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The '''mole''' is a unit of measurement in [[chemistry]]. It is defined in the [[SI]] as ''the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12.'' A mole of a substance has a mass in grams which is equal to the mass of a single atom or molecule of the substance in [[atomic mass units]].
In [[chemistry]] and [[physics]], the '''mole''' is an [[SI]] base unit of [[amount of substance]], used to signify how much or how many--just as one would use "one kilogram" or "one dozen". The unit is abbreviated '''mol'''. The word "mole" is derived from "gram '''mole'''cular weight", the original term. Industrial chemists and chemical engineers also use a "kilogram molecular weight" as equal to 1 kmol (or kg-mol) and in the [[United States of America]] as a pound-mol (lb-mol) which is equal to 453.592 gram mols.


Another way to phrase this is that a mole is '''''the molecular mass in grams'''''. Using the carbon-12 isotope, a mole of carbon-12 is 12 grams.
Informally the mole may be defined for a pure (i.e., consisting of one kind of molecules) substance by the easily remembered sentence: 
:''the mole is the amount of substance that '''weighs''' (in grams) as much as its molecular '''weight'''''. <ref>It is preferred to write ''has mass'' instead of ''weighs'', because the verb ''to weigh'' includes the gravitation of the Earth, which varies over the surface of the Earth.</ref>


We pose the question, “How many carbon-12 atoms are needed to have a mass of exactly 12 grams?
For instance, the [[molecular weight]] of the water molecule (H<sub>2</sub>O) is 18.02, and therefore one mole of pure water weighs 18.02 gram.  
That number, Avogadro's number, is the number of carbon-12 atoms in 12 grams of of carbon-12. The abbreviation for Avogadro’s number is NA. NA is defined by:


''' '' NA x (mass of carbon-12 atom) = 12 g”'''''
The definition of a mole extends to different entities: molecules, ions, atoms, electrons - any elementary substance in chemistry. The formal [[SI]] definition reads:
:''the mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of <sup>12</sup>C (carbon-12).''<ref>The <sup>12</sup>C isotope accounts for 98.89% of all carbon. It is one of two stable isotopes of the element carbon; it contains 6 protons, 6 neutrons and 6 electrons.</ref><ref name=BIPM>[http://www.bipm.org/en/si/si_brochure/chapter2/2-1/mole.html International Bureau of Weights and Measures] From the website of the [[International Bureau of Weights and Measures]].</ref> 


So, the number of entities (atoms or molecules) of a substance in one mole is known as Avogadro's constant, which is approximately 6.022 141 5 × 10<sup>23</sup>.  
The definition of the mole leads to the definition of [[Avogadro's constant]]: ''N''<sub>A</sub> is the number of entities (e.g. molecules or atoms) in one mole. ''N''<sub>A</sub> ≈ 6&times;10<sup>23</sup>/mol.


This also applies to other entities, magnesium for example. The atomic mass of magnesium is 24.305 amu,<ref>atomic mass unit</ref> the average isotopic mass of magnesium as it naturally occurs.
One mole of an [[ideal gas law|ideal gas]] occupies 22.41399 [[litre]]s at an absolute temperature ''T'' of 273.15 K = 0 &#176;C and pressure ''p'' of 101.325 kPa = 1 [[Atmosphere (unit)|atm]]. This follows from the [[ideal gas law|ideal gas equation]]:
What is the molar mass of magnesium in grams? From the equation NA x (mass of atom) = X grams” we get 1 amu = 1g/NA or 1 amu = 1.66054x10<sup>-24</sup> g. Using this we calculate for magnesium: NA x 24.305 amu x (1.66054x10<sup>-24</sup> g/amu) = 24.305 g
:<math>
V = \frac{RT}{p} \quad \Longrightarrow\quad 22.41399 =  \frac{8.314472\times 273.15}{101.325},
</math>
where ''R''  ( = 8.314472 J mol<sup>&minus;1</sup> T<sup>&thinsp;&minus;1</sup> ) is the [[molar gas constant]]. A mole of a real gas will deviate from this volume, but for many real gases 22.4 litres for a mole is a good&mdash;first&mdash;approximation.


