Download Using Chemical Equations

Section 9.1
Using Chemical Equations
Section 9.1
Using Chemical Equations
1984, Union
Carbide Plant –
10000+ killed in
Bhopal, India
by Methyl
Isocyanate
1988 Pepcon,
Henderson Rocket
Fuel Explosion
Section 9.1
Using Chemical Equations
Objectives
1. To understand more of the information given in a
balanced equation
2. To use a balanced equation to determine relationships
between moles of reactant and products
Section 9.1
Using Chemical Equations
A. Information Given by Chemical Equations
• A balanced chemical equation gives relative numbers of
reactant and product molecules (or moles) that participate
in a chemical reaction.
• The coefficients of a balanced equation give the relative
numbers of molecules (or moles).
CO(g) + 2H2(g)
CH3OH(l)
How many moles of CH3OH will be produced from 1, 2, or
5 moles of CO?
How many moles of hydrogen will react with 1, 2, or 5
moles of CO?
Section 9.1
Using Chemical Equations
B. Mole-Mole Relationships in a BALANCED Reaction
• A balanced equation can predict the moles of product
that a given number of moles of reactants will yield.
What number of moles of O2 will be produced by
the decomposition of 2, 1, or 5.8 moles of water?
Section 9.1
Using Chemical Equations
B. Mole-Mole Relationships using Factor Analysis
• The mole ratio allows us to convert from moles of one
substance in a balanced equation to moles of a second
substance in the equation. Use factor analysis and mole
ratios to calculate the answers below:
CH4(g) + 2O2(g)
CO2 + 2H2O(l)
4.2 moles of CH4 reacts with how many moles of O2?
4.2 moles of CH4 will produce how many moles of H2O?
(WOC P310 Q8, 10)
Section 9.1
Using Chemical Equations
Objectives
1. To learn to relate masses of reactants and products in a
chemical reaction
2. To perform mass calculations that involve scientific notation
Moles
Grams
Moles
Grams
Section 9.1
Using Chemical Equations
A. Mass Calculations
If 9.0 grams of water are
decomposed to hydrogen and
oxygen, what mass of oxygen
gas is formed?
(WOC p311 Q15,16)
Section 9.1
Using Chemical Equations
C. Mass Calculations: Comparing Two Reactions
• For equal masses of methane (CH4) and propane
(C3H8), which will use more oxygen when fully burned?
"Happy Mole Day to You" Chemistry Song
Section 9.1
Using Chemical Equations
Double Bacon Cheeseburgers and Stoichiometry
• 1 Double Bacon Cheeseburger needs 1 bun, 2 patties, 2 slices of
cheese, 4 strips of bacon
• For 5 Double Bacon Cheeseburgers how many units of each
ingredient do I need?
• Fill in the final column below with how many complete burgers I can
make:
Buns
Patties
Cheese
Slices
Bacon
Strips
2
4
4
8
4
4
8
16
16
30
32
48
How Many
Burgers?
Section 9.1
Using Chemical Equations
Objectives
1. To understand the concept of limiting reactants
2. To learn to recognize the limiting reactant in a reaction
3. To learn to use the limiting reactant to do stoichiometric
calculations
4. To learn to calculate percent yield
Section 9.1
Using Chemical Equations
A. The Concept of Limiting Reactants
• Stoichiometric mixture
– N2(g) + 3H2(g)  2NH3(g)
Section 9.1
Using Chemical Equations
A. The Concept of Limiting Reactants
• Limiting reactant mixture
– N2(g) + 3H2(g)  2NH3(g)
Section 9.1
Using Chemical Equations
A. The Concept of Limiting Reactants
• For a Limiting reactant mixture the number of moles are
not balanced to match the reaction equation
– N2(g) + 3H2(g)  2NH3(g)
– Limiting reactant is the
reactant that runs out
first
– When the limiting
reactant is exhausted,
then the reaction stops
Limiting Reactants Game
Section 9.1
Using Chemical Equations
B. Calculations Involving a Limiting Reactant
Section 9.1
Using Chemical Equations
B. Calculations Involving a Limiting Reactant
Section 9.1
Using Chemical Equations
C. Percent Yield
• Theoretical Yield
– The maximum amount of a given product that can be
formed when the limiting reactant is completely
consumed.
• The actual yield (amount produced) of a reaction is usually
less than the maximum expected (theoretical yield).
• Percent Yield
– The actual amount of a given product as the percentage
of the theoretical yield.
Section 9.1
Using Chemical Equations
According to her pre-lab theoretical yield calculations a
student’s experiment should have produced 1.44g of
magnesium oxide. When she weighed her product after
reaction, only 1.23g of magnesium oxide was present.
What is the student’s percent yield?