Today, 23rd October, is Mole Day – which might put you in mind of small, furry, burrowing animals. However, they don’t even seem to have a commemorative day of any kind; we’re actually talking about the mole in chemistry, a quantity that essentially allows us to ‘count’ atoms and molecules in a more convenient way. This is a fundamental concept, and one that all chemists utilise.
The problem that chemists have with atoms and molecules is that they’re a little on the small side, which can make counting them more than a little difficult. To put the size of atoms into perspective, if we scaled up the size of a helium atom so it was the size of one of the full stops in this article, scaling up an average gerbil’s size by the same factor would give you one terrifying, Earth-sized gerbil. When carrying out chemical reactions, we often want to get some idea of how much of the product we can expect to get from a reaction, but the enormous numbers of atoms or molecules involved in even a small sample would make these calculations incredibly unwieldy.
This is where the mole comes in. It’s defined as the number of atoms (or molecules) present in one mole of a substance: 6.022 x 1023. French scientist Jean Perrin proposed naming this number after Amedeo Avogadro, the Italian scientist widely credited with being the first to realise that the volume of a gas, at a given temperature and pressure, is directly proportional to the number of atoms or molecules, regardless of the gas’s identity. Hence, we know 6.022 x 1023 as ‘Avogadro’s number’ (NA).
Avogadro’s number is defined as being the number of atoms present in 12 grams of carbon-12. One mole of any substance, then, contains exactly this number of atoms or molecules. Obviously, dependent on the mass of the atoms or molecules, the total mass of one mole of a substance can vary – one mole of water weighs very marginally over 18 grams, whilst one mole of the salt you have in your kitchen, sodium chloride, weighs 58.4 grams.
This can seem confusing for students of chemistry at first; as an analogy, it’s easy to compare the term ‘mole’ to the more everyday term, ‘a dozen’. When someone says ‘a dozen’, you know they mean twelve of something. In the same way, chemists know that, when they’re referring to a mole, they’re talking about 6.022 x 1023 of something. Similarly, if you have a dozen elephants and a dozen mice, you’d expect the dozen elephants to have a significantly greater mass than the dozen mice (unless you’ve got your hands on some freakishly small elephants). However, despite the difference in mass, you still have a dozen of each. In the same way, the same number of moles of different elements and compounds can be different masses.
That’s what a mole is – but why do we need it? Well, for starters, it makes expressing amounts of chemicals a lot easier. We don’t have to represent the number of molecules of a compound we have, and use the large numbers that that would entail, and we can instead use moles in our calculations to simplify them. We can find the number of moles of any mass of any compound, simply by dividing the mass we have of the compound by the ‘molar mass’ – that is, the sum total of the masses of the atoms that make up the compound. Additionally, using moles can help us easily predict the masses of products we can obtain from reactions.
In summary, then, the mole is a vital tool for chemists – probably the reason there’s a day to commemorate it, whilst its furrier namesakes go the whole year unappreciated.
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