This means that a mole of magnesium atoms has a mass of 24.305 grams. This example shows that the atomic mass of any element can be interpreted in two ways: (1) the average mass of a single atom in atomic mass units (amu) or (2) the average mass of a mole of atoms in grams. For magnesium, (1) the average mass of a single magnesium atom is 24.305 amu or (2) the average mass of a mole of magnesium atoms is 24.305 g;
===Examples===


Correspondingly: a mole of hydrogen, molecular mass 1.0079 is 1.0079 grams, a mole of lithium, molecular mass 6.94, is 6.94 grams. Molecules also have the same measure. A molecule of water, H<sub>2</sub>O is two hydrogen (at 2 times 1.0079) and one oxygen (15.9994) for a combined molecular mass of 18.0152. So a mole of water would contain 18.0152 grams.
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]]). [[Avogadro's number|Recalling]] that
<ref>[http://www.iun.edu/~cpanhd/C101webnotes/quantchem/moleavo.html The Mole Concept (Avogadro's Number)] N..De Leon, Indiana University, Northwest</ref>
: 1 u = 1/''N''<sub>A</sub> gram,
we find that one mole of B weighs  ''N''<sub>A</sub> &times; ''M''(B) u = ''M''(B) gram. For example, ''N''<sub>A</sub> molecules of water (B = H<sub>2</sub>O) have mass 18.02 gram.  


One mole of an [[ideal gas law|ideal gas]] occupies 22.414 [[litre]]s at "standard temperature and pressure" (273.15K and 101.325 kPa).
Let us give another example. The [[atomic mass|standard atomic weight]] of [[magnesium]] ''M''(Mg) =  24.3050 u, which is the mass of a magnesium atom averaged over its isotopes weighted by natural abundance.<ref>Evidently, a single magnesium atom with an average mass does not exist. A single atom is by definition a pure isotope and has an isotopic&mdash;not an average&mdash;mass. Yet it is convenient to do calculations with a fictive atom of average mass.</ref>  One mole of magnesium is ''N''<sub>A</sub> &times; ''M''(Mg) u = 24.3050 gram.
This example shows 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 unit]]s (u) or (2) the average mass of a mole of atoms in grams. For magnesium, (1) the average mass of a single magnesium atom is 24.3050 u or (2) the average mass of a mole of magnesium atoms is 24.3050 gram.


The word "mole" is shortened from "gram '''mole'''cular weight", the original term. Industrial chemists also used a "kilogram molecular weight", equal to 1000 mole.
Further examples: a mole of [[hydrogen]] molecule, standard atomic weight of H is 1.00794, has the mass 2&times;1.00794 = 2.01588 gram. A mole of [[oxygen]], standard atomic weight 15.9994, has the mass 31.9988 gram. <ref>[http://www.iun.edu/~cpanhd/C101webnotes/quantchem/moleavo.html The Mole Concept (Avogadro's Number)] N..De Leon, Indiana University, Northwest</ref>
 
To explain the usefulness of the mole concept we consider the following example of a chemical reaction:
: 2 A + 6 B &rarr; 2 AB<sub>3</sub>
This equation is expressed in numbers of atoms and molecules, which are impractical to measure or count directly. However, if we multiply the equation by Avogodro's number, it is expressed in macroscopic quantities. The equation tells us then that 2 moles of A react with 6 moles of B to give 2 moles of AB<sub>3</sub>.
 
In general, the molecular masses of the compounds A [ = ''M''(A)] and  B [ = ''M''(B)] are known and hence also is the molecular mass  of AB<sub>3</sub> [ = ''M''(AB<sub>3</sub>)]. The reaction equation can be translated thus: 2''M''(A) gram of A reacts with 6''M''(B) gram of B to give 2''M''(AB<sub>3</sub>) gram of AB<sub>3</sub>.
 
A real life example:
:2H<sub>2</sub> + O<sub>2</sub> &rarr; 2H<sub>2</sub>O 
Using rounded numbers: 4 gram H<sub>2</sub> reacts with 32 gram O<sub>2</sub> giving 36 gram H<sub>2</sub>O.


