Talk:Mole (unit)

intro prompted by comments on the forumsGareth Leng 12:46, 28 November 2007 (CST)

Re: "(clear language intro??)" No question Gareth, it does read more clearly.--Thomas Simmons 14:41, 28 November 2007 (CST)

not a textbook

I rewrote the introduction, and removed the worked-out example from the article. The example added unnecessary bulk to the article. Anthony Argyriou 18:16, 28 November 2007 (CST)

Can't agree with the rationale here for the removal of the 'textbook' style text. The assertion that this is not a textbook to make sweeping deletions, and I do mean large patches of text, is to be discussed here, not simply a decision to be reached by a single author. I have taught chemistry at the high school and the undergraduate level and can say from experience that as a basic concept in chemistry it is imperative that it be comprehended. Writing for more than one level of understanding is very much a function of a viable encyclopaedia. I have had a number of students here in New Zealand and in the States as well as teachers at those levels read this for their comments and the extended version definitely reaches a broader range of levels of understanding. The amended intro has received varied comments which indicates to me that it is reaching a higher level of understanding--basically those who already know what a mole is. It is the extended text that is giving vital concrete information so that the concept is generalised by the reader and that is needed to support the technical level of the introduction. In other words, if you know what a mole is, why consult an encyclopaedia? Furthermore, many of our better articles are certainly capable of being used as textbook texts. The biology article for example, has received good feedback from students and it certainly uses an expository register as one would expect in a competently written textbook. Please bring any objections to text and context to the discussion page before making such large deletions, as is the normal procedure here on CZ. --Thomas Simmons 16:17, 29 November 2007 (CST)

RE: Paul Wormer's rewrite: I think it reads very well. Explains some of the older references the reader might encounter to a greater degree as well. --Thomas Simmons 19:11, 4 December 2007 (CST)

I've gone through and corrected a number of mistakes and clarified the text of the article. The "textbook" examples are rather confusing and have fairly low content - if there are to be examples, they should be clear and more practically useful. I have separated the examples into a separate section, but have not made any changes. Anthony Argyriou 23:20, 7 December 2007 (CST)

• Why did you delete the few words that remind the reader that a mole of water is 18 mL? Further you undid the high school-type mnemonics in the "loose" definition: a mole weighs as much as the molecular weight. I put both in inspired by the forum discussion on this subject, which taught me that this sort of thing is appreciated by non-scientists.--Paul Wormer 03:39, 8 December 2007 (CST)

The density of liquid water isn't relevant to the discussion of the concept of mole. The result that one mole of any gas occupies the same volume is far more interesting, and should stay, but pointing out that one mole of particular liquid will have a volume numerically equal to its mass is potentially confusing as most other liquids do not have this property. I'd missed the parallelism between "weight" and "molecular weight", though I'm not sure quite how useful that is, as it propagates an error (mass != weight). Unfortunately, the term "molecular mass" is uncommon, even though "atomic mass" is used as much as "atomic weight". Anthony Argyriou 10:44, 8 December 2007 (CST)

This makes the point that any constable or editor would have made, discuss these changes when they are substantive. If the density of water can clarify the issue then it is a supporting concept. So assertions about what is relevant and what is far more interesting are to be discussed here before deletions are made. Tangents are not necessarily distracting if they help the reader understand. --Thomas Simmons 17:44, 21 December 2007 (CST)

• (1) Most people have an idea about an amount of water, and 18 mL (if I had known American measures, like teaspoon, etc. I would have added that) illustrates how much a mole of water is. (2) Lots of people, even scientists, still say something "weighs" some grams. But I added a sentence (that you removed) explaining that one shouldn't. --Paul Wormer 10:56, 8 December 2007 (CST)

He has a point, not so many years ago a lot of people were being told "molecular weight in grams." It is still being used. The reason is that most people figure weight in mass units not weight units even though they do not know it. Less of a heavy hand with the deletions please. --Thomas Simmons 17:44, 21 December 2007 (CST)

Testing readability

I had a couple of high school graduates read this, both scored well into the upper 700s on all three sections on the SAT and with four years of science and chemistry in a very good high school in the midwest. Their eyes glazed over when they got to this section: The total mass of an amount of substance is the sum of the masses of its entities. For example, consider a pure substance B of entities with molecular mass M(B) u ( u is unified atomic mass unit). Recalling that 1 u = 1/NA gram, we find that one mole of B weighs NA × M(B) u = M(B) gram.

