COVALENT BOND FORMATION

• When one nonmetal shares one or more electrons with an atom of another nonmetal so both atoms end up with eight valence electrons• There is a mutual attraction of different nuclei to the electron’s orbitals.

Characteristics of covalent compounds: MOLECULES

1.Composed of non-metals that are sharing electrons.

3. There are two forces involved:2. Composed of molecules

Intramolecular – strong covalent bond with in the molecules

Intermolecular – weak bonds between the molecules

Van der Waals forces4. Have low Melting Points (weak Van derWaals forces are breaking)

Example: Water

• Every atom has full energy levels

H OH

Intramolecular

OH H

Intermolecular

Strong covalent bonds between H & O, make up the water molecules

Weak attractive forces between different water molecules.

Intermolecular Forces

• They are what make solid and liquid molecular compounds possible.

• The weakest are called van der Waal’s forces - there are two kinds

1. (London) Dispersion forces2. Dipole Interactions

(London) Dispersion Forces

• Depend on the number of electrons • More electrons stronger forces• Bigger molecules more electrons

•Fluorine is a gas•Bromine is a liquid• Iodine is a solid

Dipole interactions

• Occur when polar molecules are attracted to each other.

• Slightly stronger than dispersion forces.• Opposites attract but not completely hooked

like in ionic solids.

H Fd+ d-

H Fd+ d-

Dipole Interactionsd+

d-

d+ d-

d+ d -

d+ d-

d+ d -

d+

d-

d + d -d+

d-

Hydrogen bonding

• Are the attractive forces caused by hydrogen bonded to F, O, or N.

• F, O, and N are very electronegative so it is a very strong dipole.

• The hydrogen partially share with the lone pair in the molecule next to it.

• The strongest of the intermolecular forces.

Hydrogen Bonding

HH

Od+ d-

d+

H HOd+d-

d+

Molecules composed of two atoms of the same element.

EX: H2 1s1 1s1

H H

Diatomic Molecules (HONClBrIF) –

** All Diatomic molecules are purely covalent or non-polar covalent because: there is an equal sharing between the two atoms.

H : H or H-H Bond angle (180)

LINEAR SHAPE

Both atoms strive to fill their 1s orbital so both hydrogen's attract the pair of bonding electrons equally

+ +:

Electrons spend time here; mutual attraction for the same electron pair.

Forms a NEW

Molecular orbital+ +* s-s bonding is the only

bonding that is non-directional

Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals

F F8 Valence electrons

8 Valence electrons

Directional Bonding: Atoms approach at 2z

Other possible bonding ( s-p, p-p, s-d or p-d)

Linear shape

(bond angle 180)

Single Covalent Bond

• A sharing of two valence electrons.• One pair of electrons shared between two

atoms

• H2 and F2 are both examples of single covalent bonding.

How many unshared pairs of electrons does I2 have?

I I6

unshared pairs

(Unshared pairs are also called lone pairs or nonbonding pairs)

Multiple Bonds

• Sometimes atoms share more than one pair of valence electrons.

• A double bond is when atoms share two pair of electrons. (4 electrons)

• A triple bond is when atoms share three pair of electrons. (6 electrons)

Double Covalent Bond

Ex: O2 :O::O: or O=OBond angle: 180 Shape: linear

Triple Covalent BondEx: N2

How many pairs of unshared electrons does O2 have? :O=O:

: :

(4)

:N:::N:N

:

Bond angle: 180

Shape: linear

How many pairs of unshared electrons does N2 have? (2)

:

:

:

::N=N:

Coordinate Covalent Bond

• When one atom donates both electrons in a covalent bond.

• Carbon monoxide• CO

OC

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide CO

OC

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide CO

OCTriple Bond

Oxygen donates a pair of electrons so both atoms now have 8 valence electrons.

:C=O:

Ex: SO2 (Sulfur dioxide)

O::O :S. .:. .: :

..O::O :S. .:. .: :

.

Ex: SO2 (Sulfur dioxide)

:O S::O:

:

::

::

Double bond

Ex: SO2 (Sulfur dioxide)

:O S::O:

:

:: : :

Double bond

Sulfur donates a pair of electrons to oxygen so all 3 atoms have complete octets!

Can this molecule be drawn another way and still be the same molecule? YES

:O S=O:

::

::

2 or more valid electron dot formulas that can be written for a molecule.

RESONANCE

Ex: Ozone (O3)

:O

:

.::

.. . ..O: O:Shares this pair of electrons

RESONANCE• 2 or more valid electron dot formulas that

can be written for a molecule.

:O::O:O:

:O:O::O::

:::

:

::

: or

:O=O O:::::

Ex: SO2 (Sulfur dioxide)

:O S::O:

:

:: : :

orO S O: ::

::: : :

:

Drawing Lewis Dot Structures• Draw a skeleton structure putting the first

atom written in the center (except Hydrogen)• Add up all the valence electrons.• Count up the total number of electrons to

make all atoms have a stable octet.• Subtract.• Divide by 2• Tells you how many bonds - draw them.• Fill in the rest of the valence electrons to fill

atoms up.

