Mole Concept : The identity of a substance is defined not only by the types of atoms or ions it contains. But by the quantity of each type of atom or ion.
For example : Water, H2O, and hydrogen peroxide, H2O2, are alike in that their respective molecules are composed of hydrogen and oxygen atoms. However, because a hydrogen peroxide molecule contains two oxygen atoms instead of the water molecule, which has only one. The two substances exhibit very different properties.
Today, sophisticated instruments allow the direct measurement of these defining microscopic traits. However, the same traits were originally derived from the measurement of macroscopic properties using relatively simple tools.
This experimental approach required introducing a new unit for the number of substances, the mole. Which remains indispensable in modern chemical science.
Definition of Mole Concept
A mole is an amount unit similar to familiar units like pair, dozen, gross, etc. It provides a specific measure of the number of atoms or molecules in a sample of matter.
A mole of a substance is that amount in which there are 6.02214076 × 1023 discrete entities (atoms or molecules). This large number is a fundamental constant known as Avogadro’s number (NA) or the Avogadro constant (Avogadro’s Hypothesis) in honour of Italian scientist Amedeo Avogadro.
This constant is properly reported with an explicit unit of “per mole,” a conveniently rounded version being 6.022 × 1023/mol.
Consistent with its definition as an amount unit, 1 mole of any element contains the same number of atoms as 1 mole of any other element.
However, the masses of 1 mole of different elements are different. Since the masses of the individual atoms are drastically different.
The molar mass of an element (or compound) is the mass in grams of 1 mole of that substance, a property expressed in units of grams per mole (g/mol).
Other Definitions of Mole Concept
Here are some more definitions related to the mole concept :
Definition-1 (mole concept)
One Latin connotation for the word “mole” is “large mass” or “bulk”. Which is consistent with its use as the name for this unit.
The mole provides a link between an easily measured macroscopic property, bulk mass, essential fundamental property, number of atoms, molecules, and so forth.
Because the definitions of both the mole and the atomic mass unit are based on the same reference substance, 12C. The molar mass of any substance is numerically equivalent to its atomic or formula weight in amu. Per the amu definition, a single 12C atom weighs 12 amu (its atomic mass is 12 amu).
According to the definition of the mole, 12g of 12C contains 1 mole of 12C atoms (its molar mass is 12 g/mol).
This relationship holds for all elements since their atomic masses are measured relative to the AMU-reference substance, 12C.
Extending this principle, the molar mass of a compound in grams is likewise numerically equivalent to its formula mass in amu.
The mole (symbol: mol) is the base unit of the amount of substance in the International System of Units (SI). It is defined as exactly 6.02214076×1023 particles, which may be atoms, molecules, ions, or electrons.
The definition of the mole was adopted in November 2018 as one of the seven SI base units. Revising the previous definition that specified one mole as the amount of substance in 12 grams of carbon-12 (12C), an isotope of carbon.
The number 6.02214076×1023 (the Avogadro number) was chosen so that the mass of one mole of a chemical compound in grams is numerically equal, for most practical purposes, to the average mass of one molecule of the compound in Dalton’s.
For example, one mole of water (H2 O) contains 6.02214076×1023 molecules, whose total mass is about 18.015 grams, and the mean mass of one molecule of water is about 18.015 Dalton’s.
For example, the chemical equation 2H2 + O2 → 2H2O can be interpreted to mean that for each 2 mol dihydrogen (H2) and 1 mol dioxygen (O2) that react, 2 mol of water (H2O) form.
The mole may also measure the number of atoms, ions, electrons, or other entities.
The concentration of a solution is commonly expressed by its molarity. Which defined as the amount of dissolved substance in mole(s) per unit volume of solution. For which the unit typically used is moles per litre (mol/L), commonly abbreviated M.
The term gram-molecule (g mol) was formerly used for “mole of molecules”, and gram-atom (g atom) for “mole of atoms”.
For example, 1 mole of MgBr2 is 1 gram-molecule of MgBr2 but 3 gram-atoms of MgBr2.
Importance of Mole Concept
- The mole concept is a convenient method of expressing the amount of a substance. Any measurement can be broken down into two parts – the numerical magnitude and the magnitude’s units. For example, when a ball’s mass is measured to be 2 kilograms, the magnitude is ‘2’ and the unit is ‘kilogram’. When dealing with particles at an atomic (or molecular) level, even one gram of a pure element is known to contain a huge number of atoms. This is where the mole concept is widely used. It primarily focuses on the unit known as a ‘mole’, which is a count of a very large number of particles.
