Intermolecular forces (4) Solid structures of simple molecular lattices: Iodine and Ice

OCR A level H432 Chemistry A (from 2017)

Learning objectives

2.2.2 Intermolecular forces

(n) explanation of the solid structures of simple molecular lattices, as covalently bonded molecules attracted by intermolecular forces, e.g. I2, ice 


(o) explanation of the effect of structure and bonding on the physical properties of covalent compounds with simple molecular lattice structures including melting and boiling points, solubility and electrical conductivity. 


Solid structures of simple molecular lattices

Simple molecular lattices are very common especially in organic chemistry.

Think of sugar molecules, alkane molecules or alcohol molecules: when they freeze they exist in simple molecular lattices.

Two examples that come from inorganic chemistry are Iodine and ice.

Iodine I₂(s)

Iodine is a grey, shiny solid at room temperature and pressure.

With seven electrons in its outer electron energy level, an atom of iodine is covalently bonded to another to form a diatomic molecule I_2.

Such an arrangement means that each iodine atom possesses a full outer energy level of eight electrons.

The bond between these two atoms is a σ bond where the electron density sits along the axis joining the two iodine atoms.

The bond is formed by the end on overlap of two atomic orbitals gving a molecular orbital that is complete with two electrons.

The picture to the left  is an attempt at a representation of this covalent bond.

A weak induced dipole—induced dipole force holds the iodine molecules in the molecular lattice since both atoms have the same electronegativity and so do not possess permanent bond polarity.

The weak intermolecular force is called by its generic name in the diagram above as a van der Waals force.

The final picture attempts to show the arrangement of iodine molecules in the crystal lattice as they are close packed together.

Ice H₂O(s)

Ice is a hard, colourless solid at temperatures below 0^oC or +273K.

It is quite brittle and breaks easily to produce sharp shards that can cut through skin.

Ice is composed of two elements of significantly different electronegativity:  oxygen (3.44) and hydrogen (2.20. 

So ice molecules are polar covalent molecules.

The covalent σ bonds that exist between oxygen and hydrogen are polar with the oxygen atom pulling the bonding electron pairs towards itself. 

Because of the covalent bond polarity molecules hydrogen bonds are found between molecules of water in ice.

There are two hydrogen bonds per water molecules that mean the structure of ice is very similar to that of diamond.

Compare the two structures here:

                                                         Diamond   

                                                                Ice

Two hydrogen bonds per molecule are possible because the molecule of water has two lone pairs on the oxygen atom.

The intermolecular forces in ice are permanent dipole forces and the much weaker induced dipole forces that we find exist between all molecules at some level. 

Substance	Structure	Bonding	Solubility in water	Mp and bp /oC	Electrical conductivity
Iodine	Diatomic molecular Lattice	Covalent	Partially soluble	387K 457K	no
Ice	Simple molecular lattice	Polar covalent	Fully miscible	273K 373K	yes

Intermolecular Forces (3) Hydrogen Bonding

Here we go again on Chemical Bonding (8) but this time its Hydrogen Bonding.

This type of chemical bonding is just as ubiquitous as van der Waals forces or permanent dipole bonding in the world of molecules.

But hydrogen bonding is much stronger than the other two intermolecular bonds.

H bonding can be as strong as 40-50 kJ per mole.

Evidence

Let’s look at some of the evidence for this type of bonding starting with a look at the boiling points of the hydrides of Groups 4, 5, 6 and 7 in the Periodic Table.

The dotted lines show the expected trend whereas the bold lines indicate the actual values. 

(Incidentally, there are no continuous variables here, these bold lines are just for purposes of clarity since Period is a categoric variable not a continuous variable like boiling point)

  

What can we see from this slide?

The boiling points of water (H₂O) ammonia (NH₃) and hydrogen fluoride (HF) seem to be out of synch with the rest of the groups’ hydrides unlike methane (CH₄) which fits the group trends.

Here is more evidence from how the density of water changes with temperature.

The first surprise looking at this plot is that the density of water at temperatures just above its melting point is not constant.

This change in density suggests that water’s structure as a liquid is changing and in fact liquid water has a structure, the molecules are not just randomly arranged or jumbled together.

