Chemical Stoichiometry
Stoichiometry - The study of quantities
of materials consumed and produced
in chemical reactions.
Atomic Masses
Elements occur in nature as mixtures of isotopes
Carbon = 98.89% 12C
1.11% 13C
<0.01% 14C
Average atomic mass = (% of each isotope)(atomic mass of each isotope)
100
Carbon atomic mass = 12.01 amu
The Mole
The number equal to the number of carbon
atoms in exactly 12 grams of pure 12C.
1 mole of anything = 6.022  1023 units of
that thing
Molar Mass
A substance’s molar mass (molecular
weight) is the mass in grams of one
mole of the compound.
CO2 = 44.01 grams per mole
Percent Composition
Mass percent of an element:
mass of element in compound
mass % 
 100%
mass of compound
For iron in iron (III) oxide, (Fe2O3)
Formulas
molecular formula = (empirical formula)n
[n = integer]
molecular formula = C6H6 = (CH)6
empirical formula = CH
Empirical Formula Determination
1. Base calculation on 100 grams of
compound.
2. Determine moles of each element in
100 grams of compound.
3. Divide each value of moles by the
smallest of the values.
4. Multiply each number by an integer to
obtain all whole numbers.
Percent Composition/ Empirical
Formula Problem
1. An ion containing only oxygen and
chlorine is 31% oxygen by mass.
What is its empirical formula?
2. What is the percent composition by
mass of the elements in the compound
NaNO3?
Chemical Equations
Chemical change involves a
reorganization of the atoms in one
or more substances.
A representation of a chemical reaction:
C2H5OH + 3O2  2CO2 + 3H2O
reactants
products
Chemical Equation
C2H5OH + 3O2  2CO2 + 3H2O
The equation is balanced.
1 mole of ethanol reacts with 3 moles of
oxygen
to produce
2 moles of carbon dioxide and 3 moles of
water
Hydrogen and Nitrogen React to Form
Ammonia According to the Equation
N2 + 3H22NH3
Schematic Diagram of the Combustion
Device Used to Analyze Substances for
Carbon and Hydrogen
Combustion Stoichiometry Problem
A gaseous hydrocarbon sample is completely burned in air,
producing 1.80 liters of carbon dioxide at standard
temperature and pressure and 2.16 grams of water.
a. What is the empirical formula for the hydrocarbon?
b. What was the mass of the hydrocarbon consumed?
c. The hydrocarbon was initially contained in a closed 1.00
liter vessel at a temperature of 32oC and a pressure of
760 mmHg. What is the molecular formula of the
hydrocarbon?
d. Write the balanced equation for the combustion of the
hydrocarbon.
Calculating Masses of
Reactants and Products
1.
2.
3.
4.
Balance the equation.
Convert mass to moles.
Set up mole ratios.
Use mole ratios to calculate moles of
desired substituent.
5. Convert moles to grams, if
necessary.
Solving a Stoichiometry Problem with a
Limiting Reactant
The limiting reactant is the reactant that is
consumed first, limiting the amounts of products
formed.
1.
2.
3.
4.
5.
Balance the equation.
Convert masses to
moles.
Determine which
reactant is limiting.
Use moles of limiting
reactant and mole
ratios to find moles of
desired product.
Convert from moles
to grams.
Limiting Reactant Problem
2 Mg(s) + 2CuSO4 (aq) + H2O (l)  2 MgSO4 (aq) + Cu2O(s) + H2 (g)
a. If 1.46 grams of Mg(s) are added to 500. Milliliters
of a 0.200 molar solution of CuSO4, what is the
maximum molar yield of H2 (g)?
b. When all of the limiting reactant has been
consumed in (a), how many moles of the other
reactant (not water) remain?
c. What is the mass of the Cu2O produced in (a)?
d. What is the value of [Mg2+] in the solution at the
end of the experiment? (Assume that the volume
of the solution remains unchanged.)
Percent Yield- the actual yield of a product
as a percentage of theoretical yield
A chemist runs the reaction described below
2 Mg(s) + 2CuSO4 (aq) + H2O (l)  2 MgSO4 (aq) + Cu2O(s) + H2 (g)
The expected yield of Cu2O from a previous
problem was 4.29 grams. The chemist is
only able to collect 3.96 grams. What is the
chemist’s percent yield?
