Chemguide: Support for CIE A level Chemistry
Learning outcome 2.3.2
This statement deals with writing equations.
Before you go on, you should find and read the statements in your copy of the syllabus.
You have to know what you start with and what you end up with. In an introductory course, you may have used word equations, but word equations are not acceptable at this level. Nevertheless, you have to have a word equation in your mind.
To take a trivial example, suppose you want an equation for burning magnesium in air or oxygen to give magnesium oxide. I am taking this because it lets me make three important points.
The unbalanced equation is:
The first point I want to make is that you have to know that certain common elements have molecules consisting of two atoms. Oxygen is one of them, but others you are likely to come across are H2, N2, F2, Cl2, Br2 and I2.
To balance the equation it often works best if you just examine the atoms from left to right.
So an initial work-through gives:
The second important point is that you must never change the formula of anything when you are doing this. You must not change the formula of the magnesium oxide to MgO2 - there is no such thing.
The third point is that once you have worked across the equation, always go back and check. In this case, if you do that you find that the magnesiums no longer balance. The final version is:
And there is one further tip that you might find useful.
When you heat copper(II) nitrate, you get copper(II) oxide, nitrogen dioxide and oxygen produced. The unbalanced equation is:
When you work across the equation checking things, it is often useful to leave things which occur in more than one place on each side until the end. They often sort themselves out. Oxygen is a case in point here.
That now leaves a tricky problem with the oxygens. There are 6 (2 lots of 3) on the left-hand side, and 7 on the right. It isn't immediately obvious how to sort that out.
The trick is to allow yourself to have half an oxygen molecule - half an oxygen molecule is just one oxygen atom.
Now there are 6 oxygen atoms on both sides.
You wouldn't normally leave it like this though. Having halves in equations feels wrong when you are starting chemistry, but they aren't uncommon. You can get rid of the half by doubling everything in the equation.
Don't worry if you need half of a diatomic molecule in a first attempt to balance an equation. Accept it and then get rid of it by doubling everything.
State symbols in equations
But quite often you will find equations which have state symbols added in brackets after each formula. There are four of these, and they tell you what physical state the substance is in.
So, for example, solid magnesium burns in steam to produce a white powder, magnesium oxide, and hydrogen gas. That would be shown using state symbols as . . .
In this case, the state symbols matter, because magnesium has only a very, very slight reaction with cold water, and produces a different product - magnesium hydroxide rather than magnesium oxide.
In an exam, if in doubt, include state symbols in each equation. But take your lead from the examiners. If they aren't using state symbols themselves in the question, you could probably safely leave them out. But if you include them, make sure that they are right.
There is one case where state symbols are almost invariable used, and that is in ionic equations - coming up below.
You usually first come across ionic equations when you are looking at precipitation reactions - for example, in the reaction between silver nitrate solution and sodium chloride solution to give a white precipitate of silver chloride.
The full equation for this reaction looks like this:
AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
But this hides something interesting.
Everything in this reaction is ionic, and most of the compounds have their free ions moving around in solution. In the silver chloride case, though, the ions have clumped together to produce a solid.
If you rewrite the equation to show the freely moving ions, it looks like this:
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) AgCl(s) + Na+(aq) + NO3-(aq)
If you look at this carefully, you will see that the sodium ions and nitrate ions aren't changed in any way. They are known as spectator ions - they simply aren't involved in the reaction.
Spectator ions are not included in ionic equations. If you include them, according to the syllabus, you will be penalised.
The proper ionic equation leaves them out, because they are unchanged in any way:
Ag+(aq) + Cl-(aq) AgCl(s)
This equation says something really useful.
Any source of silver ions in solution will react with any source of chloride ions in solution to give a precipitate of silver chloride. This is the basis for the chemical test for a chloride.
The best way of writing ionic equations for a precipitation reaction is to start by writing the formula for the precipitate on the right-hand side, and then write down the ions which formed it on the left-hand side.
Another example, not a precipitation reaction this time, is the reaction between a solid carbonate and a dilute acid. The ionic equation is:
2H+(aq) + CO32-(s) CO2(g) + H2O(l)
What this says is that any dilute acid (a source of hydrogen ions in solution) will react with any solid carbonate to produce carbon dioxide gas and water.
Another common place where ionic equations are used is in redox reactions - reactions involving oxidation and reduction.
For example, chlorine gas bubbled through a solution containing iron(II) ions oxidises them to iron(III) ions. It doesn't matter which iron(II) compound you use (for example, iron(II) chloride or iron(II) sulfate) because the chloride or sulfate ions are spectator ions.
2Fe2+(aq) + Cl2(g) 2Fe3+(aq) + 2Cl-
Don't worry too much about these redox equations for now.
During the rest of the A level course, you will have to learn how to write equations for more complex reactions involving oxidation and reduction. These are discussed in detail if you explore the redox menu on the main part of Chemguide.
You can ignore the section about oxidation states for the moment, and I wouldn't worry too much about writing redox equations for reactions done under alkaline conditions (the link from the bottom of the redox equations page).
Whether you look at that first redox equations page now or leave it until the situation arises during your course probably doesn't really matter. But whatever you do, don't go on to the alkaline conditions page until you are completely confident about these equations under acidic or neutral conditions. You will end up totally confused!
Writing equations is covered in detail in the first chapter of my A level chemistry calculations book. I can't give you any more about this without upsetting my publishers.
© Jim Clark 2019