1 The MOLE

2 Atoms and molecules are very, very small, but we still need to “count” them.

3 mole The mole is a quantity; it is used to “count” atoms and molecules

4 How we count atoms quickly l We have words that represent numbers that we use everyday. l FOR INSTANCE: –A dozen always = 12 items –A ream of paper always = 500 sheets of paper –A trio always = 3 of the item. SO IN CHEMISTRY………..

5 …. we have the mole 1 mole = 6.02 x particles AND 6.02 x is called Avogadro’s number

6 ê 1 mole of Al atoms = 6.02 x Al atoms ê 1 mole H 2 O molecules = 6.02 x H 2 O molecules ê 1 mole jelly donuts = 6.02 x jelly donuts ê 1 mole of = 6.02 x 10 23

x is a VERY large number

8 602,000,00 0,000,000,0 00,000,000

9 If Kevin took 1 mole of steps, each 15 cm long, he would walk around the earth 2.26 x times, which is 2,260,000,000,000,000 or 2,260 trillion times!

10 We cannot count 6.02 x Al atoms, it is just too too many.

11 We can, however, measure grams of Al.

grams of Al contains 6.02 x Al atoms!

grams of Al contains 1 mole of Al atoms!

14 The formula weight of any substance = mass of 1 mole of that substance in grams. (26.98 grams is the formula weight of Al.) So grams of Aluminum is ONE mole of Aluminum.

15 The formula weight of any element is the weight of an atom/molecule expressed in amu’s. (26.98 amu is also the atomic mass of Al)

16 More about amu l While amu is a valid unit for expressing weights of atoms and molecules, we are not going to be using it that often. l The reason is that it is the mass of a single unit and we are not able to use a single unit, since it is too small. l Know the definition of amu and all is well.

17 The formula weight of any compound is the sum of the formula weights of all the atoms in that compound, expressed in grams. (The formula weight of H 2 O = 18 grams per mole. (g / n ))

18 Identical Terms l The following terms on the next slide are often interchangeable. –All can be expressed in grams or amu depending on the application. –The term mass or weight is also often interchangeable. Since all of our calculations will be done on Planet Earth, the gravity is constant so we can make this simplification.

19 Here are the terms. l Formula weight: The most general of all the terms on this list. l Atomic weight (or mass): The weight of a mole of a single element. l Molecular weight (or mass): The weight of a mole of a single compound. l Molar mass: The weight of one mole of material. Just a rephrasing of above.

20 Molar conversions: The magic formula Mass in grams of an element or compound divided by molar mass = number of moles. We can make this into a formula and this is without a doubt the MOST IMPORTANT FORMULA IN CHEMISTRY.

21 The MAGIC FORMULA l The formula referred to on the earlier slide is: l Moles = Grams / Formula Weight OR n = g / M This is not a formula – This is THE FORMULA !!

22 Quantity conversions: Moles to number of particles Number of atoms or molecules divided by Avogadro’s number = number of moles. Set up a proportion !! We can all cross-multiply 1 mole = x moles 6.02 x y particles

23 Practice Fill out the following table in your notebook. Number of particlesMoles 4.56 x 10 25___________ 7.66 x 10 22___________ _____________.234 _________ 34.56

24 Practice Answers Fill out the following table in your notebook. Number of particlesMoles 4.56 x x x x

25 Moles to Liters l If the material is a gas, we can use the fact that 1 mole of gas = 22.4 Liters to find its volume. l We should set up a proportion. 1 mole = x moles 22.4 L y Liters

26 Conversion factor examples: 1 mol Al 6.02 x atoms Al 6.02 x atoms Al1 mol Al 18 g H 2 O1 mol H 2 O 1 mol H 2 O18 g H 2 O 180 g C 6 H 12 O x molecules 6.02 x molecules 180 g C 6 H 12 O 6

27 Using conversions: 2 mol Al x 6.02 x atoms Al = x atoms Al 1 mol Al 48 g H 2 O x 1 mol H 2 O = 2.7 mol H 2 O 18 g H 2 O 100 g C 6 H 12 O 6 x 6.02 x molecules = 3.34x molecules C 6 H 12 O g C 6 H 12 O 6

28 Subscripts in chemical formulas can help us represent moles! Since 1 molecule of CO 2 contains 1 atom of C and 2 atoms of O, then… 1 dozen molecules of CO 2 contain 1 dozen atoms of C and 2 dozen atoms of O, and… 1 mole CO 2 contains 1 mole C and 2 moles O ! This is a very useful relationship.

29 Molarity l Concentration of a solution, i.e. the amount of solute that is dissolved in a given amount of solvent l number of moles of solute per liter of solution l Units = mol/L OR n / L

30 Percent Composition l Percent that each part of a substance is of the whole, by mass. The whole mass is always the formula weight ! l The sum of the percents must = 100 % l % part = (mass of part / mass of whole) x 100 % l The percent composition of CO 2 is as follows… Wt. of C = 12g Wt. of O = 16g FW = 44 g % C = (12 / 44) x 100 % = % % O = ((2 x 16 ) / 44) x 100 % = %

31 Empirical Formula l Formula expressing the simplest whole number ratio of atoms in a compound l The empirical formula is often not the TRUE formula. l It is used to identify a compound in situations like forensics or learning about a new compound.

32 More on Empirical Formula l C 6 H 12 O 6 has an empirical formula CH 2 O l Found by calculating the ratio of the moles of each element in the compound. l STEPS (also found on the website) –Convert each given to number of moles. l Assume a 100 gram sample if needed. –Divide by the lowest number of moles to get a ratio. –Manipulate the values so they are whole numbers.

33 Molecular Formula l Determining the True or Actual Formula of the compound. l The empirical formula is needed first. l The weight of the molecular formula is given. –It will always be a multiple of the empirical formula. l The molecular formula is a multiple of the empirical formula. l CH 2 is the empirical formula for all of the following: C 2 H 4 C 3 H 6 C 4 H 8 C 6 H 12 and many more