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09: The Mole Key Chemistry Terms • • •
• • • • • •
Mole Conversion Examples
Mole: SI Mole: SI unit for counting (abbreviation: mol) Avogadro’s number: 6.02 number: 6.02 × 1023 Molar Mass: Mass Mass: Mass (in grams) for 1 mole of a substance (also called molecular mass, molecular weight, formula mass, formula weight) Standard temperature and pressure (STP): 0 (STP): 0°C (273K) and 1 atm Molar volume: the volume: the volume of gas for 1 mole at STP is 22.4 liters Percent composition: mass composition: mass of the element / mass of the whole sample × 100 Empirical: from Empirical: from data Empirical formula: lowest formula: lowest possible ratio of atoms in a molecule Molecular formula: actual formula: actual ratio of atoms in a molecule
How many molecules are in 0.25 moles? 0.25 mole 6.02 × 1023 = 1.51 × 1023 molecules molecules 1 mole How many grams are in 0.15 moles Fe2(SO4)3? 0.15 mol 399.91 g Fe2(SO4)3 Fe2(SO4)3 = 59.99 g Fe2(SO4)3 1 mol Fe2(SO4)3 How many liters is 4.45 × 1020 molecules at STP? 4.45 × 1020 1 mole 22.4 L molecules = 0.0166 L 6.02 × 1023 1 mole molecules
The mole
Percent Composition
The mole is a counting unit used to count very tiny particles (similar to “dozen”)
%
composition
=
mass element
× 100
mass whole
1 mole of anything = 6.02 × 1023 pieces The mass of the element and the whole can be given as in lab data, or they can be determined from the chemical formula (just as the molar mass is).
Mole relationships By knowing the relationship of the individual component to the whole thing, you can determine the amount of each individual component
Empirical Formula 1.
Examples: 1 mole of NaCl 1 mole of Na and 1 mole of Cl 1 mole of CaCl 2 1 mole of Ca and 2 moles of Cl
2. 3.
4.
Molar mass The atomic mass found on the periodic table is the mass (in grams) for 1 mole of atoms of that element. By adding the atom masses for atoms in a molecule, the molar mass of the molecule can be found.
If given percent’s, assume assume they are grams (percents add up to 100) Change all grams to moles Divide all moles by the smallest to get the lowest ratio (multiply by a factor if needed to make them whole numbers) Write the formula with the ratio as subscripts
Examples: Find the empirical formula if a sample contains Ca and Cl and is 36.1% Ca 36.1 g Ca 1 mole Ca = 0.901 mole Ca 40.08 g Ca
Be sure to distribute subscripts outside the parenthesis to each atom inside
63.9 g Ca
Examples: Find molar mass for Fe2(SO4)3 Fe 2 × 55.85 = 111.70 S 3 × 32.07 = 96.21 O 12 × 16.00 = 192.00 399.91 g/mole
1. 2. 3.
The identity of the gas doesn’t matter—just the number of particles (1 mole).
4.
Several equalities can be used for mole conversions: 1 mole of particles ↔ 6.02 × 1023 particles • mole of whole molecule ↔ moles of individual atoms • mole ↔ molar mass in grams • moles gas ↔ molar volume
RapidLearningCenter.com
0.901 mole Ca 0.901 mole
1 Ca
1.81 mole Cl 0.901 mole
2 Cl
Molecular Formula
At standard temperature and pressure (STP), 1 mole of any gas is 22.4 L.
•
= 1.81 mole Cl
CaCl2
Molar Volume
Mole Conversions
1 mole Ca 34.45 g Ca
Find empirical formula, if not given to you Find the molar mass of the empirical formula Find the ratio of the molecular formula’s molar mass (must be given to you) to the empirical formula’s molar mass. Multiple the empirical formula’s subscripts by the ratio.
Examples: If empirical formula = CH3, what’s the molecular formula if molecular mass = 30.08 g/mole? CH3 molar mass = 15.04 g/mole 30.08 / 15.04 = 2 Molecular formula = CH3 × 2 = C2H6