# Talk:Mole (unit)

## Duh?

You could read the entire article and still not understand what a mole is. Like, an encyclopedia is for people who don't know nothing to start with. Please, can we start with the definition I was taught in first year chemistry: "A mole is the atomic weight of an atom or molecule expressed in grams. Thus if the atomic weight of carbon is 12, then one mole of carbon is 12 grams. Avogadro's number is the number of atoms or molecules in a mole. Thus if carbon has atomic weight 12, one mole of carbon (12 grams) contains Avogadro's number of atoms." — Preceding unsigned comment added by 213.207.137.134 (talk) 18:49, 21 October 2015 (UTC)

## Relevance

Avagadro's Constant(The number of atoms/molecules in a mole) should be listed on this page for ease of refrence. As I just did. Also, since it's impossible to count the number of atoms/molecules in something with anything representing a reasonable accuracy, I was rather confused when someone claimed that a mole (The number of atoms so that an object weighs it's atomic weight in grams) is not based off of how many atoms there is?

Anyway, you need to know how many atoms there are if you're trying to do Thermo Dynamics equations and need to convert from molecules to moles so you can apply the Ideal Gas Law. 63.228.160.115 (talk) —Preceding undated comment added 19:57, 30 August 2009 (UTC).

## Unit?

Is a mole really a unit? In high school chemistry, I was taught that 1 mole is an amount with a value of 6.022x10^23.SamWhitey (talk) 20:31, 4 November 2008 (UTC)

Yes, it is one of the base units of the international system of units. See mole (unit). Its current definition is actually independent of the exact value of Avogadro's number (the 6.022x10^23). --Itub (talk) 15:14, 8 November 2008 (UTC)
It is a dimensionless unit.--86.125.156.119 (talk) 12:32, 31 January 2015 (UTC)

## misconception?

'It is a common misconception that the mole is defined in terms of [...] Avogadro's number'

...Is it? Look at the definition: 'The mole is defined as the amount of substance of a system that contains as many "elemental entities" (e.g., atoms, molecules, ions, electrons) as there are atoms in 12 g of carbon-12'.

How many atoms are there in 12 g of carbon-12? Avogadro's number!TAB (talk) 15:56, 9 July 2009 (UTC)

-I am in complete agreement. The mole can not be considered without discussing its connection with Avogadro's number. The two go hand in hand. The mole is defined in terms of Avogadro's number. —Preceding unsigned comment added by Isaac B Wagner (talkcontribs) 20:26, 18 October 2009 (UTC)

I agree. A mole of something 'is' defining the quantity numerically - it is akin to stating a dozen of something, only the implied number is bigger. Simply stating this is wrong without any justification is plainly POV pushing which I will tag as such. CrispMuncher (talk) 22:27, 12 November 2009 (UTC)
I find it strange that people are somehow unable to see this thread and accuse me of blanket tagging without providing a rationale. It it right here. It is not a full argument to be sure but I note that as yet no-one has yet made any assertion anywhere that the current article is correct. As such currently the consensus shown here is that there is a problem, because no-one has bothered to challenge this position. Indeed, the current article acknowledges the existence of sources that adopt this position but provided no countersources so casually discarding the assertion. That is fundamentally POV however you want to dress it up. The following sentence (..the mole is defined in terms of the Avogadro constant, rather than the other way around..) is IMHO a non-sequiter. The fact that the mole and Avogadro's constant are defined the way they are is immaterial to the central issue that the mole is a quantity representing a count of a number of discrete objects. Those that wish to brush this aside as purely an amount of substance would do well to explain the ambiguity in something like "a mole of oxygen". If a mole is a clear amount how can there be such ambiguity? CrispMuncher (talk) 19:35, 13 November 2009 (UTC)
The mole is not defined in terms of Avogadro's number. Look at the SI definition. Do you see Avogadro's number mentioned there? No. The key point that people miss is that you can use moles as amount of substance without having the slightest idea of the value of Avogadro's number. The reason for defining it without reference to Avogadro's number is that it allows for more precise measurements. For example, since the molar mass of fluorine is known more accurately than Avogadro's number, in principle you can measure one mole of fluorine more accurately by weighing it than by counting the atoms. I'm sure plenty of articles about this misconception have been published in the Journal of Chemical Education, but I don't have access right now. --Itub (talk) 20:47, 13 November 2009 (UTC)
So this is where the "current discussion" is. Hmm... seems Itub has provided a pretty good answer. As stated elsewhere on this page, the mole was defined and used as an amount long before Avogadro's number showed up. Chemists don't count atoms or molecules - they either go by mass (atomic weight) or molar volume for gases. As for your comparison to "dozen" (a common comparison in introductory chemistry) it's a "bit" off. A dozen is 12. A mole is equal to 6.0221415×1023 ... but consider the uncertainty implied by that number with seven decimal places ... what about the other sixteen? That's 6.0221415×1023 +/- 10 quintillion or so. So how is that a "count"? Bit of "ambiguity" there.
Now, if CrispMuncher cares to suggest a change in wording, we will certainly consider it. But to stuff an ugly POV tag in an article based on a misunderstanding is a bit much. I'm going to remove the tag while waiting for your suggested changes. Vsmith (talk) 00:54, 14 November 2009 (UTC)
Anyone who says the mole is a count should start by counting from from one to sixty thousand million million million… Physchim62 (talk) 08:46, 14 November 2009 (UTC)

