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The Mole
Chapter 11
By: Kathryn Johnson 2008
& BuggLady 2006
What is a mole?
Not that one….
A Mole is….. (p.310)
… an SI unit of measure used for large numbers of particles
Like the following: 1 dozen = 12 1 gross = 144 1 ream = 500 1 thousand = 1,000 1 billion = 1,000,000,000
A Mole (mol) =
= 602,213,670,000,000,000,000,000 particles
= 6.02x1023 particles (scientific notation)
= 6.02E23 particles (calculator notation)
Particles - atoms, molecules, formula units, ions,…
This number is also called ‘Avogadro's number’
Why is the number so large?
Because atoms are so small!
Why do we use the mole?
Because it is easier to talk of 4 moles than to say we have 2,410,000,000,000,000,000,000,000 or 2.41x1024 molecules of something
We use it the same way we talk of
miles instead of feet when we travel
Remember…Conversion Factors (Chapter 2)
They allow us to decide whether to divide or multiply by a value
Also called: Dimensional Analysis Unit Cancellation Factor Label
How to set one up…
Set up a fraction Place the units first
the unit you have on the bottom the other unit from the relationship on top
Scratch off (Cancel) units that are on both the top and the bottom
Repeat until you only have the units you want Then place the numbers from the relationships
into the fractions
Last…Do the Math
Multiply all the numbers on the top together
Divide by all the numbers on the bottom
Round off to correct significant figures
USE YOUR CALCULATOR!
Mole ↔ Molecules (p.311)
Set up a conversion factor!
Moles x Molecules = Molecules
Moles
Molecules x Moles = Moles
Molecules
1 mol = 6.02E23 molecules
6.02E23 molecules
1 mol
multiplication
1 mol .
6.02E23 molecules
division
Be Careful!
This will be easy at first and you will want to skip steps….
But we will be setting up lots of these fractions and you have to keep them straight
Now, lets add another step!
Molar Mass is
the mass (g)
of one mole (mol)
of a substance
Molar Mass of Elements (p.313)
The mass (g) of a mole (mol) of a particular element is equal to its Atomic Mass on the Periodic Table Use two decimal places
Carbon: 1 mol C = 12.01 g C
12.01 g/mol
Molar Mass of Compounds (p.322)
The mass (g) of a mole (mol) of a molecule is found by adding the atomic masses of all the elements, multiply by the subscripts
Water… H2O
2(1.01) + 16.00 = 18.02 g/mol or 1 mol H2O = 18.02 g
Be Careful….
Polyatomic ions with parenthesis complicate the math!
Ammonium Phosphate… (NH4)3PO4
= 3(14.01+4(1.01)) + 30.97 + 4(16.00)= 3(14.01+4.04) + 30.97 + 64.00
= 3(18.05) + 30.97 + 64.00 = 54.15 + 30.97 + 64.00 = 149.12 g/mol
Another way to do it…
Ammonium Phosphate… (NH4)3PO4
N: 3(14.01) = 42.03
H: 4∙3(1.01) = 12.12
P: (30.97) = 30.97
O: 4(16.00) = 64.00
Molar Mass = 149.12 g/mol
Mass ↔ Mole (p.314)
Now we have a new conversion factor!
Mass x Mole = Moles
Mass
Moles x Mass = Mass
Moles
1 mol = (Molar Mass) g
(Molar Mass) g
1 mol
multiplication
1 mol .
(Molar Mass) g
division
Mass ↔ Molecules (p.316)(the 2 step!)
Now we set up two conversion factors! You always have to go to Moles first!
Mass x Mole x Molecules = Molecules Mass Mole
Molecules x Mole x Mass = Mass Molecules Mole
Ratio of Moles of an Elementin 1 mole of a Compound (p. 320)
equal to the subscript beside the element
Water… H2O
2 mol H and 1 mol O .
1 mol H2O 1 mol H2O
It gets complicated…
Watch out for those polyatomic parenthesis!
Ammonium Phosphate… (NH4)3PO4
3 mol N 12 mol H .
1 mol (NH4)3PO4 1 mol (NH4)3PO4
and
1 mol P 4 mol O .
1 mol (NH4)3PO4 1 mol (NH4)3PO4
Mass Compound ↔ Element (the 3 step!)
Now we set up many conversion factors! Always go to moles!
Mass compound x Mole compound x Mole element… Mass compound Mole
compound
… x Mass element = Mass element
Mole element
Yes… it keeps going!
You can set up these conversion factors in infinite combinations to transform from one value to another
The only trick is… to go from one ‘thing’ (compound) to another ‘thing’ (element in that compound) you have to go through the mole ratio based on the subscript
... Don’t Skip Steps!
