Doc Brown's Chemistry Revision

Extra Notes on chemical bonding for advanced A level chemistry students

 Extra notes on Electronegativity, Bond Polarity and Influence on Type of Chemical Bonding

 (also notes on dipoles)

CHEMICAL BONDING INDEX

Introduction

Electronegativity is defined as the power of an atom to attract the pair of electrons in a covalent bond situation. The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical. This produces a polar covalent bond, and may cause a molecule to have a permanent dipole which is indicated by partial charges (δ+ and δ) to show that a bond is polar. You should be able to recognise when a molecule is likely to be polar and explain why some molecules with polar bonds do not have a permanent dipole.


1. Electronegativity

Electronegativity is the power of an atom to attract electron charge from another atom it is covalently bonded to. Some Pauling values of electronegativity are quoted below.

element Na Mg Al Mn Fe H Si P C S I Br Cl N O F
electronegativity 0.9 1.2 1.5 1.5 1.8 2.1 1.8 2.1 2.5 2.5 2.5 2.8 3.0 3.0 3.5 4.0

Generally speaking electronegativity increases from left to right across a period of the periodic table and decreases down a group of the periodic table.


2. Bond polarity and type of chemical bonding

Three situations are considered

(a) Non-polar covalent bond

 If the two atoms of a covalent bond have the same or very similar electronegativity, the bond is referred to as non-polar.

Examples of non-polar bonds (electronegativity difference): H-H (0), C-C (0), C-H (0.4)

(b) Polar covalent bond

If two atoms of a covalent bond have an appreciable difference in electronegativity, the result is a polar bond AND the molecule usually has a permanent dipole due to a permanent partial charge separation. This asymmetry in the electron cloud distribution is shown a delta + and a delta - (δ+ and δ).

Examples of polar bonds (electronegativity difference): H-Cl (0.9), O-H (1.3), N-H (0.9), C-Cl (0.5)

The polar bonds in polar molecules with a permanent dipole can be represented as:

Hδ+-Clδ, δO-Hδ+, δN-Hδ+, Cδ+-Cl δ

a picture of HCl , water

Note that in some circumstances a molecule with polar bonds may NOT have a permanent dipole because the individual dipoles cancel each other out because of the symmetry of the molecule.

e.g. in tetrahedral shaped tetrachloromethane CCl4, the Cδ4+ is cancelled out by the 4 symmetrically arranged Clδ– atoms.

but in chloromethane (c) doc b there is a permanent dipole due to the C-Cl bond (Cδ+-Clδ–).

Carbon dioxide O=C=O (δO=Cδ2+=Oδ) has two polar bonds, but the molecule is linear and the two dipoles cancel each other out.

Simple laboratory experiment to detect permanent polarity in a molecule

By running jets of various liquids from burettes past an electrostatically charged plastic rod you find highly polar molecules like water are quite dramatically deflected due to the permanent dipole but non-polar molecules like hexane show no deflection at all.

The plastic rod is charged by rubbing with a wool or line cloth. A 'bar' of poly(ethene) is quite good for this deflection experiment.

Dipole moment

Its unlikely you need to know about this unit, but some advanced pre–university courses do mention dipole moments.

Dipole moment (μ) is the measure of net molecular polarity, and is the magnitude of the charge (Q) at either end of the molecular dipole multiplied by the distance between the charges (d).

μ = Q x d

Dipole moments are quoted in D, debye units, 1D = 3.34 x 10–30Cm

This is a non–SI unit. C = coulomb, m = metre.

Dipole moments informs about the charge separation in a molecule. The greater the difference in electronegativities of the bonded atoms, the larger the dipole moment. Also, the larger the charge separation, the greater the dipole moment.

(c) Ionic bond

If two atoms constituting a bond have a big difference in electronegativity then an ionic bond may be formed.

In other words, rather than sharing electrons to form a covalent bond, electron transfer is completed giving rise to positive and negative ions whose electrostatic attraction constitutes the ionic bond.

Examples of ionic bonds (electronegativity difference): Na+Cl (2.1), Mg2+(F)2 (2.8)


CHEMICAL BONDING INDEX


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