## Circular argument

Consider the measurement of one mole of silicon.

This seems a little circular — the argument that the mole is useful for measuring out a mole of silicon lacks a certain something; it certainly doesn't illustrate why it should be a unit. Similarly, if you wanted a million atoms of silicon, it would be useful to know the million, but that doesn't make "million" a unit. The only reason the mole is useful in that example is because atomic weights are expressed in conventional units that are defined coherently with the mole, rather than in yoctograms. That's a rather weak argument, especially since no reason is given why the table doesn't just list silicon atoms as weighing 46.6371yg (on average). As far as I know, the main reason is in fact convention rather than anything fundamental.

Personally, I suspect that the only real argument for the mole being a unit would be a statement to the effect that the realms of bulk matter and individual atoms are qualitatively different, such that the dimensions of one don't apply to the other.--Sabik (talk) 16:23, 5 January 2009 (UTC)

I don't think the argument is circular at all. Are you saying that the mile is only useful for measuring miles of things? Or the pound is only useful for measuring pounds of things? The point of the paragraph is to show that the measurement of amount of substance does not involve the use of the Avogadro constant. On the other hand, the measurement of the mass of atoms in yoctograms does require a knowledge of the Avogadro constant: relative atomic masses are known more accurately than masses of atoms in fractions of a kilogram: THAT is the fundamental reason why the table doesn't list the masses of silicon atoms in yoctograms. Physchim62 (talk) 22:25, 5 January 2009 (UTC)
"Are you saying that the mile is only useful for measuring miles of things?" No that's not what he's saying. It does seem pretty circular to me, too. The intro talks about the misconception of the link with Avogadro's Number and claims "[i]t is not necessary to know the number of atoms or molecules which are present in order to use the mole as a unit of measurement, and indeed the first measurements of amount of substance predate modern atomic theory and any measurements of atomic weight." Okay... so it's purportedly independent of number of atoms and the concept of an atomic weight. Later, with the silion 'example': "the convenient method is by weighing. By consulting published tables, it can easily be found that the ATOMIC WEIGHT of silicon is 28.0855.". Published tables? What a cop out. We're looking up from tables the atomic weight, but the definition is supposed to be independent of all atomic theory notions? Please. How did we find out that the atomic weight of silicon is 28.0855 without any notion of Avogadro's number or atomic theory? The mole is defined as amount of 'stuff' in 12 g of C-12. Right. Now what? How do you translate that to getting the mass of equivalent amount of 'stuff' for silicon? I could see an argument involving masses of required proportions in empirical reactions with known molecular formulae to determine the mass equivalent of 1 mole of something else given the presence of 1 mole of C-12 in the reactants. I'm not sure that would be fundamentally possible for all elements in the periodic table, and I think reaction equilibria would greatly complicate precision at best and practicability at worst. Regardless, the current silicon example is terrible and circular. Essentially it says you don't need reference to atomic weights/atomic theory/Avogrado's number to use the mole of silicon. Just look up the atomic weight in a table (gee, I wonder how that was determined). It skirts the entire problem it was trying to demonstrate the solution to, i.e. how to use the definition of a mole to determine how much is 1 mole of something else. —Preceding unsigned comment added by 64.85.36.210 (talk) 01:16, 7 January 2009 (UTC)
