2.2.2 (i,j) Electronegativity and Bond Polarity

Syllabus

(i) electronegativity as the ability of an atom to attract the bonding electrons in a covalent bond; interpretation of Pauling electronegativity values

{Learners should be aware that electronegativity increases towards F in the periodic table.}

(j) explanation of:

(i) polar bond and permanent dipoles within molecules containing covalently-bonded atoms with different electronegativities

(ii) a polar molecule and overall dipole in terms of permanent dipole(s) and molecular shape

{A polar molecule requires polar bonds with dipoles that do not cancel due to their direction. E.g. H₂O and CO₂ both have polar bonds but only H₂O has an overall dipole.}

What does this mean?

Polar and non-polar bonds

In a Hydrogen molecule the shared pair of electrons is equally attracted to both atoms.

This is because both nuclei contain the same number of protons (one).

The electrons won't always be exactly halfway between the two atoms but this will be the average position.

Any pair of identical atoms covalently bonded would be like this eg. O₂, Cl₂, N₂, I₂, Br₂, F_{2 etc}

All these would contain non-polar bonds and would be non-polar molecules.

But what about a l-F molecule?

Iodine atoms contain more protons to attract the shared pair.

But Fluorine atoms are much smaller so the shared pair would be much closer to its nucleus.

And it's the proximity to the nucleus that is more important.

Added to which, A level examiners like you to mention shielding - which we can think of as inner shells screening the valence (bonding) shell from the full force of nuclear attraction.

And smaller atoms like Fluorine would also have less shielding.

All of which means that the Fluorine atom would have a bigger attraction for the bond pair than the Iodine atom would.

We say it has a larger electronegativity.

We now have a polar-bond.

A polar bond means an unequal share of electron density.

In this case the Fluorine gets the bigger share, so it becomes partly negative (δ-)

And the Iodine gets the smaller share, so it becomes partly positive (δ+)

If the difference in electronegativity was big enough then one atom would completely gain the bond pair to become a negative anion, leaving the other a positive cation.

You do not have to learn any figures.

But it is handy to know that any difference below around 0.5 isn't considered polar and anything above 1.7ish is considered ionic.

If you require electronegativity numbers, the examiner will have to provide them.

However, you should be aware of the trends across each period - increasing as the number of protons increases, the atomic radius decreases and shielding remains constant.

You should also know that electronegativity increases up each group as the atomic radius and shielding decrease.

And, you're going to need to be able to explain these trends.

Non-polar molecules with polar bonds.

There is a significant difference in electronegativity between Carbon and Oxygen atoms.

So C=O bonds are polar.

But CO₂ is linear so the two dipoles are pointing in exactly opposite directions and cancel.

This makes the molecule as a whole non-polar.

So shapes of molecules are as important to the polarity of the molecules as dipoles are.

If water was linear its dipole would also cancel and it would also be non-polar.

But water is V-shaped so the dipoles don't cancel.

There is a δ+ end and a δ- end.

C-Cl bonds are very polar. But CCl₄ is tetrahedral so again the dipoles cancel.

Making CCl₄ a non-polar molecule.

But as soon as we replace one Cl with an H atom it is no longer symmetrical.

The dipoles don't cancel any more.

So HCCl₃ is a polar molecules.

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Exam-Style Questions

1(a) Define the term electronegativity

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

(2)

(b) A bond between Nitrogen and Hydrogen can be represented as

(i) In this representation, what is the meaning of the symbol δ⁺?

………………………………………………………………………………………………………………………………………………

(ii) From this bond representation, what can be deduced about the electronegativity of hydrogen relative to that of Nitrogen?

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

(2)

2. (a) Define the term electronegativity and explain why the electronegativity values of the Group II elements Be–Ba decrease down the group. (4)

3. The table below contains electronegativity values for the Period 3 elements, except chlorine.

(a) Define the term electronegativity .......................................................................................................................................................................

(2)

(b) Explain why electronegativity increases across Period 3 ....................................................................................................................................................................... .......................................................................................................................................................................

(2)

(c) Predict values for the electronegativities of chlorine and of lithium.

Electronegativity of Chlorine ………………………………...........................................................................................………………………………..

Electronegativity of Lithium…………..........................................................................................……………………………………………………

(2)

(d) State and explain the trend in electronegativity down Group II

Trend ………………………………………………………………………………………….

Explanation………………………….....................................................................................................…………………………………….....……….…………………………………………..............................................................................…………………

(3)

(Total 9 marks )

Answers

1(a) Power (or ability) of an element / atom to attract electron pair/electrons/an electron/electron density (1)

in a covalent bond (1)

Allow attract from, withdraw in, do not allow remove from, withdraw from.

(b)(i) Electron deficient (1)

Or small, slight, partial positive charge

(ii) H < N (1)

(2) Tendency or strength or ability or power of an atom/element/nucleus to 1

attract/withdraw electrons / e– density / bonding pair / shared pair

In a covalent bond 1

(tied to M1 – unless silly slip in M1)

(If molecule/ion then = CE = 0) (NOT electron (singular) for M1)

Increase in size or number of shells or increased shielding or bonding 1

electrons further from nucleus

[NOT ‘increase in number of electrons’]

Decreased attraction for (bonding) electrons 1

(If ‘ion’ here, lose M3 and M4) (NOT ‘attraction of covalent bond’)

(Ignore reference to proton number or effective nuclear charge)

3. (a) power of an atom to attract electrons (1)

in a covalent bond (1)

(b) number of protons increases (1)

electrons in same shell (or similar shielding) (1)

(c) Electronegativity of Chlorine allow 2.8 – 3.0 (1)

Electronegativity of Lithium allow 0.91 –1.5

(d) Trend : decreases (1)

Explanation : outer electrons further from nucleus (1)

more shielding from the nucleus (1)

(or more shells or increasing radius)

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