chapter 8 covalent bonding. let’s review what do we already know? –what is a chemical bond?...

Chapter 8

Covalent Bonding

Let’s Review

• What do we already know?– What is a chemical bond?– What is an ionic bond?

Section 1

The Covalent Bond

Stability

• Lower energy is more stable

• Noble-Gas electron configuration

• Octet rule

Covalent Bond

• Atoms in nonionic compounds share electrons

• Covalent bond is the bond that results from sharing valence electrons

• Molecule is formed when two or more atoms bond covalently

Diatomic Molecules

• Two atom molecules are more stable than one atom

• H2, N2, O2, F2, Cl2, Br2, I2

HH

Hydrogen

H H

They Pair!!

Hydrogen

Oxygen

Fluorine

Fluorine

F F

Single Covalent Bonds

• One pair of valence electrons is shared– Pair may be referred to as “bonding” pair

• Also called sigma bonds– σ– Occurs when the shared pair is centered

between the two atoms

Bonding Orbital

• Localized region where bonding electrons are most likely found

Groups and Single Bonds

• Group 17

• Group 16

• Group 15

• Group 14

Homework (due Tuesday)

• Draw the Lewis structures for the following molecules– PH3

– H2S

– HCl

– CCl4– SiH4

• Challenge– Draw a generic Lewis Structure for a molecule formed

between atoms of group 1 and group 16

Homework continued

• Draw LDS for – CH4

– Br2

– C6H14 also written as CH3(CH2)4CH3

Multiple Covalent Bonds

• Bond Order – Refers to the type of bond

• Single Bond– Shares ONE pair of electrons

• Double Bonds– Two pairs of electrons are shared

• Triple Bonds– Three pairs of electrons are shared

The Pi Bond

• Multiple covalent bonds– Consist of at least one sigma and one pi bond

Strength of Covalent Bonds

• CB involve attractive and repulsive forces

• Balance of the force is upset the bond can break

• Several factors influence strength of cb

Bond Length

• Length depends on distance between bonded nuclei

• Bond length is the distance two nuclei at the position of maximum attraction– Determined by:

• Sizes of two bonding atoms• Number of electrons shared

Bonds and Energy

• Energy changes occur– When bonds are broken

• Energy is released• Need energy put in to break it

– Bond-dissociation energy » is the energy required to break a specific bond» Indicates strength of the bond

– When bonds are formed

Length and Energy

• Shorter the length the greater the energy

Energies of Chemical Reactions

• Total energy is determined from energy of bonds broken and formed

• Two types– Endothermic– Exothermic

Energies of Chemical Reactions

• Endothermic Reaction occurs when a greater amount of energy is required to break existing bonds in the reactants than is released when the new bonds formed.

• Endothermic Reaction– More energy to break a bond than energy

when bond is broken

Energies of Chemical Reactions

• Exothermic

Energyin

Energyout

Bond

Energies of Chemical Reactions

• Exothermic reaction occurs when more energy is released during product bond formation than is required to break bonds in reactants.

• Exothermic reaction– More energy is released than required to

break the bonds

Energies of Chemical Reactions

• Endothermic

Energyout

Energyin

Bond

Section Two

Naming Molecules

Binary Molecular Compounds

Example: N2O

1. First element in the formula is always named first, using the entire element name.

• What is the first element?• Nitrogen

Binary Molecular Compounds

2. The second element in the formula is named using its root and adding the suffix –ide.

1. What is the second element?• Oxygen

2. What will the name be?• Oxide

Binary Molecular Compounds

3. Prefixes are used to indicate the number of atoms of each element are present in the compound.

1. How many nitrogens do we have?• Two

2. What will the prefix be?• Di-

3. What is the prefix plus the element?– Dinitrogen

Binary Molecular Compounds

1. How many oxygens do we have?• One

2. What will the prefix be?• Mono

3. What is the prefix plus the element?• Monoxide

Binary Molecular Compounds

• What is the final answer?

How do we know what we are naming?

Pop Quiz

1. HCl

2. HClO3

3. H2S

4. H2SO4

5. H2ClO2

A. Chlorous acid

B. Sulfuric acid

C. Hydrosulfuric acid

D. Chloric acid

E. Hydrochloric acid

Match the following correctly, also note if the acid is binary or an oxyacid:

··Hint·· ClO3 is chlorate

Section Three

Molecular Structure

Molecular Formula

• Shows the elements symbols and subscripts

• PH3

Lewis Structure

HP

H

H

Space-filling Molecular Model

Ball-and-stick Molecular Model

Structural Formula

HP

H

H

Molecular Formula

• CH4

Lewis Structure

HC

H

H

H

Space-filling Molecular Model

Ball-and-stick Molecular Model

Structural Formula

HC

H

H

H

Lewis Structures

• BH3

• Nitrogen trifluoride

• C2H4

• Carbon Disulfide

• NH4+

• ClO4-

Announcement

• Print out chapter 8 review from teacher page.

• Complete by Friday (will have time in class tomorrow to work on it)

• Test Monday on sections 1,2,3

Resonance Structures

• Resonance– A condition that occurs when more than one

valid Lewis structure can be written for a molecule or ion

– Molecules and ions that undergo resonance behave as if there is only one structure

Classwork

• Page 260– #53

• Page 274– #84, 101, 102, 103, 104

• BONUS: 5 pts #137

Exceptions to the Octet Rule

• Odd number of valence electrons• Suboctets and coordinate covalent bonds

– Stable configuration with fewer than eight electrons present

– BH3

– Coordinate Covalent bond• One atom donates both of the electrons to be

shared with an atom or ion that needs two electrons to form a stable electron arrangement with lower potential energy.