For ratios of reactants, you don't care about the units in the published table at all, as long as they're consistent. In that situation, you're interested in the ratio of atomic masses of your reactants, which is dimensionless. The only situation where it'd be interesting would be if someone told you the other beaker contains one mole of the other reactant, but that's no more a reason for a separate SI unit than someone telling you that the other beaker contains a pound of the other reactant, or an ancient Greek mina.
As for whether the mole is useful for anything else, I'm actually not commenting on that - I'm commenting that the example given doesn't show such other use. Sabik (talk) 18:43, 14 January 2009 (UTC)
What is the mass of one atom of carbon-12 in SI units? Explain how you arrive at your answer without using the concept of amount of substance which, in SI units, is measured in moles. Physchim62 (talk) 19:19, 14 January 2009 (UTC)
So, essentially, a mole is useful because, in SI units, amount of substance is measured in moles. Right. (FWIW, numbers, in scientific notation when necessary, can also be used to describe amount of substance. For instance, one might have 6 apples. One does not usually talk about having 10 yoctomoles of apples, even though that would be the equivalent.)
I'm not arguing here that a mole shouldn't be a unit; I'm merely observing that the example used to motivate it does no such thing, since it is circular. The solution is to replace the example with one that does actually motivate the use of the unit. Sabik (talk) 16:02, 27 January 2009 (UTC)
Well atomic weights certainly aren't determined by "weighing" atoms! Modern methods rely on Penning traps: by measuring the frequency of the radiation emitted by silicon ions and carbon ions in a Penning trap, you can tell that a silicon atom is 28/12 times heavier than a carbon atom. Older methods usually relied on measuring the change in mass of a sample when it underwent a chemical reaction. This gives a large set of ratios, which can be converted into simpler numbers by fixing the atomic weight of one element: originally hydrogen was chosen as the standard, as it is the lightest element, but oxygen proved a more practical standard (set as O = 16) as most elements form oxides. There are big problems with stoichiometry – when you first do the measurement, you can't be sure of the formula of the oxide – but there are other methods which can be used as a check, such as the Dulong–Petit law for metals or vapour densities for gases and volatile liquids.
None of these methods depend on the fact that atoms exist (well, the Penning trap sort of assumes that they do, but that is a modern innovation), much less the mass of an atom in kilograms. The first measurements of atomic weights were published by Dalton in 1808; the first estimate of the size of a molecule was made by Loschmidt in 1865; the Avogadro constant wasn't even conceived of, let alone measured, until 1909. For one hundred years, atomic weights were based on stoichiometric ratios of reactant masses: even today, the relative masses of atoms are known to some four orders of magnitude more precisely than their absolute masses (in kilograms). Physchim62 (talk) 22:21, 7 January 2009 (UTC)
Four orders of magnitude more precision does sound useful :-) Sabik (talk) 18:43, 14 January 2009 (UTC)
Speaking of circular, I would argue that "weighing" should probably be defined as determining the weight of, so how is that not what you're doing? I agree that it's sometimes more convenient to determine them in non-SI units (for instance, oxygen atom weights) but it still just amounts to weighing with a fancy scale. —Preceding unsigned comment added by 205.175.113.173 (talk) 17:51, 28 June 2009 (UTC)