- In chemistry, a mole is defined as the amount of a substance that contains exactly 6.02214076 * 1023 ‘elementary entities’ of the given substance. The number 6.02214076*1023 is popularly known as the Avogadro constant and is often denoted by ‘NA’. The elementary entities that can be represented in moles can be atoms, molecules, monoatomic/polyatomic ions, and other particles (such as electrons).
- The word “mole” was introduced around 1896 by the German chemist Wilhelm Ostwald. Who derived the term from the Latin word moles, meaning a ‘heap’ or ‘pile. The number of moles of a molecule may not always be equal to the number of moles of its constituent elements. For example, a mole of water contains an NA number of H2O molecules. However, each water molecule has 2 hydrogen atoms and one oxygen atom. Therefore, one mole of H2O contains 2 moles of hydrogen and one mole of oxygen.
History of Mole Concept
The history of the mole is intertwined with that of molecular mass, atomic mass units, and the Avogadro number.
The first table of standard atomic weight (atomic mass) was published by John Dalton (1766–1844) in 1805. Based on a system in which the relative atomic mass of hydrogen was defined as 1. These relative atomic masses were based on the stoichiometric proportions of chemical reaction and compounds.
A chemist didn’t need to subscribe to atomic theory to make practical use of the tables. This would lead to some confusion between atomic masses and equivalent weights. Which would last throughout much of the nineteenth century.
Jones Jacob Berzelius (1779–1848) was instrumental in the determination of relative atomic masses to ever-increasing accuracy. He was also the first chemist to use oxygen as the standard to which other masses were referred. Oxygen is a useful standard, as, unlike hydrogen, it forms compounds with most other elements, especially metals. However, he chose to fix the atomic mass of oxygen as 100, which did not catch on.
Here are some mole concepts listed below
1. Nature of The Particles
The mole is essentially a count of particles. Usually, the particles counted are chemically identical entities, individually distinct.
For example, a solution may contain a certain number of dissolved molecules that are more or less independent of each other. However, the constituent particles are fixed and bound in a lattice arrangement in a solid, yet they may be separable without losing their chemical identity. Thus the solid is composed of a certain number of moles of such particles.
In yet other cases, such as diamond, where the entire crystal is essentially a single molecule. The mole is still used to express the number of atoms bound together rather than a count of multiple molecules.
Thus, common chemical conventions apply to the definition of the constituent particles of a substance.
In other cases, exact definitions may be specified. The mass of 1 mole of a substance equals its relative atomic or molecular mass in grams.
2. Molar Mass
The molar mass is the mass of 1 mole of that substance, in multiples of the gram. The amount of substance is the number of moles in the sample. For most practical purposes, the magnitude of molar mass is numerically the same as that of the mean abundance of one molecule, expressed in Dalton’s.
For example, the molar mass of water is 18.015 g/mol.
Other methods include the use of the molar volume or the measurement of electric charge. The number of moles of a substance in a sample is obtained by dividing the mass of the sample by the molar mass of the compound.
For example, 100g of water is about 5.551 mol of water.
The molar mass of a substance depends not only on its molecular formula but also on the distribution of isotopes of each chemical element present in it.
For example, the mass of one mole of calcium-40 is 39.96259098±0.00000022 grams, whereas one mole of calcium-42 is 41.95861801±0.00000027 grams, and one mole of calcium with the normal isotopic mix is 40.078±0.004 grams.
3. Molar Concentration
The molar concentration, also called molarity, of a solution of some substance is the number of moles per unit of volume of the final solution. The SI’s standard unit is mol/m3, although more practical units, such as mole per litre (mol/L), are used.
4. Molar Fraction
The molar fraction or mole fraction of a substance in a mixture (such as a solution) is the number of moles of the compound in one sample of the mixture, divided by the total number of moles of all components.
For example, if 20 g of NaCl is dissolved in 100 g of water, the amounts of the two substances in the solution will be (20g)/(58.443 g/mol) = 0.34221 mol and (100g)/(18.015 g/mol) = 5.5509 mol, respectively. And the molar fraction of NaCl will be 0.34221/(0.34221 + 5.5509) = 0.0580 7.
In a mixture of gases, the partial pressure of each component is proportional to its molar ratio
October 23, denoted 10/23 in the US, is recognized by some as Mole Day. It is an informal holiday in honour of the unit among chemists.
The date is derived from the Avogadro number, which is approximately 6.022×1023. It starts at 6:02 a.m. and ends at 6:02 p.m.
Alternatively, some chemists celebrate June 2 (06/02), June 22 (6/22), or February 6 (06.02), a reference to the 6.02 or 6.022 part of the constant.
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