For example, "entities" is used in common lanaguage as say deities, super natural beings. The use of commonly understood words is to be encouraged but when they carry wholly different meanings it is obfuscating.

Another problem is that the equations need to be written out as well. How long after they start reading equations are they still having trouble reading them? Spelling them out here would be a good way of reaching the targeted audience. Remembering that those who know what a mole is will not be reading this to learn how to calcualte molarity etc.--Thomas Simmons 17:44, 21 December 2007 (CST)

As a further test, could you ask your readers to read this older version of the article and ask them for their opinion of that text? That's the last version before examples were added. Anthony Argyriou 18:02, 21 December 2007 (CST)

edits on 21 december 2007

Thomas - the edit you've made to the article, adding:

and is used to signify how much or how many just as one would use "one kilogram" or "one dozen". It is a very large number however and it is used to indicate amounts of molecules or single atoms of substances. Since they are very small (a drop of water might contain trillions of molecules) the number is correspondingly large.

appears to detract from the article much more than improving it. Firstly, the tone really strikes me as talking down at the reader. Secondly, it introduces several inaccuracies and ambiguities. A mole is technically a count, therefore "just as one would use one kilogram" is not truly accurate - a kilogram of carbon is a different number of moles than a kilogram of uranium, while a mole of each is the same number of atoms. It's true that a drop of water contains "trillions" of molecules; in fact, a 1ml drop of water contains tens of millions of trillions of molecules. It would be better to explain the magnitude correctly than to just wave around numbers like trillions. Thirdly, the placement of the edit interrupts the previous flow of the article, and the edit essentially duplicates the former second paragraph. I'm refraining from reverting or hacking up your edit, but I hope you will take a second look at it and revise the article in a way which improves it. Anthony Argyriou 18:42, 21 December 2007 (CST)

Well, first, it may appear to detract but it does not. Having used this approach myself to good effect and having seen others using it at the high school and undergraduate level and consistently by those who get rated highly by undergraduate students for effective teaching, it is a clear and concrete way to get the message across, the concept is about a means to quantify and as such is not some strange new and complicated language. It reassures the reader that they are dealing with something they can understand and the more knowledgeable will know that the writer is adjusting it for the level it will most effectively serve, the people who are new to the concept.

And no, I do disagree, the people who will read this to learn will not feel they are being talked down to. Using clear language and building on what they already know is effective in teaching. And yes a mole is a count but more importantly it is a means to signify quantity as are the well-known examples ‘a dozen’ or ‘a kilogram’. Quantities, be they mass or enumerated are quantities. The lexical import is hypernymic.

Splitting hairs over trillions and millions of trillions could be obfuscatory but on the other hand the point is to impress upon the reader that the concept is essential when dealing with a very large number of very small things.

Concept complexity is effectively taught by building on known concepts. The introduction starts at an introductory level, makes the point that the topic is about a means to indicate quantities, drawing upon commonly known concepts and then the continuation gets more complicated having set the stage so to speak. It builds in complexity and as such allows the reader to start at the beginning and not at the middle. So, no, I can not agree that it interrupts the flow of the article. It does place it in context before continuing on.--Thomas Simmons 23:10, 21 December 2007 (CST)

Large number?

Thomas, you write:

[...] the mole [...] is used to signify how much or how many just as one would use "one kilogram" or "one dozen". It is a very large number [...]

but you mean to say that the number of molecules in a mole (Avogadro's number) is a large number. You don't say "one kilogram" or "one dozen" is a large number, do you? One mole water is 18 mL, compared to the Atlantic ocean this is not a large number. --Paul Wormer 05:38, 22 December 2007 (CST)

The antecedent is usually the subject of the prior sentence. I will make this redundant if you think it will help.--Thomas Simmons 06:03, 22 December 2007 (CST)

I agree with Anthony Argyriou that the Examples section should be deleted

I am compelled to say that I agree with Anthony Argyriou that the section on "Examples" should be deleted. I am retired chemical engineer and, like most chemical engineers, I used the word "mole" on a daily basis in oral and written communications and as an indispensible part of the process designing of oil refineries, natural gas processing facilities, chemical and petrochemical plants. I know what a mole is as well as I know my own name. But when I read that section, I found it difficult to understand. If I found it difficult, I can just imagine how difficult it would be for young students or lay people.

