Honors ChemistryChapter 10 Notes – Chemical Quantities
(studentedition)
Chapter 10 problems: 48, 60-62, 72, 74, 76, 77, 80, 81
good figures in the book to look at.... In addition to the chapter 10 problems, complete the following assignment. There are seventeen figures in this chapter. In addition to your chapter problems, make a list of what you consider “good” figures to look at (figures that you feel are helpful in some way to your understanding of chemistry). Also make a list of “useless” figures (that category is fairly self explanatory). Include these with your chapter problems. I will collect all chapter problems and this “picture assignment” on separate loose-leaf paper on the day before the test.
10.1 The Mole: A Measurement of Matter
Stoichiometry:
Formula Mass: The sum of the atomic masses of all atoms(atomic masses) in a formula.
When calculating the atomic masses, round the periodic table mass to the nearest tenth!
Atomic masses are measured in .
Examples:H2H2OCa(OH)2
Molecular mass - same thing, but truly used for only.
Gram Atomic Mass/Gram Formula Mass: Same thing as above, but the unit is .
ex - H2O formula mass (or molecular mass) =
gram formula mass=
(also called gram molecular mass - also called gram molecular weight)
The Mole: The amount of a compound or element that contains particles ofthat substance.
Atoms are tiny, so we count them in “ ”.
A mole is a “ ”.
1 mole = 1 gram formula mass = particles
Element / Moles / Atoms/Molecules/Formula Units / MassNa / 1 / 6.02 x 1023 atoms / grams
O2 / 1 / 6.02 x 1023 molecules / grams
HCl / 1 / 6.02 x 1023 molecules / grams
NaCl / 1 / 6.02 x 1023 formula units / grams
10.2 Mole-Mass and Mole-Volume Relationships
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)
Oh yeah… These are dimensional analysis problems (factor label method). As far as sig figs go, what you start with, you end with!
2 steppers
convert 13.8 g Li to moles
convert 2.0 moles Ne to g
convert 3.0 moles of Be to atoms
convert 44.8 L of O2 to moles
3 and 4 steppers
convert 128 g of O2 to molecules of O2
convert 128 g of O2 to atoms of oxygen
if there are 9.71 x 1024 atoms of oxygen in a sample of SO3, how many grams of SO3? liters?
10.3 Percent Composition and Chemical Formulas
Percentage Composition - every compound has a certain percentage of each type of atom
(we measure it by )
% 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
8.9, 8.10 and more... Determining a compound’s empirical and molecular formula
empirical formula - the lowest whole number ratio of atoms in a compound (simplest formula)
3 basic types of problems - first, two 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
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
sometimes , the information provided isn’t always the same...
example 4 - an unknown compound decomposes into 1.134 g of I2 (s), and 100.0 cm3 H2(g)
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 (mm = 144.0 g) with 66.67 % C,
11.11 % H, and 22.22 % O
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