The real circular argument comes if you try to define a mole as 6×1023 atoms or molecules – how do you count them? Imagine you have a one petaflop computer which can count an atom with every floating-point operation: it would still take you 600 million seconds (200 years) to make a measurement that a high-school student can do in a few minutes!
There is a way round the circular argument – fortunately, because the mole might be redefined in this way in the future – but you have to give up the exact relation between molar mass and atomic weight. If, under the new definition, you also admit that 6.022 141 79×1023 atoms of carbon-12 weigh approximately 12 grams (to within 50 parts per billion), things turn out OK again, at least to within the accuracy needed (and possible) in chemistry. The current measurement uncertainty in the value of the Avogadro constant would become an uncertainty in the value of the molar mass constant. But that is for the future; the current definition is as described in the article. Physchim62 (talk) 23:26, 7 January 2009 (UTC)

That's an engineering detail... essentially, it amounts to a claim that the realms of bulk matter and individual atoms are so different that it amounts to a qualitative difference. In practice, it would hardly be the first situation where the most convenient method of counting things is to weigh them in bulk. People do that with nuts and bolts. In any case, it doesn't address my complaint about the example; the example claims to show why the mole should be a unit, but it does no such thing. If you wish to replace the example with a bare statement saying that the mole is a unit because it is currently defined to be one, that would be fine, if somewhat unsatisfying. Sabik (talk) 18:43, 14 January 2009 (UTC)
It's not an engineering detail at all, it goes right to the heart of how people actually measure things. Several laboratories in the world are trying to calculate how many silicon atoms there are in sphere of silicon which weighs one kilogram: for the moment, they cannot get their result to be as accurate as other methods of measuring the Avogadro constant but they're still trying. On the other hand, the mole is used daily around the world as a measure of amount of substance without any great problems. The section states certain objections to the mole being a base unit of the SI system, which might be better treated at amount of substance except that the attacks tend to come on this article. Maybe you're not willing to admit that you can't know the mass of an atom (in everyday units) without some measure of amount of substance… Physchim62 (talk) 19:35, 14 January 2009 (UTC)
Once again, I'm mostly complaining about the example (which is circular) rather than the use of mole as a unit. Fix the example.
As for the practicalities of measurement and counting, compare with the field of astronomy — which also has many large numbers, but nobody talks about a decimole of stars (I believe that's how many there are?). While specialised units are defined, they are admitted to be specialised, for within-astronomy use only.
For the most part, though, fix the example so that it actually motivates the use of the unit… Sabik (talk) 16:02, 27 January 2009 (UTC)

According to Pieter G. van Dokkum & Charlie Conroy the number of stars in observable Universe is 0.5 mol. Neeme Vaino (talk) 04:16, 23 July 2014 (UTC)

## Prasath Santhakumaran

Some prankster added this to the article:

"Also, it was proven that Avogrado's number was stolen off another brilliant genius in Prasath Santhakumaran. In his time of death Avogrado said that it was true that he did steal the number off Prasath and that this shouldn't be told to anyone."

This is unsourced and, as far as I know, a joke in poor taste and tantamount to vandalism on its face. I am removing it forthwith. Its author deserves some sort of reprimand. Trujaman (talk) 21:28, 7 January 2009 (UTC)

## Alternate mole definitions

Apparently someone saw fit to undo the changes I made a few days ago to add a section to this article about other definitions of a mole (e.g., a kg-mol). He/she claimed that these changes were inappropriate as "These are not definitions of the mole." I must disagree: I have had numerous students confused by the use of kg-mol and lb-mol in textbooks, primarily because they too thought there was only one way to define a mole. Confusingly or otherwise, these definitions persist.

The term kg-mol is quite common among chemical engineers, and the lb-mol is even more common among chemical engineers from the USA. Software designed to assist in chemical plant design, such as PRO/II and ASPEN, all include these units. While these alternate definitions are rarely if ever seen outside the field of chemical engineering, their existence warrants at least being mentioned in an article dedicated to the definition of a mole.

