2.2.2 (d,e,f) Covalent Bonding

Syllabus

(d) covalent bond as the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

(e) construction of ‘dot-and-cross’ diagrams of molecules and ions to describe:

(i) single covalent bonding

(ii) multiple covalent bonding

(iii) dative covalent (coordinate) bonding

{‘Dot-and-cross’ diagrams of up to six electron pairs (including lone pairs) surrounding a central atom.}

(f) use of the term average bond enthalpy as a measurement of covalent bond strength

{Learners should appreciate that the larger the value of the average bond enthalpy, the stronger the covalent bond.}

{Definition and calculations not required.}

{Average bond enthalpies and related calculations are covered in detail in 3.2.1 f.}

What does this mean?

Definition

The GCSE definition of a covalent bond is still acceptable.

"an electrostatic attraction between two positively charged nuclei and a shared pair of electrons"

The attraction is strong but there is also repulsion between the two nuclei.

If the atoms get closer together then the repulsion will eventually be bigger than the attraction and the two atoms will move apart.

If they move apart too far the attraction outweighs the repulsion.

So the bond oscillates about an average bond-length - but this isn't entirely fixed.

Generally, the stronger the bond enthalpy the shorter it will be.

Dot-cross diagrams

Single bonds

This is also similar to what you learned at GCSE.

Firstly, there will be a covalent bond if (usually) two non-metal elements are involved.

To create a dot-cross diagram you need to:

Work out the electronic structure of the atoms involved.
Draw the outer shell only, pairing electrons where necessary.
Overlap the outer shells so that unpaired electrons can pair.
No electrons should be left unpaired

eg Hydrogen

Each Hydrogen atom needs to pair up its 1 electron.

So a Hydrogen molecule contains one single bond as shown.

eg. Water

Each Hydrogen atom needs to pair up its 1 electron.

Each Oxygen has 2 electrons in need of pairing.

So a water molecule contains two single bonds as shown.

The un-shared pairs are called lone pairs.

They'll be important when we come to shapes of molecules

eg. Chlorine

Each Chlorine atom has 7 electrons but only needs to pair up 1.

So a Chlorine molecule contains one single bond as shown below.

Double and Treble bonds.

It's possible for more than one pair of electrons to be shared.

Oxygen

Oxygen atoms have two unpaired electrons each.

So they share two pairs.

Making a double bond.

This is not as strong as two single bonds but is stronger than one single bond.

Nitrogen

Each Nitrogen atom has three unpaired electrons.

So they share three pairs.

In other words a treble bond forms.

Dative bonds

A dative covalent is:

"the electrostatic attraction between two nuclei and a shared pair of electrons which were both donated by one of the atoms"

An examiner will want to see the dative bond shown by a shared pair in which there are two dots (or two crosses).

This diagram shows an Ammonia molecule datively bonding to a Hydrogen ion.

Notice that the H⁺ ion has no electrons.

And that the bond that forms between them is formed from two electrons that were both in a lone-pair on the Nitrogen atom.

An examiner wants to see one of the N-H bonds with two dots to show both electrons came from the N atom.

When drawing bonds as lines, a dative bond has an arrow from the N to the H to show that it's the N atom donating both electrons

Ions

Hydroxide

In this ion we can see a normal covalent bond between the O and H atoms, represented by one dot from the O and one from the H.

The O would still have only 7 electrons unless it gained one from elsewhere.

This diagram shows it as a red cross.

Carbonate

Here we can see the right-hand Oxygen atom donating two (dot-electrons to pair with two x-electrons from the carbon and forming a double bond.

The left-hand and bottom Oxygen atoms donate one dot each to form a single bond with the Carbon.

This leaves these two Oxygen atoms short of an electron each - this lack is made right by gaining two extra electrons (empty dots) and this is why it is CO₃^2- rather than just CO₃.

In reality, there's no reason to form a double bond with any particular Oxygen atom since they're all the same.

So, the double bond moves from one to the next & back.

Powerpoint

Google Presentation

Video

Exam-Style Questions

1. (a) An Ammonium ion, made by the reaction between Ammonia molecule and a

Hydrogen ion, can be represented as shown in the diagram above.

(i) Name the type of bond represented in the diagram N-H

………………………………………………………………………………………….............................................…….

(ii) Name the type of bond represented in the diagram by N→H

………………………………………………………………………………………………………………………………………………(iii) In terms of electrons, explain why an arrow is used to represent this NH bond

……………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………..............................…………

2. (a) What is a covalent bond?

…………………………………...................................................................……………………………………………………

(1)

(b) Draw dot-and cross diagrams (outer shells only) to show the bonding in Carbon Disulphide (CS2) and Hydrogen Sulphide (H2S).

Bonding in CS2 Bonding in H2S

(2)

(b) Showing the outer electrons only, draw a dot-and-cross diagram to indicate the bonding in Calcium Oxide.

(2)

(c) (i) Give the type of bonding present in BeCl2

…………………………………...................................................................……………………………………………………

(1)

(ii) Give the type of bonding present in BaCl2

…………………………………...................................................................……………………………………………………

(1)

(iii) Explain why the type of bonding is different in these two compounds.

…………………………………...................................................................……………………………………………………

(1)

Answers

1.

(a) (i) Covalent (1)

(ii) Co-ordinate (1) (or dative)

(iii) Both / two / pair electrons come from nitrogen (1)

(a) Shared pair of electrons 1

(b) circles need not be shown

all outer electrons must be shown (1) 2

must be S, C and H, S

if omission of non-bonding electrons occurs twice, penalise first time only

non-bonding electrons needn’t be shown as a pair

(c) transfer of electron(s) (1)

correct charges on the ions (1)

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