Formation of Covalent Bonds

Lessons

Senior 4 · Chemistry

Formation of Covalent Bonds

5 lessons

Lesson 5 of 5

Polar Bonds and Intermolecular Forces

25 min

Water should be a gas

Hydrogen sulfide, $\mathrm{H_2S}$ , the gas that smells of bad eggs, has molecules almost twice as heavy as water's, and it boils at −60 °C. By that trend, water would boil far below zero and Lake Kivu would be a cloud of vapour. Instead water boils at 100 °C, insects can walk on its surface, and ice floats on it. None of this comes from the strong covalent bonds inside each molecule. It comes from weaker forces between molecules, which in turn come from how unequally the electrons in a bond are shared.

Definition

Intramolecular forces: the forces within a molecule, the covalent bonds that hold its atoms together. Intermolecular forces: the much weaker attractions between separate molecules. Melting and boiling a molecular substance overcome intermolecular forces; the covalent bonds stay intact.

Electronegativity: the ability of an atom to attract the shared pair of electrons in a covalent bond towards itself. On the Pauling scale it runs from about 0.8 (caesium) to 4.0 (fluorine).

Electronegativity and polar bonds

Table 5.1: Pauling electronegativities

H	C	N	O	F	Na	Mg	P	S	Cl	Br	I
2.2	2.6	3.0	3.4	4.0	0.9	1.3	2.2	2.6	3.2	3.0	2.7

Electronegativity increases across a period (greater nuclear charge, same shielding, smaller atom) and decreases down a group (more shells, more shielding). Fluorine, oxygen and nitrogen are the most electronegative elements.

When two identical atoms share electrons (H–H, Cl–Cl) the pair is shared equally: a non-polar bond. When the atoms differ, the more electronegative atom pulls the pair towards itself and gains a small negative charge, δ−, leaving the other atom δ+. This is a polar covalent bond, with a bond dipole: $\overset{\delta+}{\mathrm{H}}{-}\overset{\delta-}{\mathrm{Cl}}$ . The bigger the difference in electronegativity, the more polar the bond; when the difference is very large (roughly 1.8 or more), electron transfer is nearly complete and the bonding is mainly ionic. Covalent and ionic bonding are two ends of one scale.

Polar bonds, but is the molecule polar?

A molecule is polar (has a permanent dipole) if its bond dipoles do not cancel. That depends on its shape (Lesson 4). In a symmetrical molecule the dipoles cancel, even if every bond is polar.

You can see the difference: hold a rubbed plastic rod near a thin stream of water from a burette and the stream bends towards the rod. A stream of hexane, a non-polar liquid, is hardly affected.

Worked example 1: polar or non-polar?

Decide whether each molecule is polar: (a) $\mathrm{BF_3}$ ; (b) $\mathrm{CHCl_3}$ ; (c) $\mathrm{SO_2}$ .

Solution.

(a) B–F bonds are very polar, but $\mathrm{BF_3}$ is trigonal planar: three equal dipoles at 120° cancel. Non-polar.
(b) $\mathrm{CHCl_3}$ is tetrahedral, but one position holds H instead of Cl. The C–H dipole is small and points the other way, so the C–Cl dipoles are not cancelled. Polar.
(c) $\mathrm{SO_2}$ is bent (two bonding regions and one lone pair on S), so the two S–O dipoles add. Polar.

Three kinds of intermolecular force

1. London (dispersion) forces act between all molecules, polar or not. Electrons are always moving, so at any instant the electron cloud of a molecule may be slightly uneven: an instantaneous dipole. It distorts the electron cloud of a neighbour, inducing a dipole there, and the two attract. London forces are stronger when a molecule has more electrons (the halogens: $\mathrm{F_2}$ and $\mathrm{Cl_2}$ are gases, $\mathrm{Br_2}$ a liquid, $\mathrm{I_2}$ a solid) and when molecules have more surface contact: straight-chain pentane boils at 36 °C, its compact isomer 2,2-dimethylpropane at 10 °C.

2. Permanent dipole–dipole forces act between polar molecules: the δ+ end of one attracts the δ− end of another, in addition to London forces. Propanone, $\mathrm{CH_3COCH_3}$ ( $M_{r} = 58$ , polar C=O), boils at 56 °C; butane ( $M_{r} = 58$ , non-polar) boils at −1 °C.

3. Hydrogen bonds are the strongest intermolecular force. A hydrogen bond forms between a hydrogen atom bonded to N, O or F in one molecule and a lone pair on N, O or F in another. These three atoms are very electronegative and small, so the H atom is strongly δ+ (almost a bare proton) and can come very close to the lone pair.

Evidence for hydrogen bonding: boiling points of hydrides

In Group 14 the boiling points rise smoothly down the group: the molecules are non-polar, and London forces grow with the number of electrons. In Groups 15, 16 and 17 the trend is the same from Period 3 to Period 5, but $\mathrm{NH_3}$ , $\mathrm{H_2O}$ and HF are far above it. Only these three have H bonded to N, O or F, so only they form hydrogen bonds. Water is highest of all because each molecule has two δ+ hydrogens and two lone pairs, so it can form up to four hydrogen bonds.

Hydrogen bonding also explains why ice floats: in ice each molecule is hydrogen-bonded to four others in an open, tetrahedral lattice with a lot of empty space, so ice is less dense than liquid water. It explains why ethanol and sugar dissolve in water (they form hydrogen bonds with it), the surface tension of water, and the pairing of bases that holds the two strands of DNA together.