The article readership will include young students and lay people as well as advanced, experienced chemists and physicists. For the sake of those young students and the lay readers, the article should be as simple as possible without compromising any technical integrity.

These two sentences could be particularly confusing:

Futher examples: a mole of hydrogen molecule, standard atomic weight of H is 1.00794, has the mass 2×1.00794 = 2.01588 × 10⁻³ kg. A mole of oxygen, standard atomic weight 15.9994, has the mass 31.9988 × 10⁻³ kg.

Why is the mass of a hydrogen or oxygen molecule twice as much as the atomic weight of a hydrogen or oxygen atom? Does the article assume that everyone knows that hydrogen gas and oxygen gas are diatomic (meaning they have 2 atoms)? Should that not be explained?

In my opinion, the part of the Examples section that tries to explain the reactions 2A + 6B → 2AB₃ and 2H₂ + O₂ → 2H₂O really belongs in article called "Chemical stoichiometry" rather than in an article defining the word mole.

As for the notation M(A) or M(Mg) and so on, where did that notation come from? Personally, I have never seen that notation before. Would it not be better to use M_A and M_Mg and also to explain that M is the molecular mass (or, more commonly, the molecular weight)? Does the article assume that everyone reading the article will know what M is?

I would suggest this simple but correct definition a mole:

A mole of any chemical or compound is that quantity, when weighed, which will have a weight in grams that is numerically equivalent to that chemical's or compound's molecular mass (commonly referred to as molecular weight).

For readers wanting to know what the "molecular mass" is, the link in the above definition points him to the excellent molecular mass article written by Paul Wormer. If the reader wants to dig even deeper, that article has a link to the equally excellent unified atomic mass unit article also written by Paul Wormer. In other words, this mole (unit) article need not re-invent the wheel by repeating information available in Paul Wormer's two articles. In that way, this article could be a simple definition of the word "mole" without becoming involved in the strict SI definition of a mole, or "elementary entities", or Avogadro's number or other chemistry and physics textbook subjects that may have some relation to the word "mole". Please excuse me for being so long-winded. - Milton Beychok 17:50, 6 April 2008 (CDT)

You ask: As for the notation M(A) or M(Mg) and so on, where did that notation come from? Personally, I have never seen that notation before. See: http://www.bipm.org/en/si/si_brochure/chapter2/2-1/mole.html (BIPM is Bureau International des Poids et Mesures, the official international organization that settles these things.)

For the reason that there are some examples, see [1]

--Paul Wormer 18:00, 6 April 2008 (CDT)

Thanks, Paul. I guess this old dog needs to learn some new tricks. I always thought that 3(4) was 12. In other words, A(B) meant A × B. But recently I have learned that A(B) can also mean that A is a function of B and now I have just learned that M(B) means the molecular weight of B. (Sorry, I just cannot get accustomed to using "molecular mass" instaed of "molecular weight").

I think that encylopedias should explain things and if a notation indigenous to a certain discipline is used, then it should be explained for those who are not indigenous to that certain discipline. In plainer language, jargon peculiar to certain disciplines should always be defined. I am probably guilty of not always doing that in some of my articles, but I do try to follow that maxim. - Milton Beychok 22:51, 6 April 2008 (CDT)

An example of unnecessary confusion

The first paragraph in the Examples section is:

The total mass of an amount of substance is the sum of the masses of its entities. For example, consider a pure substance B of entities with molecular mass M(B) u ( u is unified atomic mass unit). Recalling that

1 u = 1/NA gram,

we find that one mole of B weighs NA × M(B) u = M(B) gram. For example, NA molecules of water (B = H2O) have mass 18.02 gram.