Exceptions to the Octet Rule

• Expanded Octets– Central atoms contain more than eight

valence electrons– Considers the d orbital– Extra lone pairs are added to the central atom

for more bonds

Section Four

Molecular Shape

Importance of Shape

• The shape can determine– Physical properties– Chemical properties

• Electron densities created by overlap of orbitals of shared electrons determine molecular shape

VSEPR Model

• Valence

• Shell

• Electron

• Pair

• Repulsion

VSEPR Model

• Arrangement that minimizes the repulsion of shared and unshared electron pairs around the central atom

• Bond Angle– Angle between bonds

Hybridization

• Hybridization– A process in which atomic orbitals mix and

form new, identical hybrid orbitals

Hybridization

• With regards to molecules that have more than two atoms

• To determine the orbital hybrid– Determine the number of e- pairs shared, and

lone pairs

Hybridization

• Count like this. . . . – 1 = s– 2 = sp– 3 = sp2

– 4 = sp3

– 5 = sp3d– 6 = sp3d2

Molecular Shapes

• Linear– Example BeCl2

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

2 2 0 sp 180

Molecular Shapes

• Trigonal Planar– Example AlCl3

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

3 3 0 sp2 120

Molecular Shapes

• Tetrahedral– Example CH4

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

4 4 0 sp3 109.5

Molecular Shapes

• Trigonal Pyramidal– Example PH3

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

4 3 1 Sp3 107.3

Molecular Shapes

• Bent– Example H2O

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

4 2 2 Sp3 104.5

Molecular Shapes

• Trigonal Bipyramidal– Example NbCl5

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

5 5 0 sp3d 90; 120

Molecular Shapes

• Octahedral– Example SF6

Total Pairs

Shared Pairs

Lone Pairs

Hybrid Orbitals

Bond Angle

6 6 0 sp3d2 90; 90

Practice Problems

• Page 264– #56 through60

Section Five

Electronegativity & Polarity

Electron Affinity, Electronegativity, and Bond Character

• Electron Affinity– The measure of the tendency of an atom to

accept electrons– How attractive an atom is to electrons– Increases with atomic number within a period– Decreases with atomic number within a group

Electron Affinity, Electronegativity, and Bond Character

• Electronegativity– Derived by comparing an atom’s attraction for

shared electrons to that of a fluorine’s atom attraction for shared electrons

– Ability of an atom to attract electrons to itself within a covalent bond

Electron Affinity, Electronegativity, and Bond Character

• Bond Character– Chemical bonds between atoms of different

elements is never completely ionic or covalent– Four Types

• Mostly ionic• Polar covalent• Mostly covalent• Nonpolar covalent

Electron Affinity, Electronegativity, and Bond Character

• Bond Character– Can be predicted using the electronegativity

difference of the elements that bond

Electronegativity Difference

Bond Character

> 1.7 Mostly ionic

0.4 – 1.7 Polar covalent

< 0.4 Mostly covalent

0 Nonpolar covalent

Polar Covalent Bonds

• Polar Covalent Bonds– An unequal sharing of valence electrons

• Partial Charge– Represented by δ (Greek letter delta) – Due to unequal sharing, partial charges result

• Partial positive—the atom with the lower electron affinity

• Partial negative—the atom with higher electron affinity

Molecular Polarity

• Covalently bonded molecules– Either polar or nonpolar

• Depends on location and nature of bonds

• Nonpolar Molecules– Not attracted by electric field

• Polar Molecules– Dipoles, with charged ends– Uneven electron density = attracted by

electric field

Polarity and Molecular Shape

• Let’s look at H2O and CCl4• What shape does water take?

– Bent

• What shape does carbon tetrachloride take?– Tetrahedral

• Draw them

H2O & CCl4

Polarity and Molecular Shape

• The symmetry in CCl4 allows for a nonpolar molecule.

• There is no symmetry in H2O, so it is polar.

• What about NH3?

– It is polar.

Properties of Covalent Compounds

• Covalent compounds have strong bonds between atoms

• Attraction forces between molecules are relatively weak

• Intermolecular forces– Many types

Properties of Covalent Compounds

• Intermolecular Forces– Between nonpolar molecules

• Force is weak• Called dispersion force or induced dipole

– Between opposite charged ends of two polar molecules

• dipole-dipole force• The more polar the molecule the stronger the force

Properties of Covalent Compounds

• Intermolecular Forces– Between hydrogen end of one dipole and a F,

O, N atom on another dipole• Hydrogen bond

• Forces and Properties– Weak forces result in relatively low melting

points– Molecular substances as gases at room

temperature• O2, CO2, H2S

• Forces and Properties– Hardness

• Depends on strength of intermolecular forces• Many covalent compounds are soft

– Example: Paraffin, found in candles

Properties of Covalent Compounds

Properties of Covalent Compounds

• Forces and Properties – Solid Phase

• Molecules align to form a crystal lattice– Similar to ionic solid– Less attraction between particles– Shape affected by molecular shape– Most information has been determined by molecular

solids

Covalent Network Solids

• Covalent Network Solids– Composed only of atoms interconnected by a

network of covalent bonds• Example: Quartz and diamonds

– Structure can explain properties• Diamond

– Tetrahedral– Strong bonds– High melting point, extremely hard

The End