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Chemistry I | © Ray Lovegrove |
Chemical
Equations – why we need them
Will I need to write and balance chemical equations to study chemisty? Yes, but don’t worry if you're out of practice or if the whole process is new to you. You will have met some equations in the first few chapters of the course so you already know some basic principles. You may find it useful to think of the valency or combining power of individual elements to help you with the first stage of equation writing, which is getting the formula correct. The table below gives the valencies of the elements and one or two other groups of elements like sulphates and nitrates.
|
1 |
2 |
3 |
4 |
|
Na+ Li+ K+ H+ Cl– Br– I–
NH4+ Ammonium |
Ca2+ Mg2+ O2– S2– SO42– Sulphate CO32– Carbonate
|
Al3+ |
C4+ |
Using the table we can see that both sodium, Na, and bromine, Br, have valencies (combining powers) of one;
this tells us that sodium bromide will have the
formula NaBr.
We can also see that oxygen has a
valency of two
so the formula for sodium oxide would need two sodium
atoms combining with one oxygen to give the formula Na2O.
As a final example, calcium has a valency of two and the
nitrate group has a valency of one;
so one calcium atom would combine with two nitrates to
give calcium nitrate with the formula Ca(NO3)2.
Many students of chemistry at all levels have found problems with writing chemical
equations. If you are one of the lucky ones who find equations not too difficult, then
this lesson will be
fairly straightforward, but if like many you find the topic challenging, you may need
to make reference quite often to the content of this chapter as you progress
through the course.
Chemists use equations for a variety of reasons, which are summarized below. Despite the difficulty that students find in using them, they are essential if you wish to get a foothold in the subject.
Equations quantify reactions, which would otherwise simply be "recipies" of reactants.
Chemical equations, like mathematical formulae, are a truly international language, irrespective of alphabet or script differences.
Chemical equations actually provide information about reactions which can not easily be conveyed otherwise.
Equations are the key to fully understanding a chemical process.
Word Equations
A word equation is a simple but effective non-quantitative way of representing a chemical reaction. The substances that you have to start with are called reactants and those formed in the reaction are called products. In between the reactants and products is usually an arrow (=>); this represents that a chemical change has taken place. If you read out a chemical equation (over the phone for instance) the arrow is called ‘gives’ and never "equals". For instance, the word equation;
sodium hydroxide solution + hydrochloric acid => sodium chloride solution + water
This would be read out loud as; ‘a solution of sodium hydroxide plus hydrochloric acid gives sodium chloride solution plus water’. Word equations have their uses at all levels of chemistry but are always less descriptive than symbol equations.
Changing words to symbols.
Using the example above as our starting word equation we can now
start on the process of changing words into symbols.
sodium hydroxide solution + hydrochloric acid =>sodium chloride solution +water
NaOH + HCl => NaCl + H2O
In this case, the term ‘equation’ is fully justified as
the same numbers of atoms occur on both sides of the arrow.
We still do not have all the information however, as the state of reactants and
products needs
to be included. To do this, we add bracketed subscripts to indicate solid (s),
liquid (l), gas (g) or
solution in water (aq) from the Latin
aqueous, meaning "in water".
We are now getting to a symbol equation that is telling us much more than a word
equation could ever do.
Let us look at another example:
Magnesium + copper sulphate solution => magnesium sulphate
solution + copper
By substituting symbols for the words, and adding state symbols we get:-
Sometimes it is necessary to balance equations so that exactly the same numbers of the atoms of individual elements occur on both sides, for instance:
Magnesium + hydrochloric acid => magnesium chloride solution +
hydrogen gas
This becomes:
Mg(s) + HCl(aq)
=> MgCl2(aq) + H2(g)
The golden rule for balancing equations is that
formulae cannot be changed; only the number of reacting particles (atoms, molecules or
ions) can be changed. We can obviously see that we have only one chlorine ion
represented in the products, whereas there are two in the reactants, and there is a
similar situation for hydrogen – we have to balance the equation thus;
To look at a more complicated situation:
aluminium + sulphuric acid => aluminium sulphate solution + hydrogen gas
Introducing symbols and states we get:
Al(s) + H2SO4(aq) =>
Al2(SO4)3(aq) + H2(g)
This is clearly an unbalanced equation as both aluminium and the sulphate ion are in greater numbers in the products. To balance this equation we need to adjust both sides;
Sometimes you can see at once what needs to be done, and at other times you need to use trial and error to balance an equation. But don't worry, it gets much easier with practice.
Ionic Equations
Throughout this course we will be considering how ions change during the course of a chemical reaction. If we use equations to help us in this we need to modify the type of symbol equations given above into ionic equations. This is fairly straightforward but, like much chemistry, lots of experience with ionic equations will make them seem easy.
Using one of the equations given above:
Mg(s) + 2HCl(aq) => MgCl2(aq) + H2(g)
We can see that some of the ions have changed, whereas others remain the same. Magnesium has changed from an element into being part of a compound, and hydrogen has changed from being part of a compound into its molecular H2 form. The chloride ions however were part of a compound in the reactants and part of a compound in the products – we call ions like this which do not change spectator ions. In ionic equations, the situation of ions changing is made clear by removing all spectator ions from the picture – this gives us:
Mg(s) + 2H+
(aq) => Mg2+(aq) + H2(g)
Note that charges have to be added to the ions after the removal of
the spectators and that state symbols should still be included.
Another example from above:
2Al(s) + 3H2SO4(aq)
=> Al2(SO4)3(aq) + 3H2(g)
Note how the hydrogen ions involved in the two sulphuric acid species become six individual hydrogen ions in an ionic equation.
One further example from above:
Mg(s) + CuSO4(aq) => MgSO4(aq) + Cu(s)
becomes:
Again, as is very often the case, sulphate ions were the spectators.
Ion-electron (half-equations)
This represents the last stage in how specific a chemical equation can get. In ion-electron equations (in most text books these are called ‘half-equations’ because that’s exactly what they are) only one element from an ionic equation is looked at in isolation.
If we look yet again at the ionic equation:
Mg(s) + 2H+(aq) => Mg2+(aq) + H2(g)
This can be converted into ion-electron equations by first considering magnesium and then turning our attention to hydrogen. The two resulting ion-electron equations must balance, like any other equation. So we need to use electrons (e-) to even things up.
2H+(aq) + 2e- => H2(g)
Notice two things; firstly the electrons in both ion-electron equations must match one
another; and secondly, it is easier to place electrons on the side of the equation that
can keep them matching as ions.
For a final example, let us look again at:
2Al(s) + 6H+(aq) => 2Al3+(aq) + 3H2(g)
This will become;
2Al(s) => 2Al3+(aq) + 6e-
6H+(aq) + 6e- => 3H2(g)
You should have a good idea of how to change an ionic equation into a pair of ion-electron equations at this point. From now on in subsequent chapters we will take the opportunity to develop these skills.
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