Let's look at that. We have:

• (1) One mole of B weighs NA × M(B) u grams
• (2) 1 u = 1/NA gram

If we substitute 2 into 1, we have:

• One mole of B weighs NA × M(B) 1/NA grams, or simply one mole of B weighs M(B) grams.

So why bring u and NA into the article at all? They are not necessary. The article should not be written only for chemists, physicists and similar professionals. Why not just simply say:

One mole of B weighs M(B) grams, where M(B) is the notation for the molecular weight of B?

That would be much less confusing to young students and to lay people. This same item of confusion was noted earlier on this Talk page by Thomas Simmons on December 21, 2007. - Milton Beychok 23:34, 6 April 2008 (CDT)

I must agree that this is overly complicated and likely to scare people away from chemistry. I always found that comparisons to the units dozen and gross always seemed to clear things up for the college students who did not have chemisty in high school. I would like to see the article start with an operational definition first, ie ~ 6.023 x 10^23 particles. Then describe that this many particles has the atom mass listed in the periodic table. Then describe that this is because of how it is defined, and so on. I would first mention 12 g= 1 mole of carbon-12, rather than using the 0.012 g, which can still be stated later. Just my 2 cents. David E. Volk 10:34, 7 April 2008 (CDT)

intro paragraph

I would like something along this line for the introductory paragraph:

The '’’mole’’’ is the SI unit that quantifies the amount of a substance, usually atoms, ions, molecules or elementary particles, that is equal to about 6.023 x 10^23 items. It functions similarly to the more common units of substance, the dozen (12) and the gross (144), but is much larger. It is widely used by chemists and physicists to determine the amounts substances to use in chemical reactions and other experiments. The mole is a convenient unit because, by definition,[1] one mole of any chemical element has the atomic mass listed on the period table for that element. Thus, one mole of carbon atoms weights 12.011 g/mol. By adding the atomic mass of the atomic substituents in a chemical, like water for example (H20), chemists can quickly determine the amounts (in grams) of hydrogen and oxygen atoms to react together to form water. They can also predict the mass of the water produced from the reaction. Because two atoms of hydrogen are needed for every atom of oxygen, one needs 2.00588 grams of hydrogen (two moles) to react with 15.9994 grams of oxygen (one mole), to yield 18.00528 grams (1 mole) of water. For industrial scale chemical reactions, chemical engineers have derived similar units called the kg-mol and the pound-mol.

[1] give the 0.012 kg definition here.

David E. Volk 11:27, 7 April 2008 (CDT)

David, I don't know about chemists, but chemical engineers would write the equation for the reaction between gaseous hydrogen and gaseous oxygen as:

H₂ + 0.5 O₂ → H₂O

or as:

2 H₂ + O₂ → 2 H₂O

The first equation above would written out as: "1 mole of hydrogen gas plus one-half mole oxygen gas yields 1 mole of liquid water".

The second equation would be written out as: 2 moles of hydrogen gas plus 1 mole of oxygen gas yields 2 moles of liquid water.

Leaving out the decimal digits for simplicity, in both equations the molecular mass of 1 mole of hydrogen would be 2, the molecular mass of 1 mole of oxygen would be 32 and the molecular mass of water would be 18.

Thus, in the first equation: 2 g of hydrogen gas + 16 gas of oxygen = 18 g of water, and in the second equation: 4 g of hydrogen gas + 32 g of oxygen gas = 36 g of liquid water

The point I want to make is that hydrogen and oxygen exist in nature as diatomic gases not monatomic gases as indicated in your suggested introductory section. Thus, most chemical (and probably other) engineers think of the molecular mass of hydrogen gas being 2 and of oxygen gas as being 32 ... instead the 1 and 16 indicated in your introductory section.

Thus, chemical engineers would not say Because two atoms of hydrogen are needed for every atom of oxygen, one needs 2.00588 grams of hydrogen (two moles) to react with 15.9994 grams of oxygen (one mole) .... Instead they would say Because two atoms of hydrogen are needed for every atom of oxygen, one needs 2.00588 grams of hydrogen (one mole) to react with 15.9994 grams of oxygen (one-half mole> .... The problem is that chemists, physicists and engineers should really undertand each other's language when writing articles in an encyclopedia. - Milton Beychok 17:23, 7 April 2008 (CDT)

The subject of the Examples section appears to have been well discussed by now

Anthony Argyriou voiced his opinion earlier on this Talk page that the Examples section should be removed.

