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I know that the mole is widely used in chemistry instead of units of mass or volume as a convenient way to express amounts of reactants or of products of chemical reactions.

I'm wondering why people in chemistry still excessively use it for their measurement? To be backward-compatible and consistent with traditional textbooks? Why they don't simply express their quantity measurement per atom, per unit volume, per molecule or etc.?

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    $\begingroup$ It's similar to why people use words like "dozen", "score", "gross", "ream", etc. $\endgroup$ – DavePhD Dec 16 '15 at 17:19
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Simply speaking, because it's an appropriate unit to use.

Let's imagine I wanted to measure the length of a rope. What would be an appropriate length to use? Inches? Centimeters? Feet, maybe? It would really be awkward to express it as 0.000189393 miles, or as 304,800,000 nanometers.

(Note: if you can't see why these units are awkward, take any page discussing things like this (e.g. biomass of certain species) and change all the units so that they're nonsense like this. Then put it away for a week and try to read it later.)

Now let's say I've changed my mind and I'd like to measure all the rope created in the world in a year. Now an appropriate unit is almost certainly miles or kilometers, and not inches or centimeters.

Let's consider something else: so far, I've been using length to measure ropes. Would it make sense to measure their combined mass instead? Maybe not for small amounts, since I think I would throttle you if you told me to cut off half a pound of rope, but for global-scale things, tons or metric tons may make sense.

On the other hand, using measurements like the average density of the rope or the combined diameters of all the ropes really wouldn't be much use at all.


What we've seen here is that when we're measuring things, there are measures that make sense (for rope, length, maybe weight) and some measures that really don't (color, average diameter, etc.). In those measures, there are units that are convenient (inches, feet), and some that aren't (nanometers).

This is exactly the problem with chemical units, but much more magnified. When I measure the energy released on hydrolysis of a sample, I'm not measuring the energy of one bond, or two bonds, or a thousand bonds, or even a million. I'm measuring a collection of so insanely many molecules that there's really no number in common language for it.

What is an appropriate way to measure things in these molecules? If I'm only interested in how much I have and not what's in it, per gram or per volume is often a good way to do it. Since when I'm measuring the temperature change of something, I don't particularly care what it's made of, measures of this (specific heat, for instance) are done in units of something per gram.

On the other hand, if I do care what my sample is made of, then I need a better way to measure it. For example, sodium chloride and calcium chloride look similar, but will have very different energies since there's three ions in a $\ce{CaCl2}$ unit and only two in an $\ce{NaCl}$ unit.

Since you can't have half a molecule, the simplest way to do it is to count the darn things. Unfortunately, measuring things per atom is a really awkward way to do things because of the problems I described above. The energy it takes to melt ice is 0.000000000000000000001029 J/molecule.

What we need then, is some count of molecules that's convenient. We could go by 1000 molecules, or a million molecules, but it's a pain in the butt to convert between that and macroscopic units (how much does 1 million calcium atoms weigh?) Now it just so happens that $6.022 \times 10^{23}$ atoms of carbon-12 weigh 12 grams, and a single carbon-12 weighs 12 amu. A single molecule of water weighs roughly 18 amu, $6.022 \times 10^{23}$ molecules of water weigh 18 grams.

Let's take a look at what using $6.022 \times 10^{23}$ molecules offers us over other units:

  • It is appropriate. $6.022 \times 10^{23}$ molecules will typically
    fall in the gram to kilogram range of substance, which can easily be measured out on a balance. It's also within the range of what you'd usually use in a lab.
  • It allows for easy conversion. You can easily tell when you have $6.022 \times 10^{23}$ molecules because your atomic mass is equal to your macroscopic mass.
  • It makes sense. Intuitively, we'd like to measure things per atom or per molecule, but doing so leads to some ugly units. Instead, if we
    measure per $6.022 \times 10^{23}$ molecules, we still are doing a
    counting-based measurement, but the numbers themselves are a lot
    nicer.

You probably know this already, but we call $6.022 \times 10^{23}$ molecules a mole. These are the advantages of using the mole as a unit, and not for backwards-compatibility. I am of the opinion that if we shed all our units today and started from scratch, we would start off with grams and liters, but we would probably start using the measure of a mole again within a year. It's just a very powerful, useful way to measure things.

P.S. This is much longer than I was initially intending and was not written while I was in the most awake state of mind. Please let me know in the comments if you think this answer is unclear, or stupid, or just plain wrong.

