102 * Lab #1

Lewis Structures

Introduction

Drawing Lewis Structures

Drawing Lewis structures is best explained by working through an example. We will draw the Lewis structure of CO₂.

1. Draw the skeletal structure for the molecule or ion, joining atoms by single bonds. The single atom will always go in the center bonds.

AB₄ AB₃ CO₂ O - C - O

2. Count the number of valence electrons. Carbon has four valence electrons (Group IV A) and oxygen has 6 valence electrons (Group VI A). Therefore, CO₂ has a total of 16 valence electrons. The charge on a polyatomic ion must be taken into consideration. Remember that the negatively charged ions (anions) have gained electrons and positively charged ions (cations) have lost electrons.

3. Each single bond in step 1 accounts for 2 valence electrons. Subtract two valence electrons for each single bond and distribute the remaining electrons as non-bonding pairs of electrons around the atoms so that each atom has eight electrons, if this is possible. (For example, a hydrogen can only have 2 electrons associated with it.) For CO₂ , you have 12 electrons (16 alence e^- minus 4 bonding e^-) to distribute.

When there are too few electrons to obtain an octet of electrons for each atom in step 3, you must then make bonding pairs of electrons out of nonbonding pairs of electrons. That is, you must form multiple bonds.

Molecular Geometry and bond angles

The pairs of electrons around the central atom are positioned as far apart as possible to reduce the electrostatic repulsion. There are different types of molecular geometry based on how many pairs of electron pairs are around the central atom.

Linear geometry with a bond angle of 180°

Tetrahedral Geometry (remember that you are working in three dimensions) with a bond angle of 109.5 °

Angular Geometry with a bond angle of approximately 120° ( remember that you have three sets or groups of electrons surrounding the sulfur atom )

Trigonal Pyramid Geometry with bond angles of approximately 109 ° . For other geometry's see your textbook. Remember that elements will empty d orbitals can expand their octets and can have more than eight electrons.

Polarity

If there is a difference in electro negativity between the two atoms joined by a bond, the bond will be polar. How polar the bond will be is determined by how great the difference in electro negativity is. A polar bond does not mean that the molecule is “polar”. For a diatomic molecule a polar bond automatically implies a polar molecule, but for other molecules the geometry of the molecule must be known before making such a decision.

Polar bonds may cancel each other out. Examples are carbon dioxide and methane.

and

Lewis structures will not predict polarity. Molecules that can be drawn to appear nonpolar must be put in three dimensions with their correct geometry to determine the polarity. Example CH₂F₂ appears to be nonpolar until the correct geometry is considered.

CH₂F₂ is definitely polar.

Resonance

With certain molecules or ions, given the atomic geometry, it is possible to satisfy the octet rule with more than one bonding arrangement. The proceeding structures are called resonance structures:

Molecules or ions that have two or more resonance structures are said to exhibit resonance. The actual bonding in such molecules or ions is thought to be an average of the bonding present in the resonance structures and (for the above example) might be represented as

The stability of molecules or ions exhibiting resonance is found to be higher than that anticipated for any single resonance structure.

Procedure

Obtain one of the CD’s from your instructor and follow the directions on the next page. Working through the examples on both the Discover Chemistry and the Zumdahl CD will assist greatly in writing the Lewis structures as well as in visualizing the molecules or ions.