Moles of Element
Mass of Element Atoms of Element
Moles of Compound
Mass of Compound Molecules of Compound
(subscript) mol of Element = 1 mol of Compound
(atomic mass) g = 1 mol 1 mol = (6.02E23) atoms
(molar mass) g = 1 mol 1 mol = (6.02E23) molecules
And we still have it easy…
wait until we start dealing with
Chemical Reactions!(Chapter 12)
Percent Composition (p.328)
percent (%) by mass of each element in a compound Review: Name → Formula (Chapter 8&9) Watch out for those polyatomic parenthesis!
% Element = (subscript)x(Atomic Mass)x100Molar Mass Compound
All the % should add up to 100!
Example: Water…. H2O
% H = 2(1.01) g H x 100 = 11.2%
18.02 g H2O
% O = 16.00 g O x 100 = 88.8%
18.02 g H2O
Check your work:
11.2% + 88.8% = 100%
Another way to do it…
Ammonium Phosphate… (NH4)3PO4
N: 3 (14.01) = 42.03 → 28.2%
H: 12 (1.01) = 12.12 → 8.1%
P: (30.97) = 30.97 → 20.8%
O: 4 (16.00) = 64.00 → 42.9%
Molar Mass = 149.12 g/mol
Practice
Determine the percent by mass of each element in calcium chloride.
Calculate the percent composition of sodium sulfate.
Which has the larger percent by mass of sulfur: H2SO4 or H2S2O8?
Empirical Formula (p.331)
The smallest whole-number mole ratio of the elements in a compound
Derived from %comp of the elements Turn Percent Composition (%) into Mass (g) Convert Mass→Mole with the Molar Mass Divide by the smallest value to get the whole
number ratios
Like a %comp in reverse
Example: 59.95% O & 40.05% S
0: 59.95 g x 1 mol = 3.747 mol = 3
16.00 g 1.249 mol S: 40.05 g x 1 mol = 1.249 mol = 1
32.07 g 1.249 mol
Therefore: SO3… Sulfur trioxide!
If the ratios are still decimals, multiply them by a small factor to make them whole (EX: 0.5 = ½, multiply by 2) (EX: 0.33 = 1/3, multiply by 3) (EX: 0.25 =1/4, multiply by 4)
Example: 48.64% C, 8.16% H, & 43.20% O
0: 43.20 g x 1 mol = 2.700 mol = 1 x 2 = 2
16.00 g 2.700 mol C: 48.64 g x 1 mol = 4.050 mol = 1.5 x 2 = 3
12.01 g 2.700 mol H: 8.16 g x 1 mol = 8.10 mol = 3 x 2 = 6
1.01 g 2.700 mol Therefore: C3H6O2
Two or more substances with different properties can have the same % composition and
empirical formula!
The simplest ratio in the empirical formula does not
indicate the actual number of elements in the compound
Molecular Formula (p.333)
• The actual numbers of atoms in a compound• Identifies the compound
• A multiple of the empirical formula• Molecular Formula = (Empirical Formula) n
n = Measured molar mass of compound_ _. Calculated molar mass of empirical formula
Example: 46.68% N & 53.32% OMeasured Molar Mass = 60.01 g/mol
0: 53.32 g x 1 mol = 3.333 mol = 1
16.00 g 3.332 mol N: 46.68 g x 1 mol = 3.332 mol = 1
14.01 g 3.332 mol Empirical Formula: NO … nitrogen monoxide
Molar Mass = 14.01 + 16.00 = 30.01 g/mol But, 60.01 g/mol = 2 therefore…
30.01 g/mol Molecular Formula: N2O2 … dinitrogen dioxide!
Review
Mol ↔ Molecules/Atoms (Avogadro's #) Mol ↔ Mass (Molar Mass) Mass ↔ Molecules/Atoms (The 2 step) Element ↔ Compound (Subscript, 3 step) % Composition Empirical/Molecular Formula
Practice…
Practice…
Practice…
Work Smart, Not Hard!
You have to be fast enough
to finish the test!
Mole Ratios
Coefficients indicate the number of moles Mole Ratios – the ratio between the numbers of moles of any
2 substances in a balanced equation
Ex: 2 Al + 3 Br2 2 AlBr3
All possible mole ratios: 2 mole Al : 3 mole Br2 2 mole Al : 2 mole AlBr3 3 mole Br2 : 2 mole Al 3 mole Br2 : 2 mole AlBr3 2 mole AlBr3 : 2 mole Al 2 mole AlBr3 : 3 mole Br2
Interpreting Mole Conversions
4 Fe + 3 O2 2 Fe2O3
You can interpret any balanced equation in terms of moles, mass, and molecules
Molecules: 4 atoms of iron react with 6 atoms of oxygen to produce 2 molecules of iron (III) oxide
Moles: 4 moles of iron react with 3 moles of oxygen to produce 2 moles of iron (III) oxide
Mass: convert to grams first
Practice Mole Ratios
For the following equations:Balance the equations
Interpret each in terms of moles, molecules, and mass
Determine all possible mole ratios 1. CaC2 + N2 CaNCN + C 2. NH3 + CuO Cu + H2O + N2
3. MnO2 + HCl MnCl2 + Cl2 + H2O 4. P4O10 + C P4 + CO 5. ZrI4 Zr + I2 6. PbS + O2 PbO + SO2
Mole to Mole Conversions Ex: Determine the number of moles of H2 produced when 0.040 moles
of K is used. Steps for solving: Step 1. Write the balanced equation:
2 K + 2 H2O 2 KOH + H2 Step 2. Identify the known (given) substances and the unknown
substances:Known: .040 moles KUnknown: moles of H2
Step 3. Determine the mole ratio of unknown substance to known substance:
1 mole H2 (use coefficients from balanced equation)2 mole K
Step 4. Convert using the known substance and the mole ratio:0.040 moles K x 1 mole H2 = 0.02 mole H2 produced
2 mole K
Practice
Handout: Mole to Mole Stoichiometry Problems
Mole to Mass Conversions Ex: Determine the mass of NaCl produced when 1.25 moles of Cl2 reacts with
sodium Steps: 1. Write the balanced equation:
2 Na + Cl2 2 NaCl 2. Identify the known (given) substances and the unknown substances:
Known: 1.25 moles Cl2Unknown: mass of NaCl
3. Determine the mole ratio of unknown substance to known substance:2 moles NaCl (use coefficients from balanced equation)1 mole Cl2
4. Convert using the known substance and the mole ratio:1.25 moles Cl2 x 2 moles NaCl = 2.5 moles NaCl
1 mole Cl2 5. Convert moles of unknown to grams of unknown: (use molar mass)
2.5 mole NaCl x 58.44 g = 146 grams NaCl 1 mole
Example: Mole to Mass
Reactants: 4 mole Fe x 55.85 g Fe = 223.4 g Fe 1 mole Fe
3 mole O2 x 32.00 g O2 = 96.0 g O2
1 mole O2
Product: 2 mole Fe2O3 x 158.7 g Fe2O3 = 319.4 gFe2O3
1 mole Fe2O3
Example, cont.
** Law of Conservation of Mass – the sum of the reactants should ALWAYS equal the sum of the product – check your work!!!
223.4 grams of Fe react with 96.0 grams of O2 to produce 319.4 grams of Iron (III) oxide
Mass to Mass Conversions Ex: Determine the mass of H2O produced from the decomposition of 25.0
grams of NH4NO3 Steps: 1. Write the balanced equation:
NH4NO3 N2O + 2H2O 2. Identify the known (given) substances and the unknown substances:
Known: 25.0 grams of NH4NO3
Unknown: mass of H2O 3. Convert grams of known to moles of known:
25.0 g NH4NO3 x 1 mole NH4NO3 = 0.312 mole NH4NO3 80.04 g
4. Determine the mole ratio of unknown substance to known substance:2 moles H2O (use coefficients from balanced equation)1 mole NH4NO3
5. Convert using the known substance and the mole ratio:0.312 mole NH4NO3 x 2 moles H2O = 0.624 mole H2O
1 mole NH4NO3 6. Convert moles of unknown to grams of unknown: (use molar mass)
0.624 mole H2O x 18.0 g = 11.2 grams H2O 1 mole
Hydrates (p.338)
Compound that has a specific number of water molecules (H2O) bound to its atoms
Solids in which water molecules are trapped Like a sponge, different amounts of water can
be trapped
Formulas and Names
1 formula unit of compound with a dot and then the attached water moleculesEx: Na2CO3∙10H2O
Name the molecule Use the prefixes in front of the
word ‘hydrate’Ex: sodium carbonate decahydrate
1 Mono-
2 Di-
3 Tri-
4 Tetra-
5 Penta-
6 Hexa-
7 Hepta-
8 Octa-
9 Nona-
10 Deca-
Anhydrous
“without water” Water is removed from the hydrate by
heating Desiccants are used to absorb excess
moisture in packages Calcium chloride Calcium sulfate
How much water attached?
How many moles of water attached to one mole of the molecule (BaCl2∙xH2O)
compare the mass before and after heating Mass hydrate – Mass anhydrous = Mass water
Convert the Mass water and Mass anhydrous to moles using the molar mass of each
Calculate the ratio x = mol H2O / mol BaCl2