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Why the Mole Concept is the Most Important High School Chemistry Concept

Why Chemists Need the Mole Concept

The mole may just seem like a number, an awkward number that may seem initially to make chemistry calculations that bit more complicated.

But it’s much more than just a number. And it’s actually there to make things simpler whenever we need to use or calculate the amount of a substance.
And the mole is used throughout your chemistry syllabus, in most chemistry calculations you will perform. It is fundamental to your understanding of kinetics, thermodynamics, entropy and much more.

Why do we need the mole, and how does it make things simpler?

Chemical equations represent the ratio that molecules react with each other. They show how many molecules of each reactant are needed for an exactly balanced equation, and how many molecules of each product substance will be produced in that reaction.

But in reality, we don’t deal with individual molecules. We can’t , of course, because they are way too small to handle or observe.
We tend to measure substances by mass or volume, usually in grams or millilitres. Even a crystal or two of a salt, or a drop of a liquid or solution still contains an immense number of molecules.

And we can’t use mass directly as a proxy for the number of molecules. Why not? Because each element’s atoms have different mass, and these have a vast range of masses. Using mass to reference the of amount of substance is seriously flawed.
A simple comparison from the periodic table shows that one gram of hydrogen contains about16 times more molecules than one gram of oxygen. If we were to compare the masses of each used in an experiment, we would be skewing the true ratio by a factor of 16.
Using volume would be even tricker because densities change with temperature.

Chemists do need to use a count of molecules, but not individual molecules. We need a unit for a specific bulk quantity of molecules.
Computer science has gigabytes to define bulk data quantities, rather than use massive numbers of individual bytes. Gigabytes makes it much easier to visualise and handle. Likewise chemistry has a bulk unit of molecules to make it much easier to use and calculate information about substances. That unit is the mole.

Listen to my explanation of the mole:

Exactly What Is a Mole?

So the mole is the standard unit used internationally to handle the number of molecules bulked up to a scale we actually use in the lab, or in a calculation.

A mole contains a huge but specific number of molecules or particles. The number of particles in a mole is 6.022 x 1023.
(Technically, a mole is defined as 6.02214076 x 1023 particles but we don’t generally use all those significant figures. Stick to four significant figures at A-Level and other high school, college or undergraduate use).

The mole is given the abbreviation “mol”.

‘One mole contains exactly 6.022 x 1023 particles’

SI definition of the mole (paraphrased). Elemental particles are usually molecules in this context.

You may hear the number 6.022 x 1023 referred to as the Avogadro number.
The Avogadro Constant, NA, is 6.022 x 1023 mol-1.

You may be wondering why this number was chosen. It certainly wasn’t chosen for ease of use, so what is its significance?

Why is the Avogadro Number So Big (and so….awkward)?

The Avogadro number wasn’t chosen for convenience, and doesn’t follow the usual practice of using factors of 103, 106, 109 and so on to refer to multiplication my 1000, 1000000 or 1000000000 etc.

It was chosen for one specific reason. It is the exact number of atoms in a specific quantity of a reference substance.

That reference substance is 12C (or “carbon-12”), a sample comprised entirely from the most common isotope of carbon.
The reference quantity of 12C is exactly 12 grams.

So, the Avogadro number 6.022 x 1023 is the exact number of atoms in exactly 12 g of 12C.

How to Use the Mole in Chemistry

The primary use of the mole is to calculate or predict the amount of a substance used, required or produced in a reaction.

You can do this for real reactions such as the practical experiments you do in the college lab. You do so along with the molar masses of the reactants or products, and we’ll look at how to calculate molar mass in a moment.

And you can use the mole in calculations such as exam questions or assignments. Virtually every chemistry calculation uses the mole or calculates something relative to the mole.

What is Molar Mass, and How is it Calculated?

The molar mass of a substance is simply the mass of exactly one mole of a substance, in grams.

To calculate molar mass, add up the atomic mass of each atom in the unit formula of a substance. The atomic mass for any atom is shown on your periodic table.

What do we mean by the unit formula? This is the smallest unit that a substance can be broken down to without it becoming a different substance. This is the same as the formula we use in a chemical equation.

Let’s consider some examples:

Water, H2O. The unit formula is straightforward because water exists as individual molecules.
The molar mass of water is the sum of the atomic mass of each atom – two hydrogen atoms and one hydrogen atom. Hydrogen has atomic mass 1.0 and oxygen has atomic mass 16.0, so the molar mass of water is 1 + 1 + 16 = 18 grams.

Oxygen, O2. Oxygen naturally exists as a diatomic molecule, rather than as individual atoms, and we must consider this when we calculate the molar mass of oxygen. A mole of oxygen is a mole of O2 molecules. The molar mass is 32 grams (because the atomic mass for each oxygen atom is 16).

Silicon dioxide, SiO2. Silicon dioxide is macromolecular, meaning it has molecules that go on and on … there is no defined size for the molecule. When we write a chemical equation for a reaction involving silicon dioxide, we use SiO2 as this is the unit formula for the smallest unit of this substance.
When calculating the molar mass we need to calculate for a mass of that + unit formula. SO we add together the atomic m + ass of one silicon and two oxygen atoms. The molar mass is therefore 60 grams (28 + 16 + 16).

How to Use Molar Mass

In many chemistry calculations we need to use the number of moles of a substance.

To calculate the number of moles just divide the mass of a substance used by its molar mass to get the number of moles used.
For example if 15 grams of CaCO3 were used in an experiment, just divide 15g by the molar mass. The molar mass of CaCO3 is 100grams, so the number of moles used is 15/100 = 0.15 mol.

(The calculation works equally well for reaction products as it does for reactants).

It’s also possible to calculate the mass of a substance used / produced if the number of moles is known. This is something you are can do after you have calculated the number of moles using n equation, such as the ideal gas equation (or any equation that references moles).
For example, imagine you used some data given in a chemistry question to calculate that the number of moles of CO2 present in a reaction is 0.3 moles. That sort of question will usually ask you to calculate the mass of the CO2 too. To do this simply multiply the number of moles by the molar mass. The molar mass of carbon dioxide is 44 grams, so the mass of 0.3 moles is 0.3 x 44 = 13.2g.

In fact you will use moles in almost every chemistry calculation. This is one reason why it is so important to understand the concept of the mole, and to become proficient when using moles in calculations.
I recommend highly that you should practice chemistry calculations to get comfortable with the many different applications of moles, as well as getting used to the types of calculations you are expected to be able to do.
If you’re studying A-Level, Advanced Highers or any program with a final exam you can usually find past papers on your exam board’s website; you should be able to download and practice these easily:

OCR Chemistry (A) AS Level OCR Chemistry (A) A Level

OCR Chemistry (Salters) AS Level OCR Chemistry (Salters) A Level

AQA Chemistry AS and A-Level

SQA Advance Higher Chemistry SQA Higher Chemistry

WJEC Chemistry AS and A Level Eduqas Chemistry AS and A Level

CCEA Chemistry Advanced GCE

Edexcel Chemistry A Level

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