1. 1 Chemistry 11 Chapter 4 - The MOLE

2. 2 Relative Atomic Mass Dalton, concerned with how much one element could combine with a given amount of element, hypothesized:  Molecules are made of “atoms” of various elements  If compound B contains twice the mass of element X as does compound A, then compound B must contain twice as many atoms of X  Simple compounds are made up of only one atom of each of the two elements making up the compound (this was later proved to be wrong) John Dalton (1766-1844)

3. 3 Dalton did not attempt to figure out the mass of an individual atom of any element. (atoms are SO small!) Instead, he simply assigned an arbitrary mass to each element, assuming that hydrogen was the lightest element (which had a mass of “1”) Experiments showed that C was 6x heavier than H, so C was assigned a mass of “6” (similarly O had a mass of “16” as it was 16x heavier) In this way, Dalton was eventually able to calculate “relative masses”

4. 4 Nowadays... The standard used for measuring atomic mass is Carbon-12 (instead of hydrogen), since C is a very common element and thus readily available. Relative atomic mass is measured in units of Atomic Mass Units (a.m.u.), which is equal to 1/12 the mass of the C-12 atom Eg. If C-12 has a mass of 12.0000 amu, then H has a mass of 1.007825 amu

5. 5 Other contributions... Joseph Gay-Lussac (1778) experimented with N and O. He found that for all compounds formed from N and O, the ratios of volumes of gases used were whole number ratios: Eg. 100 mL N 2 + 100 mL O 2 forms NO (Ratio 1:1) 100 mL N 2 + 200 mL O 2 forms NO 2 (Ratio 1:2) Avogadro (1811) then made the assumption that “equal volumes of different gases, at constant temperature and pressure, contain the same number of particles” (this is called Avogadro’s Hypothesis) In other words, if 1 L of gas A reacts with 1 L of gas B, then there are exactly as many particles of A present as B. The molecule formed is AB; if 2 L of gas A reacts with 1 L of gas B, then there are TWICE as many atoms of A as B, so A 2 B will form Joseph Gay-Lussac Amadeo Avogadro (1776-1856)

6. 6 The MOLE - a unit of measurement Suppose we want to make FeS molecules in such as way that there wouldn’t be any extra Fe and S atoms left. 10 Fe atoms + 10 S atoms = 10 FeS molecules 100 Fe atoms + 100 S atoms = 100 FeS molecules 10 000 Fe atoms + 10 000 S atoms = 10 000 FeS molecules BUT individual atoms are extremely small and so impossible to count out… (in fact the mass of one Fe atom is 9.27 x 10 -23 g, mass of one S atom = 5.33 x 10 -23 g) So how do we deal with this problem?

7. 7 The periodic table shows that: atomic mass of Fe = 55.8 amu atomic mass of S = 32.1 amu SO the masses of Fe and S atoms are in the ratio: mass Fe = 55.8 amu mass S 32.1 amu Provided equal numbers of Fe and S atoms are used, then the ratio will always be 55.8:32.1 Eg.mass of 1000 Fe atoms = 1000 x 55.8 = 55.8 mass of 1000 S atoms 1000 x 32.1 32.1

8. 8 SO to make FeS we could weigh out and react 55.8 g of Fe with 32.1 g of S Hence in chemistry the concept of the MOLE (mol)- a unit of measurement representing the AMOUNT OF SUBSTANCE Experimentally, it has been found that... 1 mol particles= 6.02 x 10 23 particles [particles = atoms, ions, molecules, electrons etc.] 6.02 x 10 23 is known as AVOGADRO’S NUMBER (N A ) (named in his honor - he didn’t invent it) Analogies:1 pair = 2 1 dozen = 12 1 ream = 500 (usually paper) SO1 mol = 6.02 x 10 23

9. 9 Putting it into use... If 1 mol = 6.02 x 10 23 particles then 1 mol of C atoms = 6.02 x 10 23 C atoms The MOLAR MASS is the mass of ONE MOLE of particles in grams ElementAtomic Mass shown on Molar mass of element Periodic Table C12.0 amu12.0 g Fe55.8 amu55.8 g O16.0 amu16.0 g 1 mol C = 12.0 g of C 1 mol O = 16.0 g O 1 mol Fe = 55.8 g of Fe