REACTIONS OF THE GROUP 1 ELEMENTS WITH WATER

This page looks at the reactions of the Group 1 elements - lithium, sodium, potassium, rubidium and caesium - with water. It uses these reactions to explore the trend in reactivity in Group 1.

The Facts

General

All of these metals react vigorously or even explosively with cold water. In each case, a solution of the metal hydroxide is produced together with hydrogen gas.

2X(s)  +  H₂O(l)    2XOH(aq)  +  H₂(g)

This equation applies to any of these metals and water - just replace the X by the symbol you want.

In each of the following descriptions, I am assuming a very small bit of the metal is dropped into water in a fairly large container.

Details for the individual metals

Lithium

Lithium's density is only about half that of water so it floats on the surface, gently fizzing and giving off hydrogen. It gradually reacts and disappears, forming a colourless solution of lithium hydroxide. The reaction generates heat too slowly and lithium's melting point is too high for it to melt (see sodium below).

Sodium

Sodium also floats on the surface, but enough heat is given off to melt the sodium (sodium has a lower melting point than lithium and the reaction produces heat faster) and it melts almost at once to form a small silvery ball that dashes around the surface. A white trail of sodium hydroxide is seen in the water under the sodium, but this soon dissolves to give a colourless solution of sodium hydroxide.

The sodium moves because it is pushed around by the hydrogen which is given off during the reaction. If the sodium becomes trapped on the side of the container, the hydrogen may catch fire to burn with an orange flame. The colour is due to contamination of the normally blue hydrogen flame with sodium compounds.

Potassium

Potassium behaves rather like sodium except that the reaction is faster and enough heat is given off to set light to the hydrogen. This time the normal hydrogen flame is contaminated by potassium compounds and so is coloured lilac (a faintly bluish pink).

Rubidium

Rubidium is denser than water and so sinks. It reacts violently and immediately, with everything spitting out of the container again. Rubidium hydroxide solution and hydrogen are formed.

Caesium

Caesium explodes on contact with water, quite possibly shattering the container. Caesium hydroxide and hydrogen are formed

Summary of the trend in reactivity

The Group 1 metals become more reactive towards water as you go down the Group.

Explaining the trend in reactivity

Looking at the enthalpy changes for the reactions

The overall enthalpy changes

You might think that because the reactions get more dramatic as you go down the Group, the amount of heat given off increases as you go from lithium to caesium. Not so!

The table gives estimates of the enthalpy change for each of the elements undergoing the reaction:

X(s)  +  H₂O(l)    XOH(aq)  +  ½H₂(g)

You will see that there is no pattern at all in these values. They are all fairly similar and, surprisingly, lithium is the metal which releases the most heat during the reaction!

Digging around in the enthalpy changes

When these reactions happen, the differences between them lie entirely in what is happening to the metal atoms present. In each case, you start with metal atoms in a solid and end up with metal ions in solution.

Overall, what happens to the metal is this:

X(s)    X⁺(aq)  +  e^-

You can calculate the overall enthalpy change for this process by using Hess's Law and breaking it up into several steps that we know the enthalpy changes for.

First, you would need to supply atomisation energy to give gaseous atoms of the metal.

X(s)    X(g)

Then ionise the metal by supplying its first ionisation energy.

X(g)    X⁺(g)  +  e^-

And finally, you would get hydration enthalpy released when the gaseous ion comes into contact with water.

X⁺(g)  +  (aq)    X⁺(aq)

If we put values for all these steps into a table, they look like this (all values in kJ / mol):

So why isn't there any pattern in these values? If you look at the various bits of information, you will find that as you go down the Group each of them decreases:

• The atomisation energy is a measure of the strength of the metallic bond in each element. This is falling as the atom gets bigger and the metallic bond is getting longer. The delocalised electrons are further from the attraction of the nuclei in the bigger atoms.

• The first ionisation energy is falling because the electron being removed is getting more distant from the nucleus. The extra protons in the nucleus are screened by additional layers of electrons.

• The hydration enthalpy is a measure of the attraction between the metal ions and lone pairs on water molecules. As the ions get bigger, the water molecules are further from the attraction of the nucleus. The extra protons in the nucleus are again screened by the extra layers of electrons.

What is happening is that the various factors are falling at different rates. That destroys any overall pattern.

It is, however, possible to look at the table again and find a pattern which is useful.

Drawing something useful from these figures

Let's take the last table and just look at the energy input terms - the two processes where you have to supply energy to make them work. In other words, we will miss out the hydration enthalpy term and just add up the other two.

Now you can see that there is a steady fall as you go down the Group. As you go from lithium to caesium, you need to put less energy into the reaction to get a positive ion formed. This energy will be recovered later on (plus quite a lot more!), but has to be supplied initially.

This isn't literally the activation energy for the reaction, because atomisation energy and ionisation energy have quite specific meanings which don't apply here. But the actual activation energy will be affected by these two factors - the need to break the atom away from the lattice and its ionisation.

Both of these things become easier as you go down the group, and so the reactions get faster.

So although lithium releases most heat during the reaction, it does it relatively slowly - it isn't all released in one short, sharp burst. With caesium, on the other hand, although it doesn't release quite as much heat overall, it does it extremely quickly - and you get an explosion.

Summarising the reason for the increase in reactivity as you go down the Group

The reactions become easier as the energy needed to form positive ions falls. This is in part due to a decrease in ionisation energy as you go down the Group, and in part to a fall in atomisation energy reflecting weaker metallic bonds as you go from lithium to caesium. This leads to lower activation energies, and therefore faster reactions.