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Thread: Moles

  1. #1
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    Moles

    Why do chemists use moles as units of quantity?

    https://en.wikipedia.org/wiki/Mole_(unit) says "the mole is not a true metric (i.e. measuring) unit, rather it is a parametric unit."

    It seems confusing to use a parametric unit when metric units of weight are available (grams etc).

    For example, an oceanography paper states "the molar ratios of C to Fe are of the order of 10,000 to 100,000:1." How could such a molar ratio be more informative than a weight ratio?

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    They are much easier to deal with when balancing reactions. Easier to say one mole of C reacts with 1 mole of O2 to make a mole of CO2, rather than 12.01g of carbon reacts with 32.00g of diatomic oxygen.

    Same with your example - it is saying for each atom of Fe there are 10-100K atoms of C. Which, when looking at reactions, is more useful.

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    Chemists are interested in the combining ratios. One mole of Oxygen will reliably combine with two moles of Hydrogen in forming H2O, whatever the weights involved.

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    Moles are not units of weight, they are units of amount.

    A mole of carbon-12 atoms are 6.02 x 1023 carbon atoms. A mole of hedgehogs are 6.02 x 1023 hedgehogs, though they would have a much greater weight than a mole of carbon atoms.

    For the reaction of 2 H2 + O2 -> 2 H2O (for example), you could express this as 4 grams of hydrogen and 32 grams of oxygen to produce 36 grams of water. But that does not give an obvious connection to the stoichiometry of the reaction. Expressing it as 2 moles of H2 and 1 mole of O2 reacting to form 2 moles water does relate to the stoichiometry.

    The conversion of chemical substances from weights to moles is with the atomic weight or molecular weight of the substance, and which has units of grams/mole.

    The wikipedia article on the mole, gives a pretty good explanation and the history of the concept.
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    In my early years I felt that chemists use moles so that no one else could figure out their secret formulas ;-).

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    Looking around the internet, I like this website for the history of the concept.

    Long before the mole concept was developed, there existed the idea of chemical equivalency in that specific amounts of various substances could react in a similar manner and to the same extent with another substance. Note that the historical equivalent is not the same as its modern counterpart, which involves electric charge. Also, the historical equivalent is not the same as a mole, but the two concepts are related in that they both indicate that different masses of two substances can react with the same amount of another substance.

    The idea of chemical equivalents was stated by Henry Cavendish in 1767, clarified by Jeremias Richter in 1795, and popularized by William Wollaston in 1814. Wollaston applied the concept to elements and defined it in such a way that one equivalent of an element corresponded to its atomic mass. Thus, when Wollaston's equivalent is expressed in grams, it is identical to a mole. It is not surprising then that the word "mole" is derived from "molekulargewicht" (German, meaning "molecular weight") and was coined in 1901 or 1902.
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    Quote Originally Posted by Spacedude View Post
    In my early years I felt that chemists use moles so that no one else could figure out their secret formulas ;-).
    You can play with the minds of the uninitiated with that: "I made the Avogadro Run in less than 12 moles!"

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    Quote Originally Posted by DonM435 View Post
    You can play with the minds of the uninitiated with that: "I made the Avogadro Run in less than 12 moles!"

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    So what’s a mole of moles?

    (slinks over to the “when you just have to make a joke” thread).


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    There seems to be some reasonable evidence, more than one line, for ~ 2 trillion galaxies. Assuming 300 billion stars/ galaxy .....

    It's pretty amazing the number of ways Avagadro's number has been determined, including sky scattering.
    Last edited by George; 2018-Jan-23 at 12:09 AM.
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    Quote Originally Posted by schlaugh View Post
    So what’s a mole of moles?
    xkcd's answer
    What would happen if you were to gather a mole (unit of measurement) of moles (the small furry critter) in one place?
    —Sean Rice

    Things get a bit gruesome.
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    Thanks all, makes good sense. My interest in molar chemistry comes from climate physics, where weight is a more accessible framework to a broad audience than stoichiometry.

    I used an excel converter linked here to answer my specific question on iron and carbon. It explains that an iron atom weighs 4.65 carbon atoms. That would mean the 100K molar ratio mentioned in the article I read means addition of an iron atom could theoretically enable reactions involving up to 465,000 times its weight in carbon in an iron limited sea location.

    Use of molar ratios is an issue I have come across several times in technical papers, where I have had to convert to weight to make sense of the argument. I have also seen scientists get this conversion wrong, which seems an easy mistake to make. It makes me wonder if when chemists are writing for a broader audience they should more often convert their atomic ratio calculations into weight measures to make them more accessible.

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    Quote Originally Posted by Robert Tulip View Post
    It makes me wonder if when chemists are writing for a broader audience they should more often convert their atomic ratio calculations into weight measures to make them more accessible.
    Maybe.