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Your statement that the "mole is widely used in chemistry instead of units of mass or volume.." is misleading. There are seven base units of measurement in the SI system and your question alludes to two, mass and amount. The other 5 are length, time, temperature, electrical current, and luminosity.

Consider going out for a typical (American) breakfast of bacon and eggs. There are two diners that you like to visit; one serves small-sized eggs and bacon strips while the other uses jumbo eggs and offers canadian bacon. At each place, you order "Two scrambled eggs with three slices of bacon." But do you get the same amount of food at each diner?

A rash of bacon weights about 22 grams and a slice of canadian bacon is a bit closer to 50 g, so three slices of each weigh 66 and 150 grams, respectively.

A small chicken egg weighs about 43 grams while a jumbo egg weighs 71 g. These weights are with the shell, and you visit good diners that give you scrambled eggs without the shell, but the point should be clear: while the amount of items ordered is the same, the weight of items is different.

Colloquially, we use the term amount when we are interested in answering a question beginning with "How much...". The units of our response oftentimes depends on practical considerations about how we measure the substance. From the deli, I'd order a pound of ham and two balls of mozzarella. In this case using units of weights or amount could be valid (I'd like 12 slices of ham and 100 grams of mozzarella.) If I'm waiting in line, however, I would say "there are 3 patrons before me" not "I have to wait for 425 pounds of people before I can be served."

Last a comment about the mole. The International Bureau of Weights and Measures has a good definition of the mole that doesn't need to be repeated here. To carry out the food analogy just one more step, the mole is the same type of word as a dozen. When you go to the store to prepare for a large breakfast, do you think to yourself, "I'm going to buy twenty four eggs" or "I'm going to by 2 dozen eggs". The word dozen, which stands for 12, is a way to simplify or group our numbers. Along the same lines, it's much easier to say 12 grams of carbon contains 1 mol of carbon atoms instead of $6.022 \times 10^{23}$ atoms.

In summary The mole is used extensively in the sciences because we need a unit that describes an amount of substance, which is different from the mass of a substance or how much space a given amount of substance occupies.

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  • $\begingroup$ Haha, one (1) is a good enough unit of amount I guess, but I'm no expert. $\endgroup$ – m93a Feb 18 at 10:16
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The mole is a really good bridge, in a manner of speaking, between MASS vs. NUMBER, when talking about things that are really really small ! Once you balance several dozen equations, especially for a practical lab use----like, to make something----you readily see that , in order to determine the proper masses for a reaction to go forward, you need some kind of 'bridge' to convert a huge sum of microscopic energies [ molecule to molecule ] into a quick, readable format. The mole does just that : It gives you a way to quickly read the "energy ledger book" for reactions. If $\ce{C2H6 + O2 -> CO2 + H2O}$, you know that it takes 1 mole of C2H6 to do that reaction , and that your product will be 1 mole of $\ce{CO2}$ and 1 mole of $\ce{H2O}$. IT ALSO MEANS you can weigh out the molar mass of 1 mole of $\ce{C2H6}$, which is ______ (you do the math, it's on the periodic table for you for each element---how convenient is that !), and you will have assurance that you are using the right amt of reactant !

From this standpoint it does not take long to see that the mole is a pretty indispensable tool.

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Consider this: you are comparing the effectiveness of acetic acid vs stearic acid (solid) as a drug. The former is a liquid and per molecule weighs about 5x less.

If you used 1 g of each, you would be adding different numbers of molecules. If you added 1 mL of each, same problem and compounded by differences in density. Since mole is a per molecule measurement, it allows you to make real comparisons between compounds. For example 1 mole of acetic acid will have the same number of molecules as 1 mole of stearic acid even though the weights, volumes, and phases might be different. Why biology journals still report concentration as ppb or ppm is beyond me, because if there is a large difference in MW weight between compounds, you cant make comparisons.

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  • $\begingroup$ It actually makes perfect sense to report some concentrations as ppm or ppb. Technically, it is only a very small molar fraction — but nonetheless measureable. And not only biologists do it, chemistry journals do so to. (Remind me to link a JACS paper using it, as soon as I run across one again.) Of course, one should never confuse ppm for a weight per volume measurement; it is always molar fraction. Other than that, $+1$. $\endgroup$ – Jan Oct 23 '16 at 19:47

protected by Loong Nov 14 '16 at 18:45

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