=Notes=
=Notes=
<div class="references-small" style="-moz-column-count:3; column-count:3;">
{{reflist}}
<references/>
</div>


==Sources==
==Sources==
*{{cite web|url=http://www.sizes.com/units/mole.htm|title=mole|publisher=Sizes.com|date=2006-11-07|accessdate=2007-05-11}}
*{{cite web|url=http://www.sizes.com/units/mole.htm|title=mole|publisher=Sizes.com|date=2006-11-07|accessdate=2007-05-11}}[[Category:Suggestion Bot Tag]]

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In chemistry and physics, the mole is an SI base unit of amount of substance, used to signify how much or how many--just as one would use "one kilogram" or "one dozen". The unit is abbreviated mol. The word "mole" is derived from "gram molecular weight", the original term. Industrial chemists and chemical engineers also use a "kilogram molecular weight" as equal to 1 kmol (or kg-mol) and in the United States of America as a pound-mol (lb-mol) which is equal to 453.592 gram mols.

Informally the mole may be defined for a pure (i.e., consisting of one kind of molecules) substance by the easily remembered sentence:

the mole is the amount of substance that weighs (in grams) as much as its molecular weight. [1]

For instance, the molecular weight of the water molecule (H2O) is 18.02, and therefore one mole of pure water weighs 18.02 gram.

The definition of a mole extends to different entities: molecules, ions, atoms, electrons - any elementary substance in chemistry. The formal SI definition reads:

the mole is the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of 12C (carbon-12).[2][3]

The definition of the mole leads to the definition of Avogadro's constant: NA is the number of entities (e.g. molecules or atoms) in one mole. NA ≈ 6×1023/mol.

One mole of an ideal gas occupies 22.41399 litres at an absolute temperature T of 273.15 K = 0 °C and pressure p of 101.325 kPa = 1 atm. This follows from the ideal gas equation:

where R ( = 8.314472 J mol−1 T −1 ) is the molar gas constant. A mole of a real gas will deviate from this volume, but for many real gases 22.4 litres for a mole is a good—first—approximation.

Examples

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 us give another example. The standard atomic weight of magnesium M(Mg) = 24.3050 u, which is the mass of a magnesium atom averaged over its isotopes weighted by natural abundance.[4] One mole of magnesium is NA × M(Mg) u = 24.3050 gram. This example shows 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 of a mole of atoms in grams. For magnesium, (1) the average mass of a single magnesium atom is 24.3050 u or (2) the average mass of a mole of magnesium atoms is 24.3050 gram.

Further examples: a mole of hydrogen molecule, standard atomic weight of H is 1.00794, has the mass 2×1.00794 = 2.01588 gram. A mole of oxygen, standard atomic weight 15.9994, has the mass 31.9988 gram. [5]

To explain the usefulness of the mole concept we consider the following example of a chemical reaction:

2 A + 6 B → 2 AB3

This equation is expressed in numbers of atoms and molecules, which are impractical to measure or count directly. However, if we multiply the equation by Avogodro's number, it is expressed in macroscopic quantities. The equation tells us then that 2 moles of A react with 6 moles of B to give 2 moles of AB3.

In general, the molecular masses of the compounds A [ = M(A)] and B [ = M(B)] are known and hence also is the molecular mass of AB3 [ = M(AB3)]. The reaction equation can be translated thus: 2M(A) gram of A reacts with 6M(B) gram of B to give 2M(AB3) gram of AB3.

A real life example:

2H2 + O2 → 2H2O

Using rounded numbers: 4 gram H2 reacts with 32 gram O2 giving 36 gram H2O.

Notes

  1. It is preferred to write has mass instead of weighs, because the verb to weigh includes the gravitation of the Earth, which varies over the surface of the Earth.
  2. The 12C isotope accounts for 98.89% of all carbon. It is one of two stable isotopes of the element carbon; it contains 6 protons, 6 neutrons and 6 electrons.
  3. International Bureau of Weights and Measures From the website of the International Bureau of Weights and Measures.
  4. Evidently, a single magnesium atom with an average mass does not exist. A single atom is by definition a pure isotope and has an isotopic—not an average—mass. Yet it is convenient to do calculations with a fictive atom of average mass.
  5. The Mole Concept (Avogadro's Number) N..De Leon, Indiana University, Northwest

Sources

  • mole. Sizes.com (2006-11-07). Retrieved on 2007-05-11.