Covalent Bonding and Molecular Structure (OWLBook Chapter 8) Covalent Bonding Lewis Structures Bond Properties Electron Distribution VSEPR and Molecular Shapes Bond Polarity OWL Due 27-November (I am aware of some of the problems in Chapter 8, they are being addressed) MarvinSketch 5 Tutorial Exam #4 (Chapters 8 and 9) about 7-December

Covalent Bonding Coulombs Law (Ionic Bonds) ◦ Describes the forces most involved in ionic bonds  Bonds based almost solely on electrostatic forces

Covalent Bonding Metallic Bonding ◦ Forces between metal atoms’ nuclei and electrons dominate bonding the atoms together Covalent Bonding ◦ Forces between nuclei and electrons in two atoms in a molecule. It is generally described as sharing of a pair of electrons between to atoms to form a single bond

Evaluating Covalent Bonding Lewis Structures ◦ Used to examine the number of valence electrons and how they may bond ◦ Generally, atoms prefer to have eight valence electrons around them when bonded  When atoms are not bonded (by themselves)  Group number for the “A” or main group group elements  More complicated for transition elements ◦ Elements preferring other than 8?  H (2), B (6), S, P, Br

Drawing Lewis Structures For atoms and ions ◦ Determine the number of valence electrons  Usually the group number  Can be evaluated from the electron configuration also ◦ Write the atomic symbol ◦ Arrange electrons about the symbol ◦ Add or subtract electrons for ions  Enclose in brackets and add charge

Chlorine and chloride ? Zinc and zinc ion ?

Lewis Structures and Molecules Lewis structures help describe modes of bonding AND shapes Try to follow the octet (8) rule when it is appropriate ◦ Evaluate the total number of electrons ◦ Determine which is the central atom (usually the single one)  Furthest from fluorine (lowest e- affinity) ◦ Add one pair (bond between atoms) ◦ Distribute electrons  Evaluate compliance with the octet rule

Examples…. Remember that a line ( - ) is the same as a pair of electrons (:) ◦ Hydrochloric acid ◦ Ammonia ◦ Boron trifluoride ◦ Ammonium ion ◦ Sulfate ion

Why are Lewis Structures Important?

Ice ?

Methane (CH 4 ) Carbon Dioxide ? Carbon Monoxide Methane (CH 4 ) Carbon Dioxide ? Carbon Monoxide

Resonance Structures Different (visually) but ordinarily equivalent Lewis Structures Help distribute electrons (and thus charge) throughout an ion or molecule ◦ Stabilizes the ion or molecule Help satisfy the Octet Rule Very important in Organic Chemistry

CH 3 COO - Valence electrons: 2(4) + 3(1) + 2(6) + 1 = 24 C is the central atom. A double bond is needed between C—O. There are two equivalent places for it, so two resonance structures are required.

Carbonate Ion ?

Bond Properties Bond length (or bond distance) is the distance between nuclei in a bond. Bond order is, defined in terms of the Lewis formula, the number of pairs of electrons in a bond. Bond energy is, defined in terms of thermodynamics, the energy required to break a bond (separate the nuclei so they no longer can “detect” each others presence) in a gas- phase molecule.

Copyright © Cengage Learning. All rights reserved. 9 | 21 Consider the propylene molecule: 134 pm 150 pm The shorter bond is the double bond; the longer bond is the single bond. One of the carbon–carbon bonds has a length of 150 pm; the other 134 pm. Identify each bond with a bond length.

Electron Distribution Formal Charge Electronegativity Polarity Resonance

Formal Charge Represents the distribution of charge in a bond ◦ Relative to the # of valence electrons each atom had originally ◦ A value assuming all electrons in a bond are equally shared Usefully in determining where electrons might be in a Lewis Dot Structure for a molecule Total of Formal Charges = Overall Charge on Molecule (0) or Ion (+ or -)

Electronegativity The tendency of an atom to draw electrons in a bond towards it ◦ Electrons are not normally shared equally

The difference in electronegativity between the two atoms in a bond is a rough measure of bond polarity. Using electronegativities, arrange the following bonds in order by increasing polarity: C—N, Na—F, O—H. For Na—F, the difference is 4.0 (F) – 0.9 (Na) = 3.1. For C—N, the difference is 3.0 (N) – 2.5 (C) = 0.5. For O—H, the difference is 3.5 (O) – 2.1 (H) = 1.4. C—N < Bond polarities: O—H < Na—F

Resonance, Formal Charge and Electronegativity All of these factor into the distribution of electrons in a molecule or ion. ◦ Formal charge shows the best distribution of electrons- more towards more electronegative atoms ◦ Electronegativity describes bond polarity ◦ Resonance helps stabilize molecules and ions Which is more likely?

In general…… Lowest value of formal charges ◦ 1,0,-1 is preferred over 3, 0, -3 Charge distributed somewhat throughout the molecule Negative formal charges on more electronegative elements Practice Problem 8.4.3

VSEPR (Valence Shell Electron Pair Repulsion) Theory Allows us to predict the shapes of non-metal containing ions and molecules Central atom is surrounded by structural electron pairs Structural e- pairs can be ◦ Non-bonding (lone pairs) ◦ Bonding (between atoms) Electron pair geometry is the arrangement of structural e- pairs around the central atom(s) Molecular geometry is the arrangement of atoms around the central atom(s)

Electron Pair Geometry (if all electrons are structural it is the same as molecular geometry)

Molecular Geometry Different than electron pair geometry when there are lone pairs Lone pairs want to be physically as far apart as possible in 3-dimensional space

Molecular Polarity For a molecule to be polar, it must contain polar bonds, and they must be distributed asymmetrically Polarity impacts ◦ Reactivity ◦ Crystal structure ◦ Solubility ◦ Properties  BP, MP, etc.

Molecular Polarity Determining Polarity ◦ Draw Lewis Structure ◦ Determine Molecular Shape ◦ Assign polarity to bonds ◦ Determine if there is an uneven distribution of charge The dipole moment is measure of polarity. Let’s draw these molecules.