Recall:

Ionic Compounds and "Ionic bonds"

In ionic compounds the crystal is held together by electrostatic forces (the

attraction between the + and - charges).

In a crystal of NaCl there are no NaCl "molecules." The crystal consists of an array of Na⁺ and Cl^-ions.

CaF₂ is an ionic compound. There are no CaF₂ "molecules" in the crystal. However, there must be twice as many F^-ions as Ca²⁺ ions in the crystal.
 
 

Lewis Dot Symbols and Formulas

Lewis dot formulas help us visualize how the octet rule works in the formation of molecules.

In Lewis dot formulas the atom is represented by its symbol and the valence

electrons are represented by dots placed around the symbol of the element.
 
 

Covalent Bonds (sharing electrons)

Elements that don't make ions easily can obtain the noble gas configuration by sharing electrons.

H· ·H gives H:H

The ":" stands for a shared pair. Each H atom thinks it has the He configuration. The shared pair of electrons is called a covalent bond.

Lewis dot formulas were invented to help us show how electrons are shared.

· = valence electron

- = shared pair = :

write H₂ as H:H or H- H.

A shared pair of electrons is called a covalent bond.

Examples of molecules,

F₂ ,  NH₃,   CCl₄,   CH₃OH,   CHClBrI,   C₂H₆

Where are the shared electrons?

They are in molecular orbitals. Molecular orbitals are electron clouds that encompass more than one nucleus. (We will talk more about molecular orbitals in Chapter 10.)

Atoms can share more than one pair

Double Bonds

It is possible and common for some single bonds to be formed with both of the electrons in the electron pair coming from the same atom. Such a bond is called a coordinate covalent bond.

Examples of molecules with coordinate covalent bonds are:

Rules for Drawing Lewis Dot Structures

One way to draw Lewis dot structures is to write the atoms (with their Lewis dot valence electrons) in the proper arrangement to form the molecule.

Then use circles to show how the electrons are paired to form bonds.

Form all the single bonds first.

Then form any double or triple bonds.

There is an alternate method to draw Lewis dot structures given beginning on page 383 of our text. The book's method works, but it does not make it quite as obvious how the electrons are shared.

In either case we have to know the arrangement of the atoms in the molecule. Practice and experience help us determine how the atoms are arranged, but there are some general rules that are helpful.

NH₃

ClO^-

NO₂⁺

PO₄^3-

For neutral compounds:

C always has four bonds

N always has three bonds

O always has two bonds

F and H always have one bond

For negative molecular ions:

N^(-) will have two bonds

O^(-) will have one bond

For positive ions:

N⁽⁺⁾ will have four bonds

O⁽⁺⁾ will have three bonds

Isoelectronic Series

We often find groups of molecules with the same electronic structure. We say that these species are isoelectronic.

Examples:

NO⁺, N₂, CO, CN^-

BH₄^-, CH₄, NH₄⁺

NH₃, H₃O⁺

CO₂, OCN^-, SCN^-, N₂O, NO₂^-, COS, CS₂

Resonance

Several of the Lewis dot structures we have written could have been written in two or more ways.

For example:

SO₂, SO_3, O₃, HNO₃, NO₃^-, CO₃^2-, etc.

The answer is: none of them and all of them.

That means that none of the structures is correct by itself, but that the correct structure is a "mixture" of all of them.

This process of "mixing" structures is called resonance.

The mixture is sometimes said to be a "linear combination" of all the possible structures.

The official name for this linear combination or mixtures is: resonance hybrid.

The individual structures which combine to make the resonance hybrid are referred to as "contributing structures."

Resonance confers extra stability on a molecule. This is because electrons don't like to be confined so the more room you give an electron to move around the more stability you gain.

The rules are:

You can't move nuclei to make a resonance form.

You can only move electrons.

All of the contributing forms must be valid Lewis dot structures.