And for all of us telling you that the density of water is 1kg/m³ or 1 g/cm³, it isn’t at all just  look at the variation in those density figures.

Second is the fact that there is a maximum density at 4^oC.

What’s all this telling us about water?

The obvious thing is that at the bottom of your garden pond in winter when the air temperature is -1^oC the water at the bottom of the pond is warmer at 4^oC and your fish are not frozen but still alive!!

And then we can look at evidence from DNA replication how is it that the two strands of the DNA double helix can unzip and zip up so easily?

How come carboxylic acids are often found to have double their actual molar mass?

Or how do the secondary protein structures of the alpha helix and beta pleated sheet form?

Or 4–nitrophenol has a different molar mass and melting and boiling points to its analogy isomer 2–nitrophenol?

All these phenomena (and there are others here and here) are evidence for hydrogen bonding.

Let me explain:

Boiling points of ammonia, water and hydrogen fluoride

Take the seemingly anomalous boiling points of ammonia , water and hydrogen fluoride.

These boiling points cannot just be due to the permanent dipole forces between these polar molecules nor just the result of the temporary dipoles set up as the electrons shift from place to place in the molecule.

Both the permanent dipoles forces and the temporary dipole or London forces are too weak to account for the sudden hike in boiling point.

Another force must be involved.

And this force has to be stronger than either London forces or permanent dipole forces for the hike in boiling point to be so great.

This intermolecular force we call a hydrogen bond first recognised by– yes you guessed it: Linus Pauling—in that famous book of his “The Chemical Bond”.

Here is what Pauling said of the hydrogen bond in the late 1930’s:

“It was recognised some decade ago that under certain conditions an atom of hydrogen is attracted by rather strong forces to two atoms instead of only one so that it may be considered to be acting as a bond between them.  This is called the hydrogen bond.”

Initially it was thought that these atoms attracted to the hydrogen atom were the highly electronegative atoms of fluorine, nitrogen or oxygen. 

Therefore, in ammonia a hydrogen bond could form between two ammonia molecules and the strength of this bond account for ammonia’s higher boiling point.

You can see that in the diagram the hydrogen bond between the two molecules consists of the hydrogen from one molecule attracted to the nitrogen of the other incorporating the lone pair.

So the bond has a degree of covalency. 

The bond angle of this hydrogen bond contrary to what you may have been taught is not 180^o but 160^o.

In water the hydrogen bond angle is 180^o

The diagram below shows the bonding in ice.

Water molecules in ice are hydrogen bonded twice (look for two dotted lines in the diagram )which leads to an open structure and therefore a lower density than that of liquid water.

As ice melts its hydrogen bonded structure collapses and molecules randomly pack closer together.

The result is that  water reaches its maximum density at 4^oC

And hydrogen fluoride has this structure:

DNA

DNA has two strands that are bonded together via the base pairs.

These base pairs bond using hydrogen bonds as in this diagram:

Cytosine and guanine fit together through the formation of hydrogen bonds as do adenine and thyamine.

Carboxylic acids

The two molecules dimerise by forming hydrogen bonds between the carboxyl groups as in the diagram. 

The α-helix  and the β-pleated sheet

These two protein secondary structures are supported by the formation of hydrogen bonds between the amino acid residue side chains as in these two diagrams:

The dotted lines in the diagrams are the positions of the hydrogen bonds.

Nitrophenol

Here we see in 2-nitrophenol the intramolecular hydrogen bond creating a hexagonal ring structure.

Name	Structure	Melting point oC	Boiling point oC
2-nitrophenol		44	215
4-nitrophenol		112	279

The effect of this internal hydrogen bonding is to reduce the melting and boiling points of the 2- isomer relative to the 4- isomer because there is no intermolecular hydrogen bonding in the 2- isomer.

You can see in 2-nitrophenol that the H bond is not linear but 120^o

In the protein structures, the H—bond is between the N—H group of one peptide and the C=O group of another i.e. N—H•••O=C where the three dots represent the H—bond.

So the two electronegative atoms do not have to be the same for an H—bond to form. 

In this reference here you can read of some of the latest thinking on H—bonds.

This is the reference to Linus Pauling’s “The Chemical Bond” extract.

Here is a recent discussion about the definition of the hydrogen bond from the scientific periodical Nature.