Atomic weight mole concept

The mass of one atom of any element is infinitessimal and is impossible to measure on any existing balance. A more convenient mass unit was needed for laboratory work, and the concept of the mole emerged, where one mole of an element is a quantity equal to the atomic weight in grams. One mole of carbon, for example, is 12.01 grams, and one mole of iron is 55.85 grams. 

If we know the atomic weight of an element on the carbon-12 scale, we can use the mole concept and Avogadro s number to calculate the average mass of one atom of that element in grams (or any other mass unit we choose). 

Although the mole is defined as a mass in grams, the concept of the mole is easily extended to other mass units. Thus, the kilogram mole (kg mol) is the usual molecular or atomic weight in kilograms, and the pound mole (lb mol) is that in avoirdupois pounds. When the mass unit is not specified, the gram mole (g mol) is intended. Molecular weight M is a pure number. 

Today, Avogadro s name is most often associated with his number. The mole and the modem definition of the number of particles in it, however, were not directly his invention. His fame hes in a single simple statement, made in 1811 .. . the number of integral molecules in gases is always the same for equal volumes." This concept opened the door to understanding atomic weight and the formulas for chemical compounds. (The modem version of this concept is one mole of any gas always occupies the same volume [22.4 L] at 0°C and I atm of pressure.)... 

You can also use the mole concept to calculate the empirical formula of a compound using the percentage composition data for that compound — the percentage by weight of each element in the compound. (The empirical formula indicates the different types of elements in a molecule and the lowest whole-number ratio of each kind of atom in the molecule. See Chapter 7 for details.)... 

The number 6.022 X 10 is called Avogfldro s number in honor of Amadeo Avogadro (1776-1856), an Italian scientist who made important contributions to the concept of atomic weights. As we have seen, this number represents the number of atoms or molecules in a specific sample of an element or compound. Because of its importance in calculations, the number of particles represented by Avogadro s number is given a specific name it is called a mole, abbreviated mol. 

The mole concept can also be applied to particles that are molecules instead of atoms. The compound carbon dioxide consists of molecules that contain one carbon atom, C, and two oxygen atoms, O. The formula for the molecule is CO2. The molecular weight of the molecule is calculated as shown earlier by adding together the atomic weight of one carbon atom and the atomic weight of two oxygen atoms ... 

This early coverage of basic concepts such as the mole, atomic weights, ions and ionic charges, molecules, and formulas and composition provides an effective groundwork for subsequent study and calculations in chemistry. 

The mole is often referred to as a chemist s unit of quantity. Counting atoms is a difficult process and beyond the scope of most calculators, but measuring the mass of a sample is easy when we can relate the number of atoms in a sample to its mass. This is the unique purpose of the mole. A mole of any substance is its molecular formula weight expressed in grams. Avogadro s number s a universal constant that states the number of molecules in a mole Nq = 6.023 x 10 molecules/mole. One mole (abbreviated mol) of any element (chemical compound) has the same number of chemical particles as one mole of another element (chemical compound). In other words, 1 mole of any compound contains 6.02 x 10 molecules. Review the following problem using the mole concept. 

The Mole Concept The mole concept allows us to determine the number of atoms or molecules in a sample from its mass. Just as a hardware store customer wants to know the number of nails in a certain weight of nails, so we want to know the number of atoms in a certain mass of atoms. Since atoms are too small to count, we use their mass. 

One result of this apparently simple concept is that the atomic weight can be used to interconvert grams and moles. Since the units of atomic weight, as defined earlier, include both grams and moles, it can be used as a conversion factor to convert grams to moles, and vice versa. The dimensional analysis would appear as follows for the conversion from grams to moles ... 

The idea of chemical equivalents was stated by Henry Cavendish in 1767, clarified by Jeremias Richter in 1795, and popularized by William Wollaston in 1814. Wollaston applied the concept to elements and defined it in such a way that one equivalent of an element corresponded to its atomic mass. Thus, when Wollaston s equivalent is expressed in grams, it is identical to a mole. It is not surprising then that the word mole is derived from molekulargewicht (German, meaning molecular weight ) and was coined in 1901 or 1902. see ALSO Avogadro, Amedeo Cannizzaro, Stanislao Cavendish, Henry Gay-Lussac, Joseph-Louis. 

The term mole was first introduced by Wilhelm Ostwald in 1901. It is derived from the Latin for mass, hump, or pile (the term molecule, introduced by Pierre Gassendh in the early seventeenth century has the same root presumably it means a mass of atoms). Specifically, Ostwald used the term to represent the formula weight of a substance in grams 36.5 g of HCl is one mole. The formal definition of the mole adopted by the Fourteenth Conference Generale des Poids et Mesures in 1971 is the amount of a substance of a system that contains as many elementary entities as there are atoms in 0.012 kilograms of carbon-12. The rich irony is that Ostwald fiercely resisted the atomic concept at the time Boltzmann committed suicide in 1906 but his mole is now defined explicitly in terms of atoms. 

Amedeo Avogadro (1776-1856) was an Italian physicist who made many early contributions to the concepts of molecular behaviour and relative molecular mass (formerly known as molecular weight) (Figure 1.48). His most critical contribution was making the distinction between atoms and molecules. His hypothesis was based on the careful experimental work by Gay-Lussac and John Dalton s atomic theory. He trained and practised as a lawyer, but later became Professor of Physics at Turin University. As a tribute to him, the number of particles in one mole of a substance is known as the Avogadro constant (formerly known as the Avogadro number). 

Biomass, e.g., the cells that are formed in this example, is another example of a material whose structure firequently is so complex that the concepts of moles and molecular weight caimot be employed. The atomic ratios of the elements in a specific kind of cell often are constant. However, live cells are constantly growing and dividing, so the molecular weight is not constant firom cell to cell. 

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