    It depends on the topic and the audience. There are also traditions in some related disciplines. For example, in metallurgy and ceramics, compositions and impurities are usually described by weight % or weight ppm, not molar quantities. I don't know why it was originally done like that, and why it is still done that way, but it is.
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    Quote Originally Posted by Swift View Post
    A mole of hedgehogs are 6.02 x 1023 hedgehogs, though they would have a much greater weight than a mole of carbon atoms.
    Never really understood the whole deal with hedgehogs. I mean, why can't they just share the hedge?

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    Quote Originally Posted by Swift View Post
    Moles are not units of weight, they are units of amount.
    I say they are units of number rather than "amount". I consider
    "amount" to be a more general term which includes all quantities
    and all units. Three and a half quarts is an amount of volume,
    nine yards is an amount of length or distance, seven amperes is
    an amount of electric current, 41 minutes is an amount of time,
    32 calories is an amount of energy. Three dozen is a number.
    Twelve moles is a number. Both of those are also amounts, but
    they are more specifically numbers. A mole is more specifically
    a number of particles. A mole is not a particle, in my opinion, so
    you can't have a mole of moles.

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    Which begs the question, are atoms particles?

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    Quote Originally Posted by grapes View Post
    Which begs the question, are atoms particles?
    Perhaps they are really icles, so they are part icles and part not.
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    In technical applications, the weight is usually preferred over the moles because it is measured more easily. However, in thermodynamics the number of particles is the more fundamental quantity. It has already been pointed out that in chemical reaction equations it is the number of atoms which must be balanced, not the involved weights.

    But there are also other thermodynamic properties which depend on the number of particles rather than their weight. In solutions, for example, these properties even have their own name: colligative properties. If you dissolve a substance in water, the resulting solution has different properties than the pure water. One of these changes is the depression of the freezing point: the solution freezes at lower temperature than the pure water. The degree of depression depends on the number of dissolved particles, not on any of their other properties (such as their weight). If you dissolve the same number of other particles (maybe heavier molecules), you get the same freezing point depression. (Some caveats with respect to ideality etc apply.)

    You can use this to determine the molar mass of a soluble substance: You dissolve a known weight of the substance in water and observe the resulting depression of the freezing point. This tells you into how many particles the substance been dissolved, and together with the known weight this gives you the molar mass.

    If you dissolve a substance with known molar mass and find twice the depression you expected, you know that the substance dissolves not only into individual molecules but that each molecule also dissociates into two smaller units, so that you ended up with twice as many individual particles in the solution as expected.

    Other colligative properties include the elevation of the boiling point or the osmotic pressure.


    There are other thermodynamic relationships which become obvious when the relevant quantities are expressed per mole rather than per kilogram. For example, if you look at the molar heat capacities of the noble gases, expressed in Joule per Kelvin and mol, you will find the same value for all of them: Cp = 20.786 J/(K mol). This is because in thermal equilibrium and at the same temperature, each particle - regardless of its mass - assumes the same thermal energy on average, not each kilogram. If you look at the specific heat capacities of the noble gases, expressed in Joule per Kelvin and kilogram, no obvious pattern appears.

    Regards,
    Thomas

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    One of Einstein's papers from 1905 was on diffusion, and determining Avogadro's number....he was pretty good, and it took a while to find it from Avogadro original conjecture.. ( he died not knowing its value). pete

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    Quote Originally Posted by Robert Tulip View Post
    Thanks all, makes good sense. My interest in molar chemistry comes from climate physics, where weight is a more accessible framework to a broad audience than stoichiometry.

    I used an excel converter linked here to answer my specific question on iron and carbon. It explains that an iron atom weighs 4.65 carbon atoms. That would mean the 100K molar ratio mentioned in the article I read means addition of an iron atom could theoretically enable reactions involving up to 465,000 times its weight in carbon in an iron limited sea location.

    Use of molar ratios is an issue I have come across several times in technical papers, where I have had to convert to weight to make sense of the argument. I have also seen scientists get this conversion wrong, which seems an easy mistake to make. It makes me wonder if when chemists are writing for a broader audience they should more often convert their atomic ratio calculations into weight measures to make them more accessible.
    I remember when decimal currency came in here in the UK.

    There were all kinds of currency converter cards etc. People would convert the decimal prices into pounds, shillings and pence, work out their change, then convert the pounds shillings and pence back to decimal. Sounds like you're doing a similar thing here

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    Quote Originally Posted by kzb View Post
    I remember when decimal currency came in here in the UK.

    There were all kinds of currency converter cards etc. People would convert the decimal prices into pounds, shillings and pence, work out their change, then convert the pounds shillings and pence back to decimal. Sounds like you're doing a similar thing here
    No, not at all. I raised the example of how iron added to the ocean will increase photosynthesis. I found papers that explain the chemistry in molar terms of the ratio of iron and carbon, and have converted these to weight. No back conversion involved.

