Chapter 3 – Stoichiometry
Section 3.1 – Counting by Weighing
-when measuring large quantities, it is often easier to measure by weight rather than counting
-because not every item being counted is identical, we rely on the concept of average mass
-this is how we count atoms and molecules in chemistry
Section 3.2 – Atomic Masses
-the modern system of atomic masses uses carbon-12 as the standard, with a mass of exactly
12 atomic mass units
-masses of other elements are determined by comparing them with carbon-12 using a mass
spectrometer
See figure 3.1-diagram of mass spectrometer
-average atomic masses are always used which are the weighted averages of all of the isotopes
of an element
-no atom actually has the mass of the average atomic mass, but the value allows us to do
stoichiometric calculations with relatively high accuracy
Section 3.3 – The Mole
-defined as the number equal to the number of carbon atoms in exactly 12 grams of pure
carbon-12
-Avogadro’s number=6.022x1023
-mole was defined as that number so that a sample of an element with a mass the same as the
element’s atomic mass only in grams would contain the above number of atoms
Section 3.4 – Molar Mass
-molar mass = mass of one mole of a compound
-found by adding together atomic masses of the components
Section 3.5 – Learning to Solve Problems
-the process is the most important part to understand about the example problems and
homework problems that you will do
-the first step is always analyzing what information is provided in the problem and what the
final goal is
-then, start making a map (figuratively or literally) of a plan to get from start to finish
-always make sure to check that your answer makes sense; this will save you for easy mistakes
Section 3.6 – Percent Composition of Compounds
-mass percent (or weight percent) of an element in a compound is found by comparing the
mass of that element in one mole of the compound to the total mass of one mole of the
compound
Section 3.7 – Determining the Formula of a Compound
-sample is usually analyzed by decomposing it to its component elements or simple compounds
that can then be weighed
-find the fraction by mass of each element (or percent composition)
-assume you are working with 100 g of the compound and find the smallest whole-number
ratio and therefore the empirical formula
-numbers that are very close to whole numbers should be rounded, otherwise, multiply each
number by an integer to create whole numbers
-able to use molar mass (if known) to find the molecular formula
See page 93 for step-by-step directions for determining empirical and molecular formulas
Section 3.8 – Chemical Equations
Chemical Reactions
-describes a chemical change
-law of conservation of mass tells us that all atoms must be accounted for on each side of the
equation
-bonds are broken or formed, or reorganized – nothing is created or destroyed
The Meaning of a Chemical Equation
-equation often gives the physical states of the reactants and products
-relative numbers are shown with coefficients
Section 3.9 – Balancing Chemical Equations
-can only change coefficients, cannot change the identities of the molecules (no subscript
changes)
-can usually solve with trial and error (“by inspection”)
-to balance a chemical equation:
-determine products and reactants, along with their states
-write the unbalanced equation
-balance using the most complicated molecule first
Section 3.10 – Stoichiometric Calculations: Amounts of
Reactants and Products
-chemical equations tell us the number of moles or molecules, not the masses, but
experimentally, masses are used
-use molar masses, mole ratios, and factor labeling in order to convert the mass of one
component to the mass of another
-to calculate masses of reactants and products:
-make sure the chemical equation is balanced
-convert the known mass of reactant or product to moles
-set up mole ratio using the balanced equation
-use the mole ratio to find the number of moles of the unknown reactant or product
-convert moles to grams with molar mass
Section 3.11 – The Concept of Limiting Reagent
-chemicals are often mixed so that they run out at the same time, also known as a
stoichiometric mixture
-sometimes one reactant is consumed before the other(s), it is known as the limiting reactant
(or limiting reagent)
-must figure out which reactant is limiting so that you know how much product can be made
-crucial to use moles, not mass, to compare
-amount of product that can be made when the limiting reactant is completely used in ideal
conditions is known as the theoretical yield
-this amount is never actually made (lab error, side reactions)
-percent yield-amount actually made (actual yield) compared to theoretical yield