ask you to compare bond lengths between the two tables, but only within a table.)


Bond Energy

The bond energy is the energy required to break a chemical bond. That is, it is the energy required to split the molecule into two parts at the bond in question.

Breaking chemical bonds is an endothermic process. So DH for breaking bonds is positive. (Making a bond is the reverse of breaking a bond. Making bonds is an exothermic process, with a negative DH.) Bond energies are usually given for a mole of bonds.

We will represent the bond energy (per mole) by the symbol, D (for bond Dissociation energy). Since bond dissociation energies are endothermic, all Ds are positive.

We can use bond energies to make estimates of DHs of chemical reactions.

DHreaction = SDbonds broken - SDbonds formed.
Example:
N2(g) + 3 H2(g) ® 2 NH3(g).
The bond energies are (from the table on p 403 of our text),
DNº N = 945 kJ/mol

DH- H = 436 kJ/mol

DN- H = 391 kJ/mol.


Electronegativity

Some electron pairs in a covalent bond are not shared equally.

(The only time electron pairs in a bond are shared equally is when the bond is between two atoms of the same element.) The ability of an atom to pull electrons toward itself is called electronegativity.

The more electronegative atoms get a larger share of the electrons.

General Trends

Electronegativity increases from left to right across a row on the periodic table.

Electronegativity decreases as you go down a column on the periodic table.

 The most electronegative element is fluorine. The least electronegative element is Fr.

Electronegativity differences tell us whether bonds are ionic or covalent. (There is a table of electronegativities in your text - page 410. We do not memorize electronegativities.)

If the electronegativity difference is 1.9 or less, the bond is (polar) covalent.

If the electronegativity difference is 2.0 or greater, the "bond" is ionic.

Polar covalent means that the electrons are not shared equally.

Examples:

NaCl    3.0  -  0.9 = 2.1

MgO    3.5  -  1.2 = 2.3

O2       3.5  -  3.5 = 0.0

C-N    3.0  -  2.5 = 0.5

C-H    2.5  -  2.1 = 0.4

Polar covalent bonds form a dipole. We say that they have a dipole moment.

A dipole moment is a charge multiplied by a distance (the distance between the separated charges).

d+ ¬ d ® d-


(d+ and d- are equal and opposite partial charges.)

Dipole moment = d ´ d

Dipole moment is a vector quantity. That is, it has a magnitude and a direction.

When we add dipole moments we must include the direction information.

Molecules with a dipole moment are called polar molecules.

Examples:

HF (The most electronegative element is the negative end.) H2

O2

N2

CO2

H2O

NH3

CH4

 So far we have only been talking about the polarity of individual bonds. What is of most importance to us is the polarity of a molecule. That is, is one end of a molecule more negative that the other?

We can't determine the polarity of a molecule until we learn how to determine the shapes of molecules.
 

Molecular Shapes

The shape of a molecule is very important in determining the chemical and physical properties of the molecule (compound).

We will determine the shapes of molecules using a method called VSEPR, which stands for Valence Shell Electron Pair Repulsion.

VSEPR is quite good at predicting the shapes of molecules - particularly involving elements of the first row. To do better you have to go to much more sophisticated - and complicated - theories or to actual physical measurements of the shapes.

The main idea of VSEPR is that electron pairs are pretty stable and do not like to be crowded. Electron pairs try to get as far away from each other as possible (recognizing that they have to stay close to the atom because it is the positive charges on the nucleus which prevent them from "flying" away).

We use the Lewis dot formulas to tell us how many electron pairs there are and where the pairs are.

Around any given atom in a molecule we have several types of electron pairs:

Lone pairs - these are electron pairs which are not involved directly in bonding.

Single bonded pairs - these are electron pairs which give a single bond between two atoms.

Single bonded pairs which are part of a double bond or triple bond.
 

(The electron pairs in double and triple bonds are "doing different things." One of the pairs is a normal single bond, but the other pair or other two pairs are arranged differently around the atom. We will look at this in more detail in Chapter 10.)
  For purposes of VSEPR we only count the single bond part of a double or triple bond. The other (double bonded) pair or (triple bonded) two pairs do not contribute to the shape of the molecule. The principal is then that the lone pairs and single bonded pairs (including the single bond parts of double or triple bonds) will try to arrange themselves as far away from each other as possible.