Chemistry
Chapter 10 Note Packet
Mr. Fedell
(student’s edition)
Chapter 10 problems: see handout!
Check out the sample problems in this chapter – they really show how to do all the work you
need to know how to do! Also check out the tables and figures – they help!
10.1 The Mole: A Measurement of Matter
Atoms are tiny, so we count them in “bunches” - a mole is a “bunch of atoms”
mole (definition) - The amount of a compound or element that contains 6.02 x 1023 particles of
that substance.
1 mole = 1 molar mass = 6.02 x 1023 particles
Avagadro’s Number = 6.02 x 1023
Molar Mass - the sum of the atomic masses of all atoms in a formula
Round to the nearest tenth! (measured in amu or grams)
ex - H2
H2O
Ca(OH)2
molar mass is a term that can be used for atoms, molecules (covalent compounds or elements)
and formula units (ionic compounds)
Official names may also be:
Formula mass (ionic compounds)
Molecular mass (covalent compounds and diatomic elements)
Atomic weight, Atomic mass, grams formula weight, etc.
examples:
1 mole Na =
1 mole O2 =
1 mole HCl =
1 mole NaCl =
atoms =
molecules =
molecules =
formula units =
10.2 Mole-Mass and Mole-Volume Relationships
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g
g
g
g
Step 1 – be able to calculate molar mass (aka – formula mass, molecular weight, atomic
weight, atomic mass, gram formula weight, etc.)
Add atomic weights from the periodic table – round to the nearest 10th place
ex. CH4 =
.
don’t forget hydrates! MgSO4 7 H20 =
Step 2 – be able to convert between grams, moles, particles, and liters
Mole Relationships - For starters in chemistry, we have to be able to convert between moles,
grams, and molecules/atoms of substance (also liters when we work with gases)
use the “mole map
_____ mole
_____ atom/molecule
_____ gram
_____ liter (at STP)
convert 13.8 g Li to moles
convert 2.0 moles Ne to g
convert 2.15g Mg to moles
convert 2.15g Mg to atoms
convert 1.55L of CO2 to molecules
convert 5.27 x 1019 molecules C6H12O6 to grams
convert 32.0 g of CH4 to moles, liters, molecules, total atoms, atoms of H
2
10.3
Percent Composition and Chemical Formulas
Lots of common names:
dry ice
epsom salts
milk of magnesia
chemical formulas show the relative #’s of atoms in a chemical compound
ex.
C12H22O11
Pb(NO3)4
(NH4)2CrO4
monatomic ions – end in “ide” – look up charges on periodic table
compare F (
)
to F-1
compare S (
)
to S-2
Percentage Composition - every compound has a certain percentage of each type of atom
(we measure it by mass)
% composition =
ex 1 - calculate % composition if a compound contains 24 g of Carbon and 64 g of Oxygen
ex 2 - what is the percentage composition of Ba(OH)2
empirical formula - the lowest whole number ratio of atoms in a compound (simplest formula)
3 basic types of problems - first, three rules
1. Divide (%’s or grams) by the gram atomic mass
2. Divide the resulting #’s by the smallest of those numbers
3. Multiply by 2 or 3 only if a whole number ratio isn’t the result of step 2
example 1 - calculate the empirical formula if there is 52.17 % C, 13.04% H, and 34.78 % O
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example 2 - calculate the empirical formula is there is 26.56 % K, 35.41 % Cr, and 38.03 % O
example 3 - Find the empirical formula if a sample contains 5.6 g N and 12.8 g O
Finding Molecular Formula - the same steps as empirical formula with one additional step - use
the gram formula mass of the empirical formula and its relationship to the gram formula mass of
the molecular formula to find what number to multiply the empirical formula by to find the
molecular formula (sounds more complicated than it is)
example 5 - Find the molecular formula of a compound (mw = 144.0 g) with 66.67 % C,
11.11 % H, and 22.22 % O
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