Covalent Bonding

Chapter 8

Covalent Bonding

8.1 Molecular Compounds

Molecules and Molecular Compounds

Covalent Bonds

Define: A bond formed when two atoms share a pair of electrons

Molecule

Define: A neutral group of atoms joined together by covalent bonds.

Diatomic Molecule:

Define: A molecule consisting of two like elements. Ex: O2

Molecular Compound:

· a compound composed of molecules. (nonmetals only, including H)

·tend to have relatively low boiling/melting points ( 300o)

Molecular Formulas

·is the chemical formula of a molecular compound

·depicts the number of atoms of each atom present in a molecule

8.2 The Nature of Covalent Bonding

The Octet Rule in Covalent Bonding

In forming covalent bonds, the octet rule is still obeyed. Electron sharing usually occurs so that atoms attain the electron configuration of the noble gases.

Single Covalent Bonds

Define: Two atoms held together by a sharing of a single pair of electrons.

The simplest compound with covalent bonds: H2

Orbital Picture:

Lewis Dot Picture:

Each hydrogen gets the stable electron configuration of helium.

Covalent bonds occur between elements which want electrons:

Another example: F

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A pair of shared electrons may also be represented by a dash ( ─ ).

These are called structural formulas.

Another example: H2O

Nitrogen trichloride: NCl3

Double and Triple Covalent Bonds

Example: the oxygen molecule, O2

In order to write a Lewis Structure, following the octet rule, oxygen must share 4 electrons.

Define:

Double Covalent Bonda bond which involves the sharing of two pairs of electrons

also:

Triple Covalent Bonda bond which involves the sharing of three pairs of electrons

Example: the nitrogen molecule, N2

Coordinate Covalent Bonds

·a special type of covalent bond

Defined:

A covalent bond in which one atom contributes both bonding electrons; the shared pair comes from one of the bonding atoms. Once formed, it acts as a normal covalent bond.

Example: carbon monoxide, CO

To indicate a coordinate covalent bond, an arrow is used in the structural formula, showing which atom contributed both electrons.

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Lewis Structures for Polyatomic Ions

Polyatomic ions contain covalent bonds within themselves, even though they are part of an ionic compound

Example NH4+

Many polyatomic ions are negatively charged meaning electrons must be added

Example: OH-

Example: SO32-

Bond Dissociation Energies

Atoms form bonds to become stable (lower energy)

Energy is required to break atoms apart

Bond Dissociation Energy:

Defined:The energy required to break a chemical bond

Example:

HH + 435kJH· + H·

See Table 15.4

Note:Double bonds have higher bond energies than single bonds

C─O356k1

C═O736kJ

CO1074kJ

Bond Energies may be used to predict ΔH of reactions

Example:

If H─Cl bond energy = 431 kJ,

calculate AH for:H2 + Cl2 2HCl

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Resonance

Defined:

A phenomenon that occurs when two or more equally valid electron dot structures can be written for a molecule or ion.

Example: ozone, O3

·both structures are different and valid

·the bonding changes while positions of the atoms do not

·the actual structure is considered to be an average or hybrid of both

Generally, the more resonance forms that can be drawn, the more stable the molecule or ion.

Exceptions to the Octet Rule

Sometimes, molecules exist which do not follow the octet rule. If there is an odd number of electrons or more/less than 8 valence electrons, the octet rule cannot be followed.

Example: NO2

Note:Both structures have an unpaired electron

Substances with one or more unpaired electrons are said to be paramagnetic

·are attracted by an external magnetic field

Substances with all paired electrons are said to be diamagnetic

·not attracted by an external magnetic field

Oxygen, O2 is found to be paramagnetic

most likely, oxygen is a resonance hybrid of:

Other exceptions to the Octet Rule:

BF3PCl5

BeCl2SF6

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8.3 Bonding Theories

VSEPR Theory

VSEPR Theory is used to determine the shapes of covalent molecules from their Lewis Structures.

Consider 5 cases:

Case #1AB2

Dot formula:

Case #2AB3

Dot formula:Picture:

Case #3AB4

Dot formula:Picture:

Case #4AB2

Dot formula:Picture:

Case #5AB3

Dot formula:Picture:

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8.4 Polar Bonds and Molecules

Bond Polarity

Not all covalent bonds are the same

·the character of the bond depends on the atoms which are joined

·the character of the molecule depends on the bonds

When atoms share electrons, they may not share them equally

·some atoms have a stronger pull on the electrons

Two types of Covalent Bonds

Nonpolar Covalent Bond: a covalent bond in which the electrons are equally shared

Examples:H2, O2, N2...

·atoms have an equal pull on the electrons

Polar Covalent Bond:a covalent bond in which the electrons are unequally shared

·one atom has a greater attraction for electrons

Electronegativity

·the measure of the attraction an atom has for electrons within a bond (see table )

Example:HCl

To indicate this:

Polar Molecules

Polar molecules: one end is more electronegative than the other.

When polar molecules are placed between oppositely charged plates, they tend to orient themselves with respect to the positive and negative plates. P.239 Fig. 8.24

H2O

Other molecules:

BeH2PCl3

CH4

The amount of polarity depends on the differences in electronegativities.

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Four levels of Polarity

Δ EN / Type of Bond / Example
0.0 - 0.4 / Covalent (nonpolar) / H─H
0.4 - 1.0 / Covalent(moderately polar) / H─Cl
1.0 - 2.0 / Covalent (very polar) / H─F
2.0 / Ionic / NaCl

Attractions between Molecules-Intermolecular Attractions

Ionic substances are held together by strong ionic bonds

·most are solids

Covalent (molecular) substances are individual molecules

·many are gases

·some are liquids or solids

Molecular substances are held together by Intermolecular Forces (forces between molecules). These forces are weaker than bonds.

Two Types:

van der Waals Forces

a.Dipole interactions the attraction between polar molecules

b.Dispersion Forces due to interactions of electrons and nuclei of neighboring molecules

Example:

Intermolecular forces help determine melting points, boiling points, solubilities, etc.

Hydrogen "Bonds"

·are the result of attractive forces between a hydrogen on one molecule and a highly

electronegative atom (N, O, F) on another

Example: H2O

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Intermolecular Attractions and Molecular Properties

Molecular substances:

·most are gases, liquids, or lowmelting solids (below 300 oC)

·do not conduct electricity

·solubility in water varies (depends on intermolecular forces)

network solids: solids that are covalently bonded and form a network: high BP/MP (diamond, SiC)

See Table 8.4 p. 244

HW: 1,2,3,4,5,6,7,8,9,10,11,12,15,16,17,18,19,21,22,24,30,31,32,34,36,37,38,39,40,41,43,44,45,46,49,50,51,52,54,57,58,59,60,61,63,64,65,68,69,70,73,75,76,79,82

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