From Last Class…

• Ionic and covalent bonds• How to name them• Electronegativity

• Scale to determine bond type

From Last Class…

• Ionic and covalent bonds• How to name them• Electronegativity

• Scale to determine bond type

| | | 0 0.5 1.7 3.3

Structures

The Octet Rule

• All elements want to get a full outer shell• Will lose or gain electrons to get to 8 electrons

with the exception of Hydrogen and Helium• atoms bond in order to achieve the same

number of electrons as the nearest noble gas in the periodic table.

• bonded atoms are said to be isoelectronic to the nearest noble gas. (All noble gases have 8 valence electrons except He.)

Lewis Dot Structure

• Uses the octet rule• Represent the valence electrons of an

element• You can follow these rules…

1. Decide how to arrange the various atoms with respect to one another.– Although there are no definite "rules" you can follow, you usually do the following:– Go for maximum symmetry. For instance, if you have one of Atom A and four of Atom

B, put Atom A in the middle and arrange the four B atoms symmetrically around it.– Carbon is always a central atom in a Lewis structure. In compounds with more than

one carbon atom, the carbon atoms are joined in a chain.– Hydrogen atoms always go on the outside. Not too hard to figure out why!!– Halogen atoms form only single bonds when oxygen is not present, therefore they are

usually placed on the “outside”. When halogen atoms are on the “outside”, they are always joined by single bonds and always obey the octet rule.

– Oxy-acids are acids containing oxygen (how surprising!). Some examples are HNO3, H2SO4, H3PO4 and HClO3. When writing the Lewis structure for an oxy-acid, always bond the hydrogen(s) to an oxygen.

2. Count all the valence electrons. – Determine the total number of valence electrons in the compound.

3. Place two electrons (a sigma covalent bond) between the central atom and each of the outer atoms

4. Place "lone pairs" of electrons about each terminal (outer) atom [except H atoms] to satisfy the octet rule.

5. Count how many electrons you have used to this point. Place the extra electrons on the central atom(s) in pairs.

6. If the central atom does not have an “octet”, try forming double or even triple bonds until it does.

Time to try a few

VSEPR Theory

• What it stands for: Valence Shell Electron Pair Repulsion Theory

• Predicts what the molecules will look like in 3D• Developed in 1957 by Canadian chemist, Ronald Gillespie

• Based on the idea that the geometry of a molecule is determined primarily by repulsion among the pairs of electrons associated with the central atom. The pairs can be bonded or un-bonded (lone pairs)

• Only valence electrons of the central atom influence the molecular shape

LONE PAIR

BONDED PAIR

1) Pairs of electrons in the valence shell of a central atom repel each other.

2) These pairs of electrons tend to occupy positions in space that minimize repulsions and maximize the distance of separation between them.

Grab a MolyMod Kit!

-There are three kinds of repulsions between valence electrons of a molecule

-LP LP repulsions are stronger than LP BP repulsions which are stronger than BP BP repulsions

LONE PAIR – LONE PAIR (LP, LP)

LONE PAIR – BONDED PAIR (LP, BP)

BONDED PAIR – BONDED PAIR (BP, BP)

Example #1SHAPE – Bent

-The two lone pairs have a stronger repulsion than the bonded pairs (hydrogen) so the Hs and pushed downwards

Example #2

SHAPE – Tetrahedral

ANGLE – 109.5º

- Used to show the 3D shape of CH4

* Shapes of Covalent Molecules with NO Lone Pairs of Electrons Around the Central Atom*

Bond Angle - 180º# OF Electron Pairs – 3Examples – BeCl2, CO2, HgCl2

Bond Angle - 120º# of Electron Pairs – 3Examples – BCl3, BF3, SO3

Bond Angle – 109.5º# of Electron Pairs – 4Examples – CH4, ClO4, PO4, SO4

Bond Angle – 90º + 120º# of Electron Pairs – 5Examples – PCl5

Bond Angle - 90º# of Electron Pairs – 6Examples – SF6

Planar triangular Tetrahedral

Trigonal bipyramidal Octahedral