Worked example 2: comparing boiling points

Ethanol, $\mathrm{CH_3CH_2OH}$ , and methoxymethane, $\mathrm{CH_3OCH_3}$ , have the same formula, $\mathrm{C_2H_6O}$ ( $M_{r} = 46$ ). Ethanol boils at 78 °C and methoxymethane at −24 °C. Explain the difference.

Solution. Both have the same number of electrons, so their London forces are similar, and both are polar. But ethanol has an O–H group: its δ+ hydrogen forms hydrogen bonds with the lone pairs on the oxygen atoms of neighbouring molecules. In methoxymethane every hydrogen is bonded to carbon, which is not electronegative enough, so it cannot form hydrogen bonds between its own molecules. Much more energy is needed to separate ethanol molecules.

Answer: hydrogen bonding between ethanol molecules, absent in methoxymethane, raises ethanol's boiling point by about 100 °C.

Warning

Common mistakes. (1) Saying a substance boils when its covalent bonds break: boiling separates molecules; the bonds inside them stay. (2) Calling the O–H bond in water a hydrogen bond: a hydrogen bond is between molecules. (3) Saying any molecule with H forms hydrogen bonds: the H must be bonded to N, O or F. (4) Ignoring London forces: they act in every molecular substance and can be larger than dipole–dipole forces in big molecules.

Safety. In the charged-rod demonstration, hexane is highly flammable: no flames anywhere in the room, small volumes only, and work in a fume cupboard.

Tip

To compare boiling points, ask three questions in order: (1) Does it form hydrogen bonds (H on N, O or F)? (2) Is it polar (permanent dipole)? (3) How many electrons and how much surface contact (London forces)? The answers usually decide the order.

The Organic Chemistry Tutor: Intermolecular Forces - Hydrogen Bonding, Dipole Dipole Interactions - Boiling Point and Solubility (11 min). Watch how the type of force is identified for each molecule and then used to rank boiling points.

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Practice

Define electronegativity and explain its trend across Period 2 and down Group 17.
Use Table 5.1 to put these bonds in order of increasing polarity: C–H, O–H, N–H, H–F, C–Cl.
Predict whether each molecule is polar: $\mathrm{CS_2}$ , $\mathrm{PCl_3}$ , $\mathrm{CH_2Cl_2}$ , $\mathrm{SF_6}$ . Give reasons.
Explain why the boiling points of the halogens increase from fluorine to iodine.
Draw two molecules of ammonia and show a hydrogen bond between them, with partial charges and the lone pair.
Explain why HF (20 °C) boils at a higher temperature than HCl (−85 °C), but HI (−35 °C) boils higher than HCl.
Explain why ice is less dense than liquid water, and why this matters for fish in a lake.

Show answersHide answers

The ability of an atom to attract the bonding pair in a covalent bond. It increases across Period 2 (more protons, same shielding, smaller atoms) and decreases down Group 17 (more shells, more shielding, bonding pair further from the nucleus).
Differences: C–H 0.4, C–Cl 0.6, N–H 0.8, O–H 1.2, H–F 1.8. Order: C–H < C–Cl < N–H < O–H < H–F.
$\mathrm{CS_2}$ linear, dipoles cancel: non-polar. $\mathrm{PCl_3}$ trigonal pyramidal (one lone pair): polar. $\mathrm{CH_2Cl_2}$ tetrahedral but with two different kinds of bond, dipoles do not cancel: polar. $\mathrm{SF_6}$ octahedral, dipoles cancel: non-polar.
The molecules are non-polar, so only London forces act. From $\mathrm{F_2}$ to $\mathrm{I_2}$ the number of electrons increases, so the instantaneous and induced dipoles are larger and the London forces stronger.
N–H···N: the δ+ H of one molecule lined up with the lone pair on the δ− N of the other, drawn as a dotted line, close to 180°.
HF forms hydrogen bonds (H bonded to F), much stronger than the dipole–dipole and London forces in HCl. From HCl to HI there is no hydrogen bonding; the number of electrons rises, so London forces increase and outweigh the falling polarity.
In ice each molecule is hydrogen-bonded to four others in an open tetrahedral lattice with empty space; melting breaks some hydrogen bonds and the molecules pack closer. Ice floats and forms an insulating layer on top, so the water beneath stays liquid and fish survive.

Quick check

Which of these molecules is polar?

Key points

Electronegativity rises across a period and falls down a group; unequal sharing makes a bond polar (δ+ and δ−).
A molecule is polar only if its bond dipoles do not cancel: $\mathrm{H_2O}$ , $\mathrm{NH_3}$ , $\mathrm{CHCl_3}$ are polar; $\mathrm{CO_2}$ , $\mathrm{CCl_4}$ , $\mathrm{BF_3}$ are not.
Intermolecular forces, weakest to strongest: London forces (all molecules; more electrons, more contact), permanent dipole–dipole, hydrogen bonds (H on N, O or F to a lone pair on N, O or F).
Hydrogen bonding explains the high boiling points of $\mathrm{NH_3}$ , $\mathrm{H_2O}$ and HF, the low density of ice and the solubility of ethanol in water.

1Water should be a gas

2Electronegativity and polar bonds

3Polar bonds, but is the molecule polar?

4Three kinds of intermolecular force

5Evidence for hydrogen bonding: boiling points of hydrides

Water should be a gas

Electronegativity and polar bonds

Polar bonds, but is the molecule polar?

Three kinds of intermolecular force

Evidence for hydrogen bonding: boiling points of hydrides