Thomas Simmons ran a readability test on some students and said their "eyes glazed over" when they got to the Examples Section.

User:Milton Beychok (that's me), stated that he agreed with Anthony Argyriou that the Examples section be removed and pointed out that omission of the fact that hydrogen and oxygen are diatomic gases created confusion. He also suggested that the discussion of reactions in the Example section really belonged in an article about chemical stoichiometry.

User:Milton Beychok, in another comment, pointed out a specific example of how unnecessarily complicated was a part of the Example section.

User:David E. Volk commented that " I must agree that this is overly complicated and likely to scare people away from chemistry."

It appears to me that there is obviously a concensus that the Examples section as it now stands is very confusing. If there are no further comments within a day or two, I plan to remove that section. Milton Beychok 16:39, 10 April 2008 (CDT) CZ author and editor

Derek Harkness won't be pleased. He suggested (i) an example where the mole is good for (ii) a titration example for this. Since I hadn't a titration example handy I added a simple (or so I thought) example from chemical kinetics. If we cannot write anything on chemistry that may scare high-school kids, I have to reconsider my relation to CZ. --Paul Wormer 20:38, 10 April 2008 (CDT)

Look, we have "tutorial" and "advanced" subpages. There's no reason to make this unpleasant. Surely some compromise can be made. Maybe it just needs to be relocated within the cluster. There's no reason these alternatives can't be considered; let's not limit ourselves. --Robert W King 20:40, 10 April 2008 (CDT)

Paul, I have re-written the Examples section in a way which I believe is more explanatory and hopefully less confusing. I changed the usage of X^-3 kg to X grams. I also removed the last part of the section which I believe is really much better suited to an article on chemical stoichiometry. In the spirit of compromise, I would like to propose this as the revised section on Examples:

"The atomic weight of a pure substance B is denoted as M(B) and is expressed in unified atomic mass units which are denoted by u. For example, the atomic weight of magnesium (Mg) is 24.305 u and is denoted as M(Mg).

Avogadro's number (≈ 6.02 × 10²³), denoted by N_A, is the number of entities, such as atoms or molecules, in a mole of substance.

The unified atomic mass unit u is, by definition, the reciprocal of Avogadro's number N_A and can be expressed as 1 u = 1/N_A, in grams.

Thus, the mass of a mole of Mg is:

(Number of atoms in a mole of Mg} × (M(Mg), in u) × (1/N_A, in grams/u) = grams

and numerically, we have:

( 6.02 × 10²³ ) × ( 24.304 ) × 1 / ( 6.02 × 10²³ ) = 24.304 grams

This example illustrates the fact that the atomic mass of any element can be interpreted in two ways: (1) the average mass of a single atom in unified atomic mass units (u) or (2) the average mass, in grams, of a mole of the atoms.

As a similar example, water (H₂O) has two atoms of hydrogen and 1 atom of oxygen. The atomic weight of a hydrogen atom, denoted as M(H), is 1.008 u, and the atomic weight of oxygen, denoted as M(O), is 15.999 u. Thus, the atomic weight of water is (2)(1.008) plus 15.999, which is 18.015 u and the mass of a mole of water would be:

( 6.02 × 10²³ ) × ( 18.015) × 1 / ( 6.02 × 10²³ ) = 18.015 grams

As further examples, hydrogen and oxygen gases exist in nature as the diatomic molecules H₂ and O₂. A mole of gaseous hydrogen (H₂), with an atomic weight for H of 1.008 u, has a mass of 2.016 grams and a mole of gaseous oxygen (O₂), with an atomic weight for O of 15.999 u, has a mass of 31.998 grams."

I hope this meets with your agreement. - Milton Beychok 03:00, 11 April 2008 (CDT)

Does anyone else have an opinion regarding my proposed re-write? If so, please post your thoughts. - Milton Beychok 01:20, 15 April 2008 (CDT)