Examples:

• M. S. Peters, K. Timmerhaus, and R. E. West, Plant Design and Economics for Chemical Engineers, Fifth Edition. McGraw Hill (2002-2003). [Uses kg-mol, written as such.]
• J. M. Douglas, Conceptual Design of Chemical Processes. Boston: McGraw Hill (1988). [Uses lb-mol, written as mol, throughout.]

Kaiserkarl13 (talk) 21:46, 20 January 2009 (UTC)

The pound-mole (also written without the hyphen) is a separate unit, although the link currently redirects here. The "kilogram-mole" is just an obsolescent name for the kilomole. Both of them are units of measurement of amount of substance that can be found in some US texts, but neither of them is the same unit as the mole. The confusion only arises if you pretend that they are. Physchim62 (talk) 22:12, 20 January 2009 (UTC)
The fact that pound-moles and kilogram-moles are separate units but are still called moles was my point! The fact that kilomoles and kilogram-moles are the same is a result of the fact that kilo(grams) and (kilograms) are the same, but the kg-mol is in no way an obsolete unit---it just happens to have the same meaning as a kmol and takes more characters to type. The purpose of the added text is to point out that there is more than one definition of a mole---with the standard SI unit of a mole being defined in terms of the gram. There's nothing pretend about it, though---the units kg-mol and lb-mol still occur, and they are no less valid as definitions of a mole. They just aren't the ones associated with the SI definition of "the mole." Kaiserkarl13 (talk) 23:00, 20 January 2009 (UTC)
You'd never say that pounds are another way of defining grams, so why do you say that the pound-mole is another way of defining the mole? It's another way of defining a unit of amount of substance, yes, but there is no law that says that all units of amount of substance have to be called "moles": in fact, there is a good pedagogical case for pointing out the exact opposite! Physchim62 (talk) 23:14, 20 January 2009 (UTC)
I'm not contending that the specific unit referred to as the mole (i.e., the SI unit defined as the number of entities in 12 g of 12-C) is the same unit as a kilogram-mole---obviously it isn't. The point of contention is that other units of amount of substance are invariably also called moles (with some modifier), perhaps merely due to lack of creativity. An excellent example is the unfortunate use of pounds in the English Imperial System to refer to both units of force and mass (these are, of course, different but related units). I think this point of contention could be settled by adding a section titled something like, "Units related to the mole" in which other units of amount of substance are defined. This would have the advantage of clarifying that "the mole" (unit) and "a mole" (generic term for amount of substance, usually but not always referring to the SI unit) can be different things. This would also solve the problem of pound-mole and kilogram-mole currently redirecting to Mole_(unit), which doesn't mention either of these other units. I may write these articles at a later date when I have the time. Kaiserkarl13 (talk) 21:35, 21 January 2009 (UTC)
Possibly this should be "Other units called mole" or something on those lines? That would make it clear that the lb-mole and kg-mole (both of which redirect here) are not actually the mole but rather something else with a confusingly similar name… I think I'll go add that. —Preceding unsigned comment added by Sabik (talkcontribs) 16:15, 27 January 2009 (UTC)

I have re-written the section, trying to be brief and non-contentious, while giving a source and giving the style of an encylopedia. I hope this helps. Chemical Engineer (talk) 17:52, 5 February 2009 (UTC)

## Chronologically?

"(also, anachronistically, known as "Avogadro's number")" Wha-Huh? "Anachonistically", chronologically out of order? Does that mean that yesterday it was correct, but today it is not correct? Maybe it means that it will be correct in 1.66 years. I'm not asking if it is correct to call Avogadro's constant, Avogadro's number. I'm asking what time has to do with it. —Preceding unsigned comment added by 207.5.226.78 (talk) 12:16, 7 April 2009 (UTC)

When it was first measured, it was assumend to be a pure number because people didn't stop to think that its value depends on the units used to measure amount of substance. Now it is generally accepted that amount of substance is a separate dimension in any practical system of macroscopic units of measurement, and so it is proper to call it the Avogadro constant to recognise the fact that it is not a pure number, but a physical constant with a unit (mol–1). Physchim62 (talk) 13:25, 7 April 2009 (UTC)

## Etymology

The name is assumed to be derived from the word Molekül (molecule). The first usage in English dates from 1897, in a work translated from German.

