Covalent Bonding Honors Chemistry – Semester 1

Ionic vs Covalent Bonding Ionic bonding = exchange electrons between metal / non-metal

Covalent bonding = sharing electrons between non-metals

Cl

Cl-

Simple Covalent Molecules Simplest example = diatomic molecules

Repulsion between nuclei and between electron clouds

Attraction between nuclei and electron clouds

Covalent bond forms when forces balance

Shared electrons located within molecular orbitals

Energy & Stability Most individual atoms have low stability

atoms form compounds!

Unbonded atoms = high potential energy

Bonded atoms = low potential energy

Bond Length Bond length = marks point of lowest potential energy

Measured in Ångstrom (Å) or picometer (pm)

1Å = 100 pm = 1 x 10-10 meters

Nuclei vibrate, distance changes constantly

average distance = bond length

Bond Energy = energy required to break bonds in 1 mol of a

chemical compound

Energy released when atoms bond

Example:

energy released when

I mol of H2 bonds

= -436 kJ/mol

436 kJ/mol must be

supplied to break bonds

Bond Energy High bond energy = strong bonds

Strong bonds = short bond lengths

Bond energy predicts reactivity of compound

low bond energy reacts more easily

Review Questions Why is H2 more stable than individual atoms of H?

H2 has lower potential energy than H

In what type of orbital are shared electrons in a bond

located?

Molecular orbital

What happens to the potential energy of two atoms as

they approach each other to form a covalent bond?

Potential energy decreases

What happens to potential energy if the atoms get too

close to each other?

Potential energy increases

Review Questions What do you call the the distance between two atoms in

a covalent bond at which potential energy is lowest? Bond length

Which of these is likely to have covalent bonds? Recall that covalent bonds tend to form between non-metals.

H2

MgO

O2

CO2

H2O

K2SO4

CH4 (methane)

C6H12O6 (sugar)

Yes

No

Yes

Yes

Yes

No

Yes

Yes

Electronegativity &

Covalent Bonding

Helps predict type of bond between elements

Generally:

Metals less electronegative than nonmetals

Smaller atoms have greater electronegativity

Nonpolar vs Polar

Covalent Bonds

Shared electrons closer to atom with higher electronegativity!

Non-polar covalent: electron pair(s)

shared equally between nuclei of

similar electronegativity

Polar covalent: electron pair(s)

shared unequally between nuclei with

different electronegativity

Cl Cl

Cl H

Example: Polar bond HF H, F = nonmetals Covalent bond

Electronegativity values:

F = 4.0; H = 2.1

Most shared electrons with F

Partial negative charge (δ-) in F, partial positive charge in

H (δ+)

Creates dipole

Dipole = a molecule or part of a molecule that

contains both positive and negative regions

Electronegativity, Bond

Polarity & Bond Strength

The greater the difference in electronegativity:

the greater the polarity of the bond

the stronger the bonds

Electronegativity & Bond Type

This scale is only a guide!!

Some compounds may deviate!!

Practice:

Predicting Bond Types Na – F

Electronegativity difference: 3.1 ionic

C - Cl

0.5 polar covalent

Ca – O

2.5 ionic

Al - Cl

1.5 polar covalent