Lewis Structures & Molecular Geometries

Why Compounds Form Nobel gases do not form bonds because they have completely filled energy levels and do not need to gain or lose electrons. Other elements do form bonds and they do it so they too can have nobel gas electron configurations.

Valence Valence electrons = electrons in the highest energy level. These are the electrons that form chemical bonds. Sulfur – 1s 2 2s 2 2p 6 3s 2 3p 4 Oxygen – 1s 2 2s 2 2p 4

Ga = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 Determining Valence Electrons 1. Write down the electron configuration 2. Count how many electrons are in the highest s and p orbitals. You should get a # between 1 and These are the valence electrons = 3

Lewis Dot Diagrams Represent the valence e - for an atom using dots. Start at the top and place e - on each side going clockwise. Once there are e- on each side, begin to pair them up. Sodium 1 valence e - Magnesium 2 valence e -

Aluminum 3 valence e - Silicon 4 valence e - Phosphorus 5 valence e - Sulfur 6 valence e - Chlorine 7 valence e - Argon 8 valence e -

Elements in the same family have the same number of valence e - and the same Lewis dot arrangement

Bonding When chemical bonds form, some valence electrons are either shared or transferred between atoms. Only unpaired electrons participate in chemical bonds. Each unpaired electron can form 1 “covalent” bond. – Electrons are shared between them to complete their outer shell.

The Octet Rule Elements share or transfer electrons in chemical bonds to reach a stable configuration of 8 valence electrons.

Bonding Using Lewis Dot Diagrams and Share a pair of e - to form So that each atom has a full outer energy level A pair of shared e- can also be represented with a dashed line.

Exceptions to the Octet Rule In covalent bonds, atoms always share e - to reach a full valence shell of 8 valence e - …except… – Hydrogen only needs 2 e - in its outer energy level. – Boron only needs 6 e - in its outer energy level. – Some elements can form an expanded octet using empty d-orbitals to form bonds and have more than 8 valence e -.

5 Steps for Drawing Lewis Structures 1.Count the total number of valence electrons for all atoms. 2.Attach each atom to the central atom with a single bond (single bond = 2 shared electrons) 3.Complete the octet for the attached atoms by adding pairs of non-bonding electrons. 4.Complete the octet for the central atom by adding pairs of non-bonding electrons

5.Count the total number of electrons in your structure and compare to step one. – If the number of e - is the same, it is correct. – If you used too many e -, add double bond(s) and check your total again (usually add one double bond for each two electrons that you are over the total). – If there are extra electrons left over, add them as non-bonding electrons on the central atom. This is called an expanded octet.

Practice Exercises html html Let’s find the Lewis Dot Diagrams for: – H 2 – H 2 O – CH 4 – SO – NH 3 – CO 2

Valence Shell Electron Pair Repulsion Theory Abbreviated “VSEPR” Pairs of e - around an atom repel each other and will form an arrangement that minimizes this repulsion (i.e. spread as far apart from each other as possible). As a result, molecules tend to form predictable shapes. Lone pairs of non-bonding e - have greater repulsion than bonded pairs of e -.

Basic VSEPR Geometries Molecular GeometryABE NotationAtoms bonded to the central atom Non-bonding pairs on the central atom TetrahedralAB 4 40 Trigonal PyramidAB 3 E31 BentAB 2 E 2 or AB 2 E21 or 2 LinearAB 2 20 Trigonal PlanarAB 3 30

Expanded Octet Geometries Molecular GeometryABE NotationAtoms bonded to the central atom Non-bonding pairs on the central atom OctahedralAB 6 60 Square PyramidAB 5 E51 Square PlanarAB 4 E 2 42 Trigonal Bipyramid AB 5 50 SeesawAB 4 E41 T-ShapedAB 3 E 2 32 LinearAB 2 E 4 24

Practice as a Class CH 4 NH 3 H 2 O CO 2 SF 6 I 3 - AB 4 AB 3 E AB 2 E 2 AB 2 AB 6 AB 2 E 3