That seems at best incomplete. The Oxford English Dictionary says

[< French molécule (1674) < post-classical Latin molecula (P. Gassendi Syntagmatis Philosophici (a1655) II. §1. III. vi, in Opera Omnia (1658) I. 271/1) < classical Latin m{omac}l{emac}s mass (see MOLE n.2) + -cula -CULE suffix.

and then of the -CULE suffix says:

suffix, corresp. to F. -cule, ad. L. -culus, -cula, -culum, dim. suffix of all three genders: see -CULUS. In living words, the suffix underwent various phonetic changes in becoming French; e.g. articulus, orteil; auricula, oreille; cuniculus, conil; masculus, masle, mâle; but it remained as -cle after persisting consonants, as in avunculus, oncle; cooperculum, couvercle. After the latter, some words of learned origin were fashioned in -cle; e.g. article; but in modern times the L. ending has been usually adapted in F. as -cule, as corcule, cornicule, corpuscule. In English, both endings -cle and -cule are found, as corpuscle, corpuscule, crepuscle, crepuscule, animalcule, formerly also animalcle, floscule, versicle, etc. The L. endings -culus, -culum are sometimes retained unchanged: see -CULUS. The ending -cule, with connecting vowel i, is sometimes employed, after L. analogies, to form contemptuous diminutives, as poeticule: cf. criticule.

The term "dim. suffix" means "diminutive suffix", so things bearing that suffix are small. So "mole" means "mass" and "-cule" makes it diminutive, so "molecule" is a small mass.

Thus "mole" did not come from "molecule", but rather "molecule" came from "mole" by the addition of a suffix. Michael Hardy (talk) 19:49, 7 April 2009 (UTC)

PS: "Assumed" is very vague! Assumed by whom??? Some anonymous Wikipedia editor, maybe?? Michael Hardy (talk) 19:50, 7 April 2009 (UTC)

The word "mole" certainly doesn't come directly from the Latin mole! You might like to read some of the references in the article before engaging in such vain speculation. Physchim62 (talk) 18:36, 28 June 2009 (UTC)

Michael Hardy: you seem to be quoting from the OED entry for molecule, n., not that for mole, n.8, which is why you have reached the wrong conclusion. The OED tells us that mole (English; the unit) comes from Mol (German), which comes from Molekul (German) (although that last step is not explicitly stated in the Ostwald 1893 book). There's no point in tracking it back any further, since the origin of Molekul is obviously, like that of molecule, ultimately from Latin.
I do agree with your postscript, that "assumed" is vague. The assumption seems to have been made by the OED, so we should say that. --Heron (talk) 18:18, 30 January 2010 (UTC)

## National Mole Day

The radio advised me this morning that today is National Mole Day. Who’d have thought?

The mole is a wondrous thing, and easy to understand – as long as you realize it is just a number:

• If you have a dozen eggs, you have 12 eggs.
• If you have a baker’s dozen of buns, you have 13 buns.
• If you live a score of years, you live 20 years.
• If you have a Groß (or gross) of hobbits you have a dozen-dozen, or 144 hobbits.
• If you have a ream of paper, you have 500 sheets or paper
• If you have a bale of paper you have 5000 sheets of paper
• If you have a Maß (or ‘great gross’) of tacks, you have a gross-gross (a dozen-dozen-dozen-dozen) or 1728 tacks.
• If you have a Mole of atoms, you have 6.02 X 1023 atoms
• If you have a nonillion freckles, you have 1030 (or perhaps 1054, depending upon where you live) freckles
• If you have googol of edits, you have 10100 edits.

What could be simpler?

But to the question: Is "National Mole Day" encyclopedic enough to be included in a Wikipedia article?

