Ch. 4 lecture

Covalent bonds form when atoms share

electrons to complete octets.

Covalent bonds are typically between two

nonmetal atoms.

2/14/13

CHM 101 Ch 4: Covalent Compounds

Covalent Compounds

1. Element with lower group number is

named first.

2. If elements are in the same

group, element with higher period

number is named first.

3. Second element named as root + -ide.

4. Greek prefixes are used to designate

the number of atoms.

Binary Covalent Compounds – molecule that contains atoms of only 2 elements

CS2, SO3, PCl5, etc.

General Rules for Naming

You must know these!

Naming Covalent Compounds

What is the name of SO3?

1. The first nonmetal is S sulfur.

2. The second nonmetal is O named oxide.

3. The subscript 3 of O is shown as the prefix tri.

SO3 sulfur trioxide

The subscript 1 (for S) or mono is understood.


1. PCl5

2. CS2

3. P4S3

Give names for the following binary compounds:

Give molecular formulas for the following compounds:

1. Dinitrogen monoxide

2. Selenium hexafluoride

3. Dichlorine heptaoxide


phosphorous pentachloride

carbon disulfide

tetraphosphorous trisulfide

N2O

SeF6

Cl2O7

Use the table below to quiz yourself. Use the formula, write the name. Use

the name, write the formula.


Covalent Bonding

Octet Rule: Main group elements gain, lose, or share e- to achieve a

stable e- configuration with 8 valence e- (except H and He – only need

2 valence e- for stability)

Main group elements

Usually make the number of bonds necessary to have noble gas

configuration

Noble gases (except He) have 8 valence electrons

Lewis Dot Symbol - Representation of the number of valence electrons

in an atom. Usually only used for main group elements. Maximum

number is 8.

X

H He

Li Be B C N O F Ne

Covalent Bonding

Cl2 molecule

Cl: 7 valence electrons, needs to make 1 bond to have octet

Cl Lewis Dot Symbol

Lone pair of e- (unshared)

ClCl

Bonding pair of e- (shared)

ClCl

ClCl

Lewis Dot Structure

Count the “shared” e- for both atoms

Covalent Bonding

Lewis Dot Structure – shows

how the atoms of a molecule

are connected. Shows lone

pairs of electrons and bonding

pairs of electrons.

In carbon dioxide, CO2, the C atom shares 4 electrons with

each O atom in a double bond.

Covalent Bonding

Multiple Bonds – Double & Triple

In nitrogen molecule, N2, each N atom shares 6 electrons

with the other N atom in a triple bond.

Multiple Bonds – Double & Triple

Covalent Bonding

Group 7(A) 7 1 HCl

Group 6(A) 6 2 H2O

Group 5(A) 5 3 NH3

Group 4(A) 4 4 CH4

# Valence e- # Bonds for Octet Example

We can often predict the number of covalent bonds an atom will form

based on the number of valence electrons.

ClH OCl HOH

Covalent Bonding

2. Arrange atoms next to each other and connect with bonds

Central atom is usually written first in the formula and has lower group

number

(If atoms are in same group, central atom is from higher period)

Rules for Drawing Lewis Dot Structures

1. Count the total number of valence electrons in the molecule

NF3 N: 5 e-

F: 7e- x 3 = 21 e-

26 val e- total for the molecule

N F

F

F26 val e-

- 6 bonding e-

20 remaining e-

Covalent Bonding


3. Place lone pairs around each atom to satisfy octet rule, starting with

terminal atoms

20 remaining e-

-20 lone e-

0 remaining e-

N F

F

F

Make sure each atom has an octet. If it doesn’t, check your work.

Covalent Bonding

Practice: Draw Lewis Structures for CH3Br and H2Se.

If the central atom has < 8 e-…

5. Change single bonds to multiple bonds (double or triple) using lone

pairs from terminal atoms.

H2CO, CH3COOH, HCN

Covalent Bonding


IF there are electrons remaining…

4. Place leftover e- on the central atom. It is ok to exceed the octet

(have more than 8 electrons on an atom) if the central atom is in

Period 3 or higher.

SF4

Draw the Lewis structure for formaldehyde: H2CO

1. Count the total number of valence electrons in the molecule

H2CO H: 1 e- x 2 = 2 e-

C: 4e- x 1 = 4 e-

O: 6e- x 1 = 6 e-

12 val e- total for the molecule

2. Place atoms relative to each other and connect with bonds

H can only make 1 bond, so it has to be a terminal atom. C is a

very common central atom (lower group # than O).

12 val e-

- 6 bonding e-

6 remaining e-

C HH

O

Covalent Bonding

Draw the Lewis structure for formaldehyde: H2CO

3. Place lone pairs around each atom to satisfy octet rule, starting with

terminal atoms

H can only accommodate 2 electrons, which it has with the bonding

electrons. Never put lone pairs on H.

6 val e-

- 6 bonding e-

0 remaining e-

C HH

O

The central atom (carbon) has < 8 e-

4. Change single bonds to multiple bonds (double or triple) using lone

pairs from terminal atoms.

C HH

O

Carbon and oxygen now both have

“octets”.

Covalent Bonding

Lewis Structures of Ions

For anions – add electrons to total valence equal to charge

NO3– : N: 5e- x 1 = 5

O: 6e- x 3 = 18

+ 1e- (b/c of ion charge)

24 e- total

- 6 e- bonding

18 e-

- 18 e- lone pairs

0 e-O N O

O

O N O

O– Does each atom have an octet?

For ions, always put the Lewis structure in

brackets and write the charge as superscript.

How can we give N an octet?

For cations – subtract electrons from total valence equal to charge

NH4+ : N: 5e- x 1 = 5

H: 1e- x 4 = 4

- 1e-

8 e- total

- 8 e- bonding

0 e-

H N H

H

H+

Examples for you to practice: ClO3–, ClF4

+

Lewis Structures of Ions

Exceptions to the Octet Rule

1. Fewer than 8 electrons: Molecules with Be or B as central atom are often electron deficient. Be usually only needs 4 electrons for stability and B only needs 6.

BeCl2 BF3

Covalent Bonding

2. Odd # of valence electrons:

NO2:

Free radicals: atoms or molecules with unpaired electrons, highly reactive

3. More than 8 valence electrons: “expanded valence shells”

- only allowed for atoms in Period 3 and beyond

H2SO4

Exceptions to the Octet Rule

Covalent Bonding

O S O

O

O

H H

Example for you to practice: PCl5

Draw the Lewis Structure for ozone, O3:

There are 2 possible structures only differing in the location of electrons and the double bond.

Experimental data says that both bonds are identical.

The actual structure of O3 is neither of them, but a composite of the two, called a resonance hybrid. Each Lewis structure is a resonance structure of O3.

Covalent Bonding

Resonance Structures

O

O+

O-

O

O+

O-

Use a double-headed arrow for resonance structures

O

O+

O-

O

O+

O-

• Resonance structures differ only in the assignment of e- pair positions, not in atom positions.

Draw resonance structures for NO3-:

Covalent Bonding

Resonance Structures

Valence-Shell Electron-Pair Repulsion (VSEPR) Theory

Bonding and lone pair electrons surrounding a central atom repel each

other. To minimize the repulsions, electron groups are oriented as far

apart as possible.

Molecular Shape – VSEPR Theory

The properties of a molecule are heavily influenced by molecular shape.

How many electron “groups” are on a central atom? Each of the following

counts as one electron group:

• Lone pair of e-

• Single bond

• Double bond

• Triple bond

Covalent Bonding


Covalent Bonding

Ideal bond angles

Covalent Bonding

Electron pair geometry – Arrangement of e- groups (bonds and lone pairs) around the central atom.

Molecular geometry – Arrangement of atoms in space, shape of molecule. This is different than e- group geometry if lone pairs are present

Linear electron pair geometry

2 atoms attached to a central atom

180 bond angle

Can’t have any lone pairs, so electron pair geometry is same as molecular geometry