Molarity
Molarity (M) = moles of solute per volume
of solution in liters:
What’s one intrepretation of the label- “3 M HCl?”
Common Terms of Solution
Concentration
Stock - routinely used solutions prepared
in concentrated form.
Concentrated - relatively large ratio of
solute to solvent. (5.0 M NaCl)
Dilute - relatively small ratio of solute to
solvent. (0.01 M NaCl)
Solution
Chemistry
The Water
Molecule is
Polar
Polar Water Molecules Interact
with the Positive and Negative
Ions of a Salt
BaCI2 Dissolving
Preparation of a Standard Solution
A chemist whishes to prepare 1.00L of a 0.200 M
sodium hydroxide solution. Describe the steps,
with calculations, necessary to complete this task
starting with solid sodium hydroxide and distilled
water.
Dilution of SolutionsM 1 V1 = M 2 V2
(a) A
Measuring
Pipet
(b) A
Volumetric
(transfer)
Pipet
You’ve been asked to prepare 150 ml of a 0.035M
solution of sodium hydroxide from the 0.200M stock
sodium hydroxide solution prepared earlier. Detail
the steps necessary to complete this task.
Beer- Lambert Law Beer’s Law
Relates the amount of light
being absorbed to the
concentration of the
substance absorbing the light
A=abc
A = measured absorbance
a=
b=
c=
Beer’s Law Sample Problems
1. A solution with a concentration of 0.14M is measured to have an
absorbance of 0.43. Another solution of the same chemical is
measured under the same conditions and has an absorbance of
0.37. What is its concentration?
2. The following data were obtained for 1.00 cm samples of a
particular chemical. What is the concentration of a 1.00 cm
sample that has an absorbance of 0.60?
Conc.
Abs.
(M)
3. The absorptivity of a particular
chemical is 1.5/M·cm. What is the
concentration of a solution made from
this chemical if a 2.0 cm sample has an
absorbance of 1.20?
0.50
0.69
0.40
0.55
0.30
0.41
0.20
0.27
Beer’s Law Sample Problems
4. Using the data from the graphing example in question #2, what
are the concentrations of solutions with absorbances of 0.20,
0.33, and 0.47?
5. A solution is prepared to be 0.200M. A sample of this solution
1.00 cm thick has an absorbance of 0.125 measured at 470nm
and an absorbance of 0.070 measured at 550nm. Calculate the
concentrations of the following solutions:
Sample
1
2
3
4
Absorbance
0.055
0.155
0.120
0.048
Wavelength
470nm
470nm
550nm
550nm
Path length
1.00cm
1.00cm
1.00cm
5.00cm
Stoichiometry Problem Set
1.
2.
3.
Aluminum oxide is to be made by combining 5.00 g of
aluminum with oxygen gas. How much oxygen is needed in
moles? In grams? In liters?
During its combustion, ethane (C2H6) combines with oxygen
gas to give carbon dioxide and water. A sample of ethane
was burned completely and the water that formed had a mass
of 1.61 g. How much ethane, in moles and in grams, was in
the sample?
Chloroform, CHCl3, reacts with chlorine gas to form carbon
tetrachloride and hydrogen chloride. In one experiment the
reactants were initially presented in a ratio of 1 to 1 by mass;
specifically, 25.0 g of CHCl3 was mixed with 25.0 g of Cl2 (g).
Which is the limiting reactant? What is the maximum yield of
carbon tetrachloride in moles and in grams?
Stoichiometry Problem Set
4. One of the steps in one industrial synthesis of sulfuric acid
(H2SO4) from sulfur is the conversion of sulfur dioxide (SO2) into
sulfur trioxide (SO3) by this reaction:
2SO2+O2  2SO3
In one “run,” 1.75 kg of SO2 was used and 1.72 kg of SO3 was
isolated from the mixture of products. What was the percent
yield?
5. A student needs 0.250 mol of NaCl and all that is available is a
solution labeled “0.400 M NaCl.” What volume of the solution
should be used? Give your answer in milliliters.
6. Describe how to prepare 250 mL of 0.200 M NaHCO3.
Stoichiometry Problem Set
7. How many milliliters of 0.114 M H2SO4
solution are necessary to completely
neutralize 32.2 mL of 0.122 M NaOH?
8. Describe how to make 500 mL of 0.20 M
NaOH from 0.50M NaOH.