Examples

• NH3 • N - has 5 valence electrons wants

8• H - has 1 valence electrons wants

2• NH3 has 5+3(1) = 8 Valence• NH3 wants 8+3(2) = 14 Valence• (14-8)/2= 3 bonds• 4 atoms with 3 bonds

N

H

N HHH

Examples• Draw in the bonds• 3 bonds = 6 electrons (8-6=2)• All 8 valence electrons are accounted for• Everything is full

Examples• HCN C is central atom• N - has 5 valence electrons wants 8• C - has 4 valence electrons wants 8• H - has 1 valence electrons wants 2• HCN has 5+4+1 = 10• HCN wants 8+8+2 = 18• (18-10)/2= 4 bonds• 3 atoms with 4 bonds -will require multiple bonds

- not to H

HCN• Put in single bonds• Need 2 more bonds• Must go between C and N

NH C

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons (2 more to add) 10 – 8 = 2

NH C

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet

NH C

Polar Bonds

• When the atoms in a bond are the same, the electrons are shared equally.

• This is a nonpolar covalent bond.• When two different atoms are connected,

the atoms may not be shared equally.• This is a polar covalent bond.• How do we measure how strong the atoms

pull on electrons?

Electronegativity

• A measure of how strongly the atoms attract electrons in a bond.

• The bigger the electronegativity difference the more polar the bond.

• 0.0 - 0.4 Covalent nonpolar• 0.5 – 2.0 Covalent polar• >2.0 Ionic

How to show a bond is polar• Isn’t a whole charge just a partial charge• d+ means a partially positive• d- means a partially negative

(2.1) (3.0)

difference = .9 (polar covalent)

*The smaller the difference the more covalent the bond. • The Cl pulls harder on the electrons• The electrons spend more time near the Cl

H Cld+ d-

Polar Molecules

• Molecules with a positive and a negative end• Requires two things to be true The molecule must contain polar bonds This can be determined from differences in

electronegativity.Symmetry can not cancel out the effects of the

polar bonds. Must determine geometry first.

Is it polar?

• HF H=2.1F=4.0 (1.9) v. polar polar

• H2O H=2.1O=3.5 (1.4) v. polar polar

• NH3 H=2.1N=3.0 (.9) polar polar

• CCl4 C=2.5 Cl=3.0 (.5) polar non-

polar

• CO2 C=2.5 O=3.5 (1.0) v. polar non-

polar

Bond (electronegativity) Bond Molecule

MOLECULAR SHAPES

OFCOVALENT

COMPOUNDS

VSEPR THEORY

What Vsepr meansSince electrons do not like each other, because of their negative charges, they orient themselves as far apart as possible, from each other.

This leads to molecules having specific shapes.

Nonbonding electron pairs on the center atom strongly repel the bonding pairs, pushing the bonding pairs closer together

Linear

•Number of atoms = 3•Number of Bonds = 2•Number of Shared Pairs of Electrons = 2•Bond Angle = 180°

EXAMPLE:

BeF2

** All molecules with 2 atoms are linear

Bent #1

•Number of atoms = 3•Number of Bonds = 2•Number of Shared Pairs of Electrons = 2•Number of Unshared Pairs of Electrons = 2•Bond Angle = 105°

EXAMPLE:

H2O

Bent #2

•Number of atoms = 3•Number of Bonds = 2•Number of Shared Pairs of Electrons = 2•Number of Unshared Pairs of Electrons = 1•Bond Angle = 105°

EXAMPLE:

O3

Trigonal Planar

•Number of atoms = 4•Number of Bonds = 3•Number of Shared Pairs of Electrons = 3•Number of Unshared Pairs of Electrons = 0•Bond Angle = 120°

EXAMPLE:

GaF3

Pyramidal

•Number of atoms = 4•Number of Bonds = 3•Number of Shared Pairs of Electrons = 3•Number of Unshared Pairs of Electrons = 1•Bond Angle = 107°

EXAMPLE:

NH3

Tetrahedral

•Number of atoms = 5•Number of Bonds = 4•Number of Shared Pairs of Electrons = 4•Number of Unshared Pairs of Electrons = 0•Bond Angle = 109.5°

EXAMPLE:

CH4

COVALENT MOLECULES SUMMARY# ATOMS

SHAPE CENTRAL ATOM

BOND ANGLE

BONDPOLAR or NON-POLAR

MOLECULEPOLAR or NON-POLAR

2

3

3

4

4

5

Depends on bond

nolinear 180

Use

ele

ctro

nega

tivity

cha

rtD

iffer

ence

: 0-

.4 n

on p

olar

.5 –

1.7

pol

ar c

oval

ent

> 1.

7 Io

nic

linear No unshared pairs

180 Non- polar (all terminal atoms

same)bent Unshared

pairs105 polar

Trigonal planar

No unshared pairs

120 Non-polar (all terminal atoms same)

pyramidal Unshared pairs 107 polar

tetrahedral No unshared pairs 109.5

Non-polar (all terminal atoms same)