Examples: BeCl2, CO2, CS2, HCN

If there are 2 electron “groups” on the central atom…

Linear molecular geometry

Covalent Bonding


Trigonal Planar electron pair geometry

If 0 lone pairs, electron pair geometry is same as molecular geometry

Examples: BF3, SO3, NO3–, CO3

2–


120 bond angle

If there are 3 electron “groups” on the central atom and 0 lone pairs…

Trigonal Planar molecular geometry

Covalent Bonding


Tetrahedral electron pair geometry

Electron pair geometry

Molecular geometry

Example


109.5 bond angle

If 0 lone pairs, electron pair geometry is same as molecular geometry

Examples: CH4, SiCl4, SO42–, ClO4

–

If there are 4 electron “groups” on the central atom & 0 lone pairs…

Tetrahedral molecular geometry

Covalent Bonding



Molecular geometry

Example

3 atoms & 1 lone pair attached to a central atom

109.5 bond angle

Examples: NH3, PF3, ClO3–, H3O

+


If there are 4 electron “groups” on the central atom & 1 lone pair…

Trigonal pyramidal molecular geometry

Covalent Bonding



Molecular geometry

Example

2 atoms & 2 lone pairs attached to a central atom

109.5 bond angle

Examples: H2O, SCl2


If there are 4 electron “groups” on the central atom & 2 lone pairs…

Bent (Angular) molecular geometry

Covalent Bonding


Covalent Bonds & Electronegativity

Electrons are shared in covalent bonds, but they usually are not shared

equally.

One atom usually pulls the electrons more strongly than the other.

The bond between H and Cl

in HCl is covalent.

Cl pulls the shared electrons

more strongly than H.

This creates a “partially

negative” region around Cl

and a “partially positive”

region around H.

The bond between H and Cl is a polar covalent bond.

Covalent Bonding

WHY???

Covalent Bonds & Electronegativity

Covalent Bonding

The electronegativity value for an atom indicates the attraction of that atom

for shared electrons (in covalent bonds).

Electronegativity strength

usually increases as the size of

an atom decreases.

A nonpolar covalent bond is an equal or almost equal sharing of electrons.

The atoms involved have almost no electronegativity difference (0.0 to 0.4).

Examples:

Electronegativity

Atoms Difference Type of Bond

N-N 3.0 - 3.0 = 0.0 Nonpolar covalent

Cl-Br 3.0 - 2.8 = 0.2 Nonpolar covalent

H-Si 2.1 - 1.8 = 0.3 Nonpolar covalent

H-C ??? ???

Nonpolar Covalent Bonds

Covalent Bonding

Diatomic molecules are molecules that

contain 2 atoms of the same element.

They are the natural state for elements

H, O, N, Cl, Br, I, and F.

These are the only truly nonpolar covalent

bonds.

Nonpolar Covalent Bonds

Covalent Bonding

A polar covalent bond is an unequal sharing of electrons.

The atoms involved have a moderate electronegativity difference (0.5 to 1.7).

Examples: Electronegativity

Atoms Difference Type of BondO-Cl 3.5 - 3.0 = 0.5 Polar covalentCl-C 3.0 - 2.5 = 0.5 Polar covalent

O-S 3.5 - 2.5 = 1.0 Polar covalent

Polar Covalent Bonds

Covalent Bonding

An ionic bond occurs between metal and nonmetal ions, and is a result of

electron transfer from the metal to the nonmetal.

There is a large electronegativity difference (1.8 or more) between the

atoms.

Examples:

Electronegativity

Atoms Difference Type of Bond

Cl-K 3.0 – 0.8 = 2.2 Ionic

N-Na 3.0 – 0.9 = 2.1 Ionic

S-Cs 2.5 – 0.7= 1.8 Ionic

Ionic Bonds

Covalent Bonding

Covalent Bonding

A covalent bond is polar if it connects atoms with different electronegativity

values, i.e. H-Cl, C-F, C-Cl, etc.

How do you know if a molecule is polar?

- Bond Polarity

- Molecular Shape

Covalent Bonding

Molecular Polarity

Determine the molecular polarity of CO2:

C OO

Each C=O bond is polar, but the molecule is linear and the two

dipoles cancel each other. Therefore, the CO2 molecule is nonpolar.

COS (carbonyl sulfide):

C SO

C and S have the same EN. There is only 1 dipole, pointing towards

O. Therefore, the COS molecule is polar.

Covalent BondingMolecular Polarity

What about molecular shape?

H2O:

O HH

Is the molecule nonpolar?

O

H H

NO! The H2O molecule is not linear, it is tetrahedral!

Covalent Bonding

Molecular Polarity

Indicate the dipole moment (if any) for CF4 and CHF3.

Covalent Bonding

Molecular Polarity

Indicate the dipole moment (if any) for CH2Cl2.

Covalent Bonding

Molecular Polarity

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Ch. 4 lecture