stoichiometry: the study of quantitative measurements in chemical formulas and reactions chemistry...
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Stoichiometry:
The study of quantitative measurements in chemical
formulas and reactions
Chemistry – Mrs. Cameron
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Review of Chemical Equations
• The Law of Conservation of Mass applies to all chemical equations.
• All equations must be balanced to have the same number and type of atoms on both the product and reactant sides of an equation.
• Equations are balanced by adding coefficients in front of compound formulas in chemical equations.
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Consider the reaction of ethylene with oxygen:
C2H4 + 3O2 → 2CO2 + 2H2O
Can be read as:
1 molecule C2H4 + 3 molecules O2 → 2 molecules CO2 + 2 molecules H2O
OR:
1 mole C2H4 + 3 moles O2 → 2 moles CO2 + 2 moles H2O
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C2H4 + 3O2 → 2CO2 + 2H2O
• The coefficients in a chemical equation provide the ratio in which moles or molecules of one substance react with moles or molecules of another.
• The coefficients in a chemical equation provide the ratio in which moles or molecules of reactants relate to the moles or molecules of the product(s).
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C2H4 + 3O2 → 2CO2 + 2H2O
If the coefficients represent moles rather than molecules, everything in the equation is enlarged by a factor of Avogadro’s #.
For example:
Every 3 moles of O2 require 1 mole of C2H4
Every 3 moles of O2 produce 2 moles of CO2
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C2H4 + 3O2 → 2CO2 + 2H2O
• The Coefficients provide the relational quantities of reactant(s) and product(s).
• Within an equation, they can be used to form a “molar ratio” between any two substances in the reaction.
• Every 3 moles of O2 produce 2 moles of
CO2
22
23
COmole
Omole
23
22
Omole
COmoleOR
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C2H4 + 3O2 → 2CO2 + 2H2O
• Stoichiometry uses larger or smaller quantities – like cutting recipe in half or doubling or tripling a recipe
• Stoichiometry allows for calculations of multiples of the standard equations.
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(63.4 g O2)
20.32
21
Og
Omole
23
22
Omole
COmole
21
20.44
COmole
COg = 58.1 g CO2
Molar Ratio
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Calculating volume of a gas from a mass of another reactant or product.
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Simplification of Volume-Volume Problems
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Applications of Stoichiometry Problems
• Limiting Reactants– The limiting reactant “limits” the amount of
product that can be formed
– It is related to the molar ratio like the measurements in a recipe
– The limiting reactant is totally consumed or “used up” in a chemical reaction
• Excess Reactants - the reactant left over.
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N2(g) + 3H2(g) 2NH3(g)
•In this reaction, nitrogen reacts with hydrogen to form ammonia.
•For every one molecule (or mole) of nitrogen, you need 3 molecules (or moles) of hydrogen.
•Since there are only enough hydrogen molecules (3 sets of 3) to make 3 “recipes” of the ammonia, and not enough to make another, the hydrogen is the limiting reactant.
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N2(g) + 3H2(g) 2NH3(g)
•In this reaction, nitrogen reacts with hydrogen to form ammonia.
•For every one molecule (or mole) of nitrogen, you need 3 molecules (or moles) of hydrogen.
•Since there are two molecules of nitrogen left over, it is the excess reactant.
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Calculating Limiting Reactant Problems
• Balance the equation
• Complete a mass-mass problem for each reactant and the same product.
• Whichever reactant makes the least product is the limiting reactant.
• Whichever reactant makes the most product will have some leftover, or is the excess reactant.
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Example:C2H4 + 3O2 → 2CO2 + 2H2O
Ethylene reacts with oxygen to produce carbon dioxide and water. If 8.0 g of ethylene and 16.3 g of oxygen react, how many grams of water can be produced?
42
42
HC0.28HC1
gmole
42
2
HC12mole
OHmoles
20.32
21
Og
Omole
8.00 g C2H4
OHmoleOHg
2
2
10.18
2
2
32
OmolesOHmoles
OHmoleOHg
2
2
10.18 16.3 g O2
•Since the oxygen produces less product, it is the limiting reactant. The ethylene is the excess reactant.
•From this basic calculation, a chemist could also calculate how much ethylene was left over.
= 10.3 g H2O
= 6.11 g H2O
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Percent Yield Problems
• 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.
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Percent Yield Problems
In the reaction above, it was found that only 5.20 grams of water were produced. What is the percent yield of this reaction?
g
g
11.6
20.5X 100 % = 85.1% yield
An example: