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Molecules and atoms are extremely small objects - both in size and mass. Consequently, working with them in the laboratory requires a large collection of them. How large does this collection need to be? A standard needs to be introduced. This standard is the "mole". The mole is based upon the carbon-12 isotope. We ask the following question: How many carbon-12 atoms are needed to have a mass of exactly 12 g. That number is NA - Avogadro's number. Thus, NA is defined by

NA x (mass of carbon-12 atom) = 12 g

Careful measurements yield a value for NA = 6.0221367x10^+23. This is an incredibly large number - almost a trillion trillion. For example, if we stack NA pennies on top of one another how tall would the stack be? The answer is it would be so tall that the stack of pennies could reach the sun and back almost 500 million times!

A convenient name is given when there is an Avogadro's number of objects - it is called a "mole". Thus in the above example there was amole of pennies.

1 mole = NA objects

The mole concept is no more complicated than the more familiar concept of a dozen : 1 dozen = 12 objects. From the penny example above one might suspect that the mass of a mole of objects is huge. Well, that is true if we're considering a mole of pennies, however a mole of atoms or molecules is a different story. Recall that the atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 atom. Consequently we have the relation


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A mole is the amount of pure substance containing the same number of chemical units as there are atoms in exactly 12 grams of carbon-12 (i.e., 6.023 X 1023). Formerly, the connotation of "mole" was "gram molecular weight." Current usage tends to apply the term "mole" to an amount containing Avogadro's number of whatever units are being considered. Thus, it is possible to have a mole of atoms, ions, radicals, electrons, or quanta.
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