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    Quote Originally Posted by grapes View Post
    Which begs the question, are atoms particles?
    Only in special circumstances and with noble gases. Most of the time the fundamental particle of a mole will be a molecule. Like Jeff said, a mole is a number of particles.

    But "the particle" is the smallest unit of the substance in question. If you break up a compound into its atoms that doesn't count.

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    If I recall, "mole" is short for "gram-molecule," because the gram is the mass unit involved. One could use other units of mass (e.g., kilogram-molecule, pound-molecule) if that somehow turns out to be more convenient.

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    Quote Originally Posted by kzb View Post
    Only in special circumstances and with noble gases. Most of the time the fundamental particle of a mole will be a molecule. Like Jeff said, a mole is a number of particles.

    But "the particle" is the smallest unit of the substance in question. If you break up a compound into its atoms that doesn't count.
    So, if we allow "particles," we can allow "hedgehogs"?

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    Quote Originally Posted by DonM435 View Post
    If I recall, "mole" is short for "gram-molecule," because the gram is the mass unit involved. One could use other units of mass (e.g., kilogram-molecule, pound-molecule) if that somehow turns out to be more convenient.
    Don. Mole was shorthand for "gram -molecular weight" in the fifties. It was a mass in grams equal to the molecular mass in AMUs. Should have been "gram-molecular mass" in a better, more accurate chemistry book.
    If you multiply the mass of the particle (atom or molecule) by Avogadro's Number, it converts from AMU's or Daltons, to grams.
    Avogadro had the right idea, there ought to be a certain mass of a substance that would supply the same count of particles.....like a dozen, except the particles were exceedingly tiny, so the count had to be enormous (Av. #), so that you could measure it out on a balance.
    Were it not for the intervention of the great Swedish chemist, Berzelius, the idea may have languished. At a conference he chastised chemists and physicists who were bickering with Avogadro, who had used the concept successfully to explain the small whole number ratio of combining volumes of gases under STP.
    Let no good idea go unridiculed. pete

    corollary....if you divide 1 gram by Av. #, you get one AMU or Dalton.
    Last edited by trinitree88; 2018-Jan-27 at 12:05 AM. Reason: amend

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    Quote Originally Posted by grapes View Post
    So, if we allow "particles," we can allow "hedgehogs"?
    Interesting argument.

    A large number of hedgehogs does not constitute a substance. It's just a number of hedgehogs.

    On the other hand we have "yeast" which we treat as a substance and I suppose you could have a mole of yeast cells.

    BUT, fundamentally, the mole is all about reaction ratios. To use it for collections of entities which have infinitely variable reaction ratios is to use it outside of its intended purpose.

    So on balance I don't think you can have a mole of hedgehogs or yeast. You can have Avogadro's number of hedgehogs fair enough, but it's not a mole.

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    A glass of water...~ 300 ml., contains ~ 15 moles of water. If you could put your initials, miraculously on each one, and dump it in the ocean, river, lake, pond or ground to evaporate from.....and wait a few thousand years, you'd find a few of those molecules in every glass of water on earth, salty or not. There's a bunch of molecules in a mole, more than there's glasses of water on Earth.

    pete

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    Quote Originally Posted by kzb View Post
    Interesting argument.

    A large number of hedgehogs does not constitute a substance. It's just a number of hedgehogs.

    On the other hand we have "yeast" which we treat as a substance and I suppose you could have a mole of yeast cells.

    BUT, fundamentally, the mole is all about reaction ratios. To use it for collections of entities which have infinitely variable reaction ratios is to use it outside of its intended purpose.

    So on balance I don't think you can have a mole of hedgehogs or yeast. You can have Avogadro's number of hedgehogs fair enough, but it's not a mole.
    It wasn't my argument, it was a question.

    But, it's interesting.
    https://www.bipm.org/utils/common/pd...chure_8_en.pdf

    The definition of the metric base unit of mole seems to include
    When the mole is used, the elementary entities must be specified and may be atoms, molecules, ions, electrons, other particles, or specified groups of such particles.
    That "specified groups of such particles" is kinda open-ended

    Chemically, we're already using entities that are not identical, like isotopes in a mole of a substance.

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    Quote Originally Posted by Swift View Post
    Whack A Mole

  30. #30
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    Leaving aside semantic sophistry, there are quite a few quantities when normalized per mole are similar, but those per unit mass aren’t

    Information about American English usage here and here. Floating point issues? Please read this before posting.

    How do things fly? This explains it all.

    Actually they can't: "Heavier-than-air flying machines are impossible." - Lord Kelvin, president, Royal Society, 1895.



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