4.2 Covalent Bonding

Syllabus

What does this mean?

How do covalent bonds work?

Metal atoms need to lose electrons and non-metals to gain; this creates ions and leads to Ionic bonds.

But Non-metal atoms also bond to non-metal atoms.

They can’t both gain electrons. And, if they did, they would both the two negative ions formed would repel.

So another type of bond is required in which electrons are shared.

It looks as though the shared electrons should repel each other (& they do).

But the attractions from the positive nuclei are strong enough to outweigh this repulsion, pulling the atoms together.

But not too close or the nuclei will repel each other too.

The atoms are always moving.

But their average separation (the “bond-length”) represents the lowest energy state.

VIDEO

Dot-Cross Diagrams.

These will be familiar from GCSE for many students.

(Almost) all atoms react to fill their outside shells.

In practice this means that unpaired electrons will pair.

We draw only the outside shells of atoms (valence shells) because only these are involved in bonding.

#1 F2

So Fluorine (9F) has an electron configuration of (2.7)

With each F atom having 1 bond-pair and three lone-pairs (non-bonding)

So, each atom needs to pair up one unpaired electron, forming 1 bond.

#2 H2O

Hydrogen atoms have 1 valence electron. (1H)

Oxygen atoms have 6. (8O = 2.6)

So paired up we get a molecule shown right

With the O atom having 2 bond-pairs and 2 lone-pairs (non-bonding).

Questions.

Draw Dot-Cross Diagrams of the following molecules.

Cl2

CH4

NH3

HCl

C2H6

H2O2

Multiple bonds.

We know that Oxygen exists as O2.

And that each Oxygen atom has an electron configuration 2.6

If we joined these with single bonds the O atoms would have unfilled outer shells containing only 7 electrons.

So, rather than pairing one electron each they pair both.

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Note:

1. Double bonds are stronger than (but not twice as strong as) the equivalent single bond.

2. Stronger bonds are shorter than weaker bonds due to the added attraction.

Questions.

Draw Dot-Cross Diagrams of the following molecules.

CO2

N2

C2H4

C2H2



Electronegativity.

· Learn the definition

“Electronegativity = the relative ability of atoms to attract the bond-pair of electrons in a covalent bond.”

· Learn the trends in electronegativity shown below:

Why?

· The bond pair is attracted by the nuclei.

· Down a group the bond pair is further from the nuclei and the attraction is weaker = lower electronegativity.

· Across a period the atoms decrease in size, the nuclei are closer to the bond-pair and the attraction is greater = higher electronegativity.

· The distance between the nuclei and the bond pair is much more important than the number of protons in the nucleus or the level of shielding.

Bond Polarity

When one atom attracts the bond pair more strongly than another it gets an “unequal share” of the bond pair – the electrons end up closer to the more electronegative atom.

H has an electronegativity of 2.1, and F of 4.0.

So the bond pair will be pulled towards the F atom in a molecule of HF.

This means that there is more electron density on the F side –making it partially negative (∂-).

And there is less electron density on the H side, making it partially positive (∂+).

This is a polar bond.

The bigger the difference in electronegativity the bigger the Dipole Moment.

In this case the difference is 1.9 – making this bond so polar that it’s considered to be Ionic.

The IB uses a difference of 1.8 as the boundary.

VIDEOS

Questions.

Label the following covalent bonds with ∂+ and ∂- to show the dipole. Indicate any bond that would be considered to be Ionic.

1. C-H

2. N-H

3. B-Cl

4. B-F

5. Mg-F

6. H-S