1. The Mole Concept

2. Relative atomic mass Ar
Relative atomic mass is a weighted average of all the naturally occurring isotopes of an element. The standard upon which it is based is one-twelfth the mass of carbon-12. Since it is a relative mass, they are pure numbers and do not have units.

3. For molecules (covalently bonded), the term relative molecular mass is used (Mr).
For ionic compounds, use relative formula mass.
Both of these are calculated as sums of the relative atomic masses of the elements in the chemical formula.

4. Molar Mass
12 g of carbon-12 contains 6.02 x 1023 atoms of carbon-12. This number is known as Avogadro’s constant (L or NA). One mole of any substance contains 6.02 x 1023 representative particles. The molar mass (M) of any substance is calculated using the same numbers as relative atomic/molecular/formula mass with the units of g mol-1

5. Measuring Matter
There are three ways to measure matter: by counting representative particles (typically atoms, molecules or formula units), by mass (in grams), or by volume (in m3 or cm3 or dm3 for gases). The method used is usually chosen by the ease of each method and the information needed. Once a measurement has been made, it is possible to convert between the units for the other methods.

6. Measuring Matter 1 mole of a compound = molar mass (g mol-1)
1 mole of a monatomic element = x 1023 atoms
1 mole of a compound or diatomic = x 1023 representative particles
1 mole of a gas at STP = dm3
(273 K and 100 kPa)
The mole is the link between grams, the number of representative particles and liters!

7. Mole Road Map

8. Empirical & Molecular Formulas

9. Empirical Formula: a formula with the simplest whole number ratio of atoms (determined by experiment).
Molecular formula: a formula of a compound in which the subscripts give the actual number of each element in the formula

10. Here are the four formulas being used as examples:
Molecular Formula
Empirical Formula
H2O
C2H4O2
CH2O
C6H12O6
Notice two things:
1. The molecular formula and the empirical formula can be identical.
2. You scale up from the empirical formula to the molecular formula by a whole number factor.

11. Why chemists use it?
After a compound has been analyzed for percent composition, the formula can be calculated- the empirical formula. If the molar mass is also known then the molecular formula can be calculated as well.

12. How to Find Empirical Formulas
If given the percentages of each element, assume 100 grams of the substance and convert % into grams.
Convert to moles by dividing the amount in grams by the molar mass of that element.
Select the SMALLEST value and divide ALL values by this smallest one.
The results of Step 3 will either be VERY close to whole numbers or will be recognizable mixed number fractions If any result from Step 3 is a decimal mixed number, you must multiply ALL values by some number to make it a whole number. Ex: x 3, x 4, x 2, etc.

13. How to Find Empirical Formulas
5. Use these whole number results as SUBSCRIPTS and write the empirical formula, listing the elements in the order they are given in the problem. (HINT: don’t be surprised in the subscripts in some formulas are VERY large-many organic molecules are huge)

14. This can be summarized with the following poem
1. Percent to mass 2. Mass to mole 3. Divide by small 4. Multiply 'til whole

15. Molecular Formulas – are either the same as it’s experimentally determined empirical formula or it’s some whole number multiple of it. To determine the molecular formula, you must know the compound’s empirical formula AND the molar mass of the molecular compound.

16. How to Find Molecular Formulas
1.Calculate the mass of the empirical formula (which you have already found or it will be given to you )
2.Divide the known molar mass by the mass of the empirical formula.
3.Multiply that number by the subscripts of the empirical formula to get the subscripts for the molecular formula.

17. Map of Mole Land Mole Ratio Bridge Volume of a solution
Volume of a gas
Volume of a gas
1 mole = 22.4 L
Molarity = moles/Liters
Molarity = moles/Liters
1 mole = 22.4 L
Moles of THIS chemical
Moles of THAT chemical
Mole Ratio Bridge
Molar mass
Molar mass
6.02x1023 particles = 1 mole
6.02x1023 particles = 1 mole
Mass (g)
particles
Mass (g)
particles

18. Limiting and Excess Reagent
When amounts of more than one reactant are known you much first determine which is limiting and which is excess.
Limiting reagent: runs out first, determines the amount of products formed and other reactants used
Excess reagent: there will be some left over after the reaction goes to completion

19. Limiting and Excess Reagent
If two reactants are in a one-to-one mole ratio, convert each to moles. The smaller number of moles is the limiting reactant. Use the limiting reactant to determine amount of products formed and amount of other reactants used (can also be used to figure out how much is left over).

20. Limiting and Excess Reagent
If two reactants aren’t in a one-to-one ratio, use each reactant and stoichiometrically determine how much product would form. The smaller answer is the true answer and that reactant is limiting.