1.4 FORCES BETWEEN MOLECULES, HYDROGEN BONDING & SHAPES OF MOLECULES

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Hi guys! How are you? Are you ready for the next post on this topic?

Today, we are going to explore the forces between molecules. The intermolecular and intramolecular forces. Furthermore, the hydrogen bonding, present e.g. in ice and water. Lastly, we are going to examine the shapes of molecules, and the ones you are required to draw and understand for the exam. I hope this post will clear any obstacles that you came up with during your revision.

Forces between molecules

So, there are two different types of forces between molecules. And these are:

1. Intramolecular forces (stronger) - they are the forces within molecules. This is for an instance: covalent bonding, ionic bonding and metallic bonding. We talked about these bondings in the last post. Visit this post if you're not sure you can explain what each of them mean.

2. Intermolecular forces (weaker) - they are the forces between molecules. An example of these is hydrogen bonding or Van der Waals. These forces govern the physical properties of molecules.

Last unit was all about bonding and this unit is going to be about intermolecular forces. And more specifically, we'll start with Van der Waals and then go to hydrogen bonding later.

Van der Waals:

There are two types of Van der Waals forces:

1.Between molecules having a permanent dipole. This means one of them is slightly positive and the other is slightly negative. Like the examples in the previous post, where the electronegativity difference is so big that one of the atoms has more control over the electrons than the other. This doesn't change, one of the atoms is more positive and one more negative continuously.

e.g. hydrogen and chlorine, where chlorine is δ- and hydrogen δ+. As in the example above, there are Van der Waals between chlorine of one HCl molecule and hydrogen of the other HCl molecule.

However, we also have the other type of Van der Waals, which is:

2. Between molecules having induced dipole. In this example, both of the atoms have the same control over the electrons which means they both have the same electronegativity value and none of them is more positive or negative than the other.

e.g. two helium atoms; It is called induced dipole because all electrons move in an atom and there may be temporary dipoles created. For instance, one side becomes positive and the other becomes negative. Therefore, can attract adjacent molecules. There are Van der Waals forces between these two helium atoms.

The strength of Van der Waals forces increases:

-as the size of atoms increases because there are more electrons present;

-as the surface area increases;

-as the molecules are thinner and longer;

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Hydrogen bonding

Hydrogen bonding- this is stronger than Van der Waals but not stronger than intramolecular forces (covalent bonding, ionic bonding, etc.) A hydrogen bond is a bond formed between hydrogen and one of either oxygen, nitrogen or fluorine.

e.g. Hydrogen bonding in water:

You can see how it looks on the picture above. You show hydrogen bonding by the little lines between oxygen and hydrogen. In water, the water molecules (H2O) bond together from every side, this is just an example of a couple.

Properties:

1. Water has a high boiling point- due to the strong hydrogen bonds. Look at the diagram on the picture above. As you can see the hydrogen-bonded molecules have a much higher boiling temperature. The reason for this is, there are a lot of hydrogen bonds in water and they all need to be broken for boiling to occur (a lot of energy needed).

2. Ice has a high melting point - ice is a solid form of water and therefore it also has hydrogen bonding. The high melting point is also due to the strong hydrogen bonds.

Ice is also less dense than water as a result of hydrogen bonds holding water molecules apart in an open crystal lattice structure.

When the ice starts to melt, the hydrogen bonding begins to break up and the open crystal lattice collapses and density increases.

Solubility - polar molecules dissolve in polar molecules and non-polar molecules dissolve in non-polar molecules. For example, you can dissolve anything in water that contains O, N or F. Otherwise, it will not dissolve, just like oil.

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Shapes of molecules

Let's start with the basics of identifying shapes of molecules. The first thing you need to know is what the shape of molecules depends on:

1.Bonding pairs on an atom - bonded pairs of electrons;

2.Lone pairs on an atom - unbonded pairs of electrons;

I prepared an example for you, using a water molecule. The best way (at least for me, maybe you can figure out a different way that suits you better) is to draw the molecule out first. In water, we have one oxygen atom surrounded by two hydrogen atoms. Oxygen is present in group 6 on the periodic table and therefore has 6 electrons in its outer shell. Both hydrogens have one electron in their outer shells and bond covalently to the oxygen.

To conclude we have:

-two bonding pairs of electrons (the ones that bond with hydrogen);

-two lone pairs of electrons (the other 4 electrons that oxygen has, which is two pairs);

These lone pairs and bonding pairs arrange themselves, so that they minimise the amount of repulsion.

VSEPR theory - helps us predict the shapes of molecules.

In this theory, you firstly draw a dot and cross diagram and count the number of bonding electron pairs and lone electron pairs.

And here are the different shapes of molecules that are required to understand for a-level:

They are the simple shapes of molecules where you only have bonding pairs of electrons and none lone pairs. The shape changes as the number of bonded electron pairs change. Look at the picture above. I included the names of the shapes, diagrams for all the shapes and the bond angles.

I don't think there is anything else there is to say. You have to remember how these shapes look. And the best way to do this is just draw them over and over again and they will stick. Also, remember to memorise the names of the shapes.

The next picture, shows some of the more complicated shapes of molecules. And what I mean by complicated, is that we also have lone pairs present. When you have lone pairs, the shape changes a little.

e.g. 1 lone pair:

the name changes to trigonal pyramidal because we have one more bond, which is a lone pair. What also changes is the bond angle. Normally, when you have 4 bonding bonds, the bond angle is 109.5, but when it changes to 3 bonding pairs and 1 lone pair (also giving 4 pairs) the bond angle is 107.5.

e.g. 2 lone pairs:

the name for this one changes to bent tetrahedron because we have 2 lone pairs apart from the the two bonding pairs. The bond angle changes from 109.5 degrees (4 bonding pairs) to 104 degrees or sometimes written as 104.5 (2 bonding pairs and 2 lone pairs, which is also 4 pairs all together).

Shapes you need to know:

They are the shapes that you are required to know and explain in your exam. However, you may also be asked to predict and explain the shapes of other simple molecules, that have up to 6 electron pairs.

1. Boron fluoride - 3 bonding pairs.:

-trigonal planar shape;

-bond angle of 120 degrees;

2. Methane - 4 bonding pairs.:

-tetrahedral shape;

-bond angle of 109.5 degrees;

3. Ammonium - 4 bonding pairs.:

-tetrahedral shape;

-bond angle of 109.5 degrees;

4.Ammonia - 3 bonding pairs and 1 lone pair.:

-trigonal pyramidal shape;

-bond angle 107.5 degrees;

5. Sulfur hexafluoride - 6 bonding pairs.:

-octahedral shape;

-bond angle of 90 degrees;

If you just think about how many bonding pairs of electrons there are and how many lone pair of electrons there are, you can very easily figure out the shape of the molecule. If you are not sure yet, you can go back to the table with all the shapes and use this as a guide.

This is all for this topic. I hope it helps you with your revision. If you liked it, don't forget to share my blog with others. If you would like to practice some questions on this topic, don't hesitate and visit my youtube channel.

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