Mole (unit): Difference between revisions
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The definition of a mole extends to different entities: molecules, ions, atoms, electrons - any elementary substance in chemistry. The [[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> | The definition of a mole extends to different entities: molecules, ions, atoms, electrons - any elementary substance in chemistry. The [[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> | ||
The definition of the mole leads to the definition of [[Avogadro's constant]]: ''N''<sub>A</sub> | The definition of the mole leads to the definition of [[Avogadro's constant]]: ''N''<sub>A</sub> is the number of entities in one mole. ''N''<sub>A</sub> ≈ 6×10<sup>23</sup>/mol | ||
One mole of an [[ideal gas law|ideal gas]] occupies 22.414 [[litre]]s at "standard temperature and pressure" (273.15K = 0 <sup>0</sup>C and 101.325 kPa = 1 atm). Real gases | One mole of an [[ideal gas law|ideal gas]] occupies 22.414 [[litre]]s at "standard temperature and pressure" (273.15K = 0 <sup>0</sup>C and 101.325 kPa = 1 atm). Real gases deviate slightly from this amount. | ||
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 | 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 | ||
Revision as of 00:16, 8 December 2007
In chemistry and physics, the mole is an SI base unit of amount of substance. The unit is abbreviated mol. The word "mole" is derived from "gram molecular weight", the original term. Industrial chemists also used a "kilogram molecular weight", equal to 1 kmol.
Loosely speaking, the mole may be defined for a pure substance, consisting of one kind of molecules, as the amount of substance that has a mass in grams equal to the molecular weight of the molecule. For instance, the molecular weight of the water molecule (H2O) is 18.02, and therefore one mole of water is 18.02 grams of water.
The definition of a mole extends to different entities: molecules, ions, atoms, electrons - any elementary substance in chemistry. The 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).[1]
The definition of the mole leads to the definition of Avogadro's constant: NA is the number of entities in one mole. NA ≈ 6×1023/mol
One mole of an ideal gas occupies 22.414 litres at "standard temperature and pressure" (273.15K = 0 0C and 101.325 kPa = 1 atm). Real gases deviate slightly from this amount.
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. One mole of magnesium is NA × M(Mg) u = 24.3050 10−3 kg. 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.
Futher examples: a mole of hydrogen molecule, standard atomic weight of H is 1.00794, has the mass 2×1.00794 = 2.01588 −3 kg. A mole of oxygen, standard atomic weight 15.9994, has the mass 31.9988 10−3 kg. [2]
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
- ↑ 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.
- ↑ The Mole Concept (Avogadro's Number) N..De Leon, Indiana University, Northwest
Sources
- mole. Sizes.com (2006-11-07). Retrieved on 2007-05-11.