Skål - 130.20.3.152 (talk) 16:37, 23 October 2009 (UTC)

Of course, we even have an article about it: Mole Day. --Itub (talk) 14:57, 26 October 2009 (UTC)

## POV Section

"The mole as a unit" is POV (and has a disputable match between content and heading):

The most blatant is the phrasing "The second misconception" referring to something which could be taken both as true and as a matter of perspective. This is the worse as the author borders on straw-manning by equating the originally stated criticism "the mole is simply a shorthand way of referring to a large number" with "that the mole is simply a counting aid", formulations that are not equivalent (albeit related). Notably, the "second" also implies that the previously discussed item is a misconception---which is equally disputable.

Looking more in detail on:

Claim 1: I must second that the mole is dimensionless and it lacks the significance of e.g. the meter. (Notwithstanding that it may be justified as a unit.) It is better compared to % (per cent): Handy and useful, but not truly significant. Further, it is redundant when the mass and composition is known. The argumentation against this is partially specious. The number of particles of the involved elements can easily be computed when mass and composition are known (and, conversely, knowing the number of particles and composition, the mass is easy to find). Without knowing the composition, OTOH, the concept of moles would typically be useless. In effect, moles and the associated calculations are a handy help, but are still redundant. That the number of moles can be found without knowledge of mass is irrelevant, because the number of moles and mass can always be found from each other when composition is known.

(There are some special cases where the equivalency fails, e.g. in relativistic conditions; however, the calculations can be adapted and mathematical equivalency of information is preserved. Further, while the ideal gas law may give some physical characteristics based on number of particles alone, these are of limited value, and more advanced models use characteristics other than particles alone, which require that composition, or other characteristics which lead to mass or possibly a mass surrogate, is known.)

Claim 2: Whether the mole is defined in terms of Avogadro's constant, or the other way around, is irrelevant. It is a short-hand for a large and arbitrary number. The other gains from using moles could be found by using the atomic weight directly. (With the exception of some handy approximative jumps from weight to number of moles when doing head calculations.)

In short: Apart from convenience, the mole has no advantage over the number of particles (as a "vanilla" number, just like 10 % has no advantage over 0.1). Apart from convenience, very little is gained from using moles instead of mass. Michael Eriksson (talk) 13:29, 23 January 2010 (UTC)

Agree that, while there may be some interesting points, it is presented in the form of an argument, not encyclopedic. Noloop (talk) 21:00, 9 April 2010 (UTC)

## 6.0221415 × 1023 / 0.012?

The kilogram is the mass of exactly (6.0221415×10230.012) unbound carbon-12 atoms at rest and in their ground state.

This seems a strange way to write it. Is it really planned that the official definition will be written like this, rather than simply as 5.01845125 × 1025? Any idea why? — Smjg (talk) 23:06, 17 May 2011 (UTC)

If you follow up the links in New SI definitions, you can see the full story. Martinvl (talk) 06:17, 18 May 2011 (UTC)

~1.660539 ymol of quark conversions were used to produce a proton. see W_boson#Weak_nuclear_force — Preceding unsigned comment added by 193.199.19.150 (talk) 06:43, 29 March 2012 (UTC)

## Redefinition

Definitions: Section 4 and 4.1 considering a new definition of the kilogram does not belong in this article on the mole, amount of substance - the section should be removed or rather moved to the kilogram page. On the other hand it is proper to consider redefining the kilomole (kmol) as a basic SI unit rather than the mole. I see no difficulty in doing this and retaining all the definitions of Avogadro's number (and others) as they are at present. What it will do is make the basic SI definitions a coherent* set rather than include the odd man out, the mole.

• I note that Wikipedia doesn't contain this notion in its list of coherence definitions - I will endeavour to elucidate soon.

Scanrod (talk) 11:23, 19 June 2010 (UTC)

In what sense is the mole the "odd man out" at the minute? Physchim62 (talk) 11:40, 19 June 2010 (UTC)

I have hidden this.

