Download AP Chemistry Chapter 3 Notes: Stoichiometry 3.1Counting by

AP Chemistry
Chapter 3 Notes: Stoichiometry
3.1Counting by Weighing
Atoms can be counted based on their average mass.
3.2 Atomic Masses
A. C-12, the Relative Standard
1. C-12 is assigned a mass of exactly 12 atomic mass units (amu)
2. Masses of all elements are determined in comparison to the C-12 atom
3. Comparisons are made using a mass spectrometer
B. Atomic Mass (average atomic mass, atomic weight)
1. Atomic masses are the average of the naturally occurring isotopes of an
element based on their relative abundances
For carbon sample with 98.89% C-12 and 1.11% C-13:
0.9889(12 amu) + 0.0111(13.0034 amu) = 12.01 amu
2. Atomic mass does not represent the mass of any actual atom
3.3 The Mole
A. Avogadro's number
1. 6.022 x 1023 units = 1 mole
2. Named in honor of Avogadro (he did NOT discover it)
B. An element's atomic mass (expressed in grams) contains 1 mole of atoms of
that element
1. 12.01 grams of carbon is 1 mole of carbon (multiple isotopes)
2. 12 grams of carbon-12 is 1 mole of carbon-12
3.4 Molar Mass
A. The mass in grams of one mole of a compound (= molar mass)
B. The sum of the masses of the component atoms in a compound
1. Molar mass of ethane (C2H6):
Mass of 2 moles of C = 2(12.01 g)
Mass of 6 moles of H = + 6 (1.008 g)
30.07 g
C. Fundamental units (individual particles) These masses are usually amu:
1. atom (of an element) = atomic mass
2. molecule (of a covalently bonded compound) = molecular mass
3. formula unit (usually of an ionic compound) = formula mass
3.5 Problem Solving
A. Ask yourself "What am I given in the problem?"
B. Then ask "What do I want to end up with in this problem?"
C. The third step is mapping out how you will get to what you want to find.
3.6 Percent Composition of Compounds
A. Mass percent composition = 100% x mass of substance (within the formula)
compound mass
"The part, divided by the whole, multiplied by 100"
Percent composition of oxygen in CO2:
O mass
2 x 16 amu
100 x CO mass = 100 x (2x16 amu)+ 12.01 amu = 72.7% O
2
3.7 Determining the Formula of a Compound
A. Empirical formulas
1. Convert quantities of each element to moles of the element.
(Treat % as grams, and convert to moles)
2. Divide each result by the smallest value to put into the smallest whole
number ratio of atoms
3. If the ratio is not all whole numbers, multiply each by an appropriate
integer so that all elements are in a whole number ratio
B. Molecular formula
1. Find the empirical formula mass
2. Divide the known molecular mass by the empirical formula mass.
3. Multiply the empirical formula by this result to get the molecular formula
3.8 Chemical Equations
A. Chemical reactions
1. Reactants are listed on the left hand side of the arrow.
2. Products are listed on the right hand side
3. From the Law of Conservation of Mass: the number of each element on
the reactant side must equal the numbers of that element on the
product side.
B. The Meaning of a Chemical Reaction
1. Physical States
a. Solid - (s)
b. Liquid - (l)
c. Gas - (g)
d. Dissolved in water (aqueous solution) - (aq)
2. Relative numbers of reactants and products
Coefficients give the atomic/molecular/mole ratios
3.9 Balancing Chemical Equations
A. Determine what reaction is occurring
List the reactants on the left side of the arrow and products on the right
B. Write the correct formulas for all the compounds
C. Balance the equation
Begin with elements that occur in only one substance on each side.
Save elements that occur more than once per side for the end.
Double as needed when diatomic elements are involved.
D. Include phase information
2H2 (g) + O2 (g) 2H2O (l)
3.10 Stoichiometric Calculations: Amounts of Reactants and Products
Based on the quantity of one substance in a reaction you can calculate quantities
of other substances involved in the reaction.
A. Balance the chemical equation
B. Molar ratios (coefficients of unknown / known from balanced equation) are
the key to stoichiometry. Before you use these, all quantities must be
converted to moles
C. Moles of the known x molar ratio = moles of the unknown (desired) substance
D. Convert from moles to the desired units
3.11 Limiting Reagents
A. The limiting reactant (reagent) limits the amount of product that can form since
it is the first reactant to be used up (completely react).
B. Solving limiting reactant problems
1. Determine amount of expected product based on given quantity of first
reactant.
2. Repeat with given amount of second reactant.
3. The reactant producing the smallest amount of product is the limiting
reactant. The smallest answer is the expected amount of product.
4. Excess (unreacted) reagents can be found by multiplying moles of
expected product by molar ratio (excess reactant / product). The
answer is the quantity that reacts. Subtract this from the initial quantity
given in the problem.
C. Calculating Percent Yield
1. Actual yield - what you got by actually performing the reaction
2. Theoretical (expected) yield - what stoichiometric calculation says the
reaction should have produced
Actual Yield
x 100% = percent yield
Theoretical Yield