Chemistry 103A; Sections 5, 6, 7, 8; Lecture 28, 30 Oct 00

Recall:

We were talking about:

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.

If the bond is covalent and the electronegativity difference is not zero we say that the pond is polar covalent.

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

The bond in O2 is covalent, but it is not a polar covalent bond. It is a nonpolar covalent bond.

In the last two cases one of the atoms is more electronegative that the other atom. The most electronegative atom gets more than its share of the electrons so it has a partial negative charge.

In the C- N bond the N is more electronegative than the C so the N gets more than its share of the electrons. Sometimes we indicate this difference in charge distributions by writing,

(d +)C- N(d- ).

The d + means that the C atom has a partial positive charge. (This is not a full unit charge, but only a fraction of a unit charge.) The d- means that the N atom has an equal and opposite partial charge.

In the C- H bond the C is more electronegative than the H so this time the C has the partial negative charge. Write,

(d- )C- H(d +).

 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-
(Recall that 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 end?

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 form what we call a "sigma bond" (s -bond). In a s -bond the electron density builds up between the nuclei of the two atoms involved. 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 s -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, but for now we have to tell you that the second pair in a double bond forms what we call a pi bond [p -bond]. In a p -bond the electrons density is not between the nuclei, but off of the internuclear axis.

In a triple bond one of the pairs forms a s -bond and the other two pairs of electrons form TWO p -bonds. But the two p -bonds are perpendicular to each other.)
 

 For purposes of VSEPR we only count the s -bond part of a double or triple bond. The other p -bond pair in a double bond or two p -bond pairs in a triple bond do not contribute to the shape of the molecule. The principal is then that the lone pairs and s -bonded pairs (including the s -bond parts of double or triple bonds) will try to arrange themselves as far away from each other as possible.

In summary, to determine the shape of a molecule we count only lone pairs and s -bonded pairs.

The rest is simple geometry.