A decision on this proposal is expected by the CGPM in October 2011.

It cannot be true since October 2011 is now past. So was a decision made? If so, what was the outcome? If not, what are they saying now? JIMp talk·cont 01:08, 8 December 2011 (UTC)

I posted a link to the CGPM 2011: see resolution 1. I think we can now delete the section about the kilogram from the mole article. --Hroðulf (or Hrothulf) (Talk) 17:05, 12 December 2011 (UTC)
Kilogram section now removed per Scanrod's request. --Hroðulf (or Hrothulf) (Talk) 17:10, 12 December 2011 (UTC)

## Consistency between articles

In various articles around Wikipedia that make use of Avogadro Number (such as Avogadro's Number and this article, and many of the articles that use Avogadro's Number in calculations), the actual number used differs. And even more unfortunately, most of these articles have cites that "prove" their number correct. Is there any way we can enforce consistent use of the number, or make reference to one "superior" cite that provides the most up-to-date and accurate number? --Nick2253 (talk) 03:38, 6 November 2012 (UTC)

Examples please? Martinvl (talk) 04:35, 6 November 2012 (UTC)

A citation only proves that someone (which is being cited) made some statement, not the statement itself. In this case, there have been quite a few experimental measurements of the Avogrado's Number over the years. One needs to look at the year of the citation to see which is more recent. I agree that it would be good to have a tool that automatically updated the value (and the information source) in every page it appears.
177.138.210.163 (talk) 02:00, 25 March 2017 (UTC)

## Confusion

Is there sometimes the mole confused with the molar mass in some sources by abuse of language?--86.125.160.200 (talk) 15:59, 10 January 2015 (UTC)

Kindly see typo mistake in first part "For example, the chemical equation 2 H2 + O2 → 2 H2O implies that 2 mol of dihydrogen (H2) and 1 mol of dioxygen (O2) react to form 2 mol of water (H2O)." It should be 2H2+ 0 and not O2. M S DIVEKAR (talk) 17:58, 5 September 2015 (UTC)

The equation is balanced as written. Vsmith (talk) 18:34, 5 September 2015 (UTC)

## Deletion of molar mass - atomic mass ratio demonstration

I found these sentences, and I only reformatted them:

```We can derive the relationship between a mole of a substance and the mass of one atom/molecule of the substance using the definition of the mole stated above and the relationship between a gram and an amu (there are Avogadro's-number of amu in a gram of a pure substance) like so:

Mass of 1 hydrogen atom = 1.008 amu.

Mass of 1 mole of hydrogen = 1.008 ⋅ [6.022 141 29(27) ⋅ 1023] amu = 1.008 g.

Mass of 1 hydrogen atom = 1.008/6.022 141 29(27) ⋅ 1023 g.

We now compare the mass in grams of 1 hydrogen atom with the mass in grams of 1 mole of hydrogen. By dividing the mass of 1 mole of hydrogen by the mass of 1 hydrogen atom we'll obtain the scale-factor between the mass of 1 hydrogen atom and the mass of 1 mole of hydrogen (we will use the numbers that are 2 lines above this one). Since dividing by a fraction is the same as multiplying by that fraction's reciprocal, we can see that the scale-factor is:

mass of 1 mole/mass of 1 atom = 1.008 g ⋅ 6.022 141 29(27) ⋅ 1023/1.008 g = 6.022 141 29(27) ⋅ 1023 = NA

This scale-factor holds true for all pure substances. For example, a mole of water is Avogadro's-number times the mass of one water molecule (H2O).
```

But I think it is a not very useful demonstration, that overload the page and make it difficult to read. I suggest to remove it. --Fornaeffe (talk) 11:29, 7 February 2016 (UTC)

## name acquire

two basic units are named after scientists; A[mpere] and Ke[lvin], so what if this ambiguous name labeled after Jons Jacob the Be[rzel]. and mmol and Kmol can be mBe and KBe. and a katal would be Be/sec
Tabascofernandez (talk) 00:19, 15 July 2017 (UTC)