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 BondingCoulombs Law (Ionic Bonds)

◦Describes the forces most involved in ionic bonds Bonds based almost solely on

electrostatic forces

Covalent BondingMetallic 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 BondingLewis 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 StructuresFor 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 MoleculesLewis structures help describe

modes of bonding AND shapesTry 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 (CH4)

Carbon Dioxide ?

Carbon Monoxide

Resonance StructuresDifferent (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 RuleVery important in Organic

Chemistry

CH3COO-

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.

C

H

H

H CO

O

-

C

H

H

H CO

O

-

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.

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Consider the propylene molecule:

C C C

H

H

H

H

H

H

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 DistributionFormal Charge

Electronegativity

Polarity

Resonance

Formal ChargeRepresents 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 -)

ElectronegativityThe 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 ElectronegativityAll 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, -3Charge distributed somewhat

throughout the moleculeNegative 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 GeometryDifferent 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 PolarityFor 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 PolarityDetermining

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.