Covalent Bonding

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Covalent Bonding

Covalent Bonding

• Sharing of electrons between two atoms

• Both atoms are usually nonmetals

• Two electrons are shared it is a single bond

• Four e - are shared it is a double bond

• Six e - are shared it is a triple bond

• Coordinate covalent bond is when both

electrons being shared come from one of the

atoms


• O 2

Lewis structures


Lewis structures

• O 2

• These two electrons are shared, so how many electrons

does each atom have


Lewis structures

• O 2

• Since each atom only has 7 e- some electrons have to

move to share more electrons until each has 8 electrons


Lewis structures

• O 2

• Since each atom only has 7 e- some electrons have to

move to share more electrons until each has 8 electrons


Lewis structures

• O 2

• Now each atom has 8 electrons so oxygen has a double

bond


How do you know where to put

atoms in a lewis structure

• The atom that has the most bonding

potential is the one that goes in the middle.

• This is the atom that has the most unpaired

electrons

• Atoms that have only one unpaired electron

must be on the outside of the structure.

ie) H,Li,Na,Cl,F


H 2 O

• Oxygen has two unpaired e- and hydrogen

only has one so which one goes in the

middle

• Oxygen


H 2 O

• Oxygen has two unpaired e- and hydrogen

only has one so which one goes in the

middle

• Oxygen


H 2 O

• Oxygen has two unpaired e- and hydrogen only

has one so which one goes in the middle

• Oxygen

• Hydrogen only has to have two electrons because

it is not big enough to hold 8


NH 3

• Nitrogen is in the middle because it has 3

unpaired electrons


NH 3

• Nitrogen is in the middle because it has 3

unpaired electrons


CO 2

• Carbon is in the middle because carbon has

4 unpaired electrons and oxygen only has

two


CO 2

• Carbon is in the middle because carbon has

4 unpaired electrons and oxygen only has

two


CO 2

• Carbon only has six electrons and each

oxygen now has seven so electrons have to

be moved to share


CO 2

• Carbon only has six electrons and each

oxygen now has seven so electrons have to

be moved to share


CO 2

• So CO 2 has double bonds


Coordinate Covalent Bond

• NH 4

+


Where does the last hydrogen have to go


Coordinate Covalent Bond

• NH 4

+

Since we are dealing with NH4+ what we

need to add is an H+, which has no e-. It

will share with the two e- on the top of the

Nitrogen.


Coordinate Covalent Bond

• NH 4

+

Since we are dealing with NH4+ what we

need to add is an H+, which has no e-. It

will share with the two e- on the top of the

Nitrogen.


Coordinate Covalent Bond

• NH 4

+

Since both electrons came from the nitrogen,

this is a coordinate covalent bond


VSEPR

• Valence Shell Electron Pair Repulsion

• Non bonding electrons around a central

atom will cause the molecule to bend

• Basic structures are tetrahedron, pyramidal,

bent, linear, (trigonal planar, square planar,

trigonal bipyramidal, octahedral, see saw.)


• H 2 O

VSEPR Examples


VSEPR Examples

• H 2 O

• So the molecule would look like this


• NH 3

VSEPR Examples


• NH 3

VSEPR Examples


• NH 3

VSEPR Examples


• CO 2

VSEPR Examples


• CO 2

VSEPR Examples


• CH 4

VSEPR Examples


Polarity

• Polarity is a difference in electronegativity

between two atoms that causes electrons to

not be shared equally.

• This causes one part of the molecule to

carry a slight positive charge and one side to

carry a slight negative charge


Polarity

• If the difference is

– 0.0-0.4 nonpolar covalent (shared equally)


Polarity

• If the difference is

– 0.0-0.4 nonpolar covalent (shared equally)

– 0.4-1.0 moderately polar covalent


Polarity

• If the difference is

– 0.0-0.4 nonpolar covalent (shared equally)

– 0.4-1.0 moderately polar covalent

– 1.0-2.0 polar covalent


Polarity

• If the difference is

– 0.0-0.4 nonpolar covalent (shared equally)

– 0.4-1.0 moderately polar covalent

– 1.0-2.0 polar covalent

– >2.0 ionic


Polarity

• The electrons will be around the atom with

the larger electronegativity more often

because larger electronegativity means that

the atom wants the electrons more.

• This means that the atom with the larger

electronegativity will have a partial negative

charge


Polarity

• The electrons will be around the atom with

the larger electronegativity more often

because larger electronegativity means that

the atom wants the electrons more.

• This means that the atom with the larger

electronegativity will have a partial negative

charge


Polarity

• The electrons will be around the atom with the

larger electronegativity more often because larger

electronegativity means that the atom wants the

electrons more.

• This means that the atom with the larger

electronegativity will have a partial negative

charge

• The other atom will have a partial positive charge


• H 2 O

Polarity Examples


• H 2 O

Polarity Examples


• H 2 O

Polarity Examples


• H 2 O

Polarity Examples


• CO 2

Polarity Examples


Polarity of Molecule

• If a molecule has partial positives and

partial negatives the molecule may be polar,

has a positive side and a negative side. If

you can separate all the + from the – by one

plane or one line it is polar.

• Shape of the molecule will make a

difference as to whether it is or is not polar


Polarity of Molecule


Intermolecular Attractions

• Attraction between two molecules (weak

bonds between two molecules)

• Van der Waals forces are the weakest

attractions and include dispersion forces and

dipole interactions.


Intermolecular Attractions

• Van der Waals forces are the weakest

attractions and include dispersion forces and

dipole interactions.

– Dispersion forces are caused by motion of

electrons creating very small electrical charges


Intermolecular Attractions

• Van der Waals forces are the weakest

attractions and include dispersion forces and

dipole interactions.

– Dispersion forces are caused by motion of

electrons creating very small electrical charges

– Dipole interactions are when the partial positive

of one molecule interacts with the partial

negative of another


Intermolecular Attractions

• Hydrogen bonding is a stronger bond than

van der Waals.

• It is an attractive force between hydrogen of

one molecule and the unpaired electrons on

a highly electronegative atom of another

molecule.

• Between hydrogen and either F, O, or N

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