Unit 5: Chemical Quantities and Reactions
Chapter 10 – Chemical Quantities, Chapter 11 – Chemical Reactions,
Ch. 12 Stoichiometry
Objectives:
1.
Demonstrate the three methods for measuring the amount of a substance – count, mass, and volume.
2.
Understand the concept of a mole and what it represents.
3.
Convert between the mass of a substance, the number of moles that mass contains, and the number of
particles in that mass.
4.
Calculate the volume of a gas at STP.
5.
Demonstrate how chemical equations describe chemical reactions.
6.
Illustrate how to balance chemical reactions by changing coefficients.
7.
State and describe the five general types of reactions.
8.
Given the equation of a reaction, classify the reaction as one of the general types.
9.
Interpret balanced chemical equations and understand all the information present in them.
10. Perform mole-mole, mass-mass, mass-particles, volume-volume, and other stoichiometry calculations.
11. Determine the amount of product produced in a chemical reaction when taking into consideration a limiting
reagent.
12. Determine the percent yield of a chemical reaction.
Vocabulary:
Mole (noun)
Mol (unit)
Avogadro’s Number
Molar Mass (g/mol)
STP
Molar Volume (22.4L/mol)
Reactants
Products
Word equation
Chemical Equation
Coefficient
Combination / Synthesis Reaction
Decomposition Reaction
Single-replacement / displacement Reaction
Activity Series
Double-replacement / displacement Reaction
Combustion Reaction
Mole Ratio
Limiting Reagent
Excess Reagent
Theoretical Yield
Actual Yield
Percent Yield
10.1 The Mole: A Measurement of Matter
Read p. 287-296 and make your own notes by answering the following questions.
How do we measure amounts?
Stoichiometry is a method used to measure the amounts of substances involved in chemical reactions and relate
them to one another. It comes from Greek and Old English words meaning a series of steps for measuring
something. For example, a sample’s mass or volume can be converted to the number of particles in the sample.
Since atoms are so small, in order to count how many are in a sample, we group them into larger units. The group
or unit of measurement used to count number of atoms, molecules, or formula units of substances is the mole,
abbreviated mol. It is most convenient to define the mole as the number equal to the number of carbon atoms in
exactly 12 grams of pure carbon-12. The number of things in one mole is 6.02 x 1023, which is known as
Avogadro’s number. The mass of one mol of an element is called the molar mass, where 1amu is equivalent to
1mol. So the mass of 1mol of carbon is 12g since the atomic mass of carbon is 12amu. The units for molar mass
are grams per mol, g/mol.
1 dozen of donuts = 12 donuts
1 mole of donuts = 6.02 x 1023 donuts
If you had 2.80 x 1024 donuts, how many moles
would you have?
If you had 0.208 moles of donuts, how many donuts
would you have?
moles = particles x 1 mole / 6.02 x 1023 particles
particles = moles x 6.02 x 1023 particles / 1 mol
If each donut had 11 M&M candies on it, and you had 2.12 moles of donuts, how many M&M’s would you have?
What is the relationship between the atomic mass of an element and the mass of a mole of that element?
Molar mass = mass of 1 mole of a substance (g/mol)
What is the mass of 1 mole of nickel?
What is the mass of 1 mole of argon?
What is the mass of 1 mole of oxygen gas?
What is the molar mass of water?
How many moles would 275g of PCl3 contain?
10.2 Mole-Mass and Mole-Volume Relationships
Read p. 297-303 and make your own notes by answering the following questions.
mass = moles x mass / 1 mole
moles = mass x 1 mole / mass
What is the mass of 2.50 mol of iron (II) hydroxide?
How many moles does 75.9g of dinitrogen trioxide
have?
What was Avogadro’s hypothesis?
What is STP?
What is the relationship between number of particles and the volume of a gas, at STP?
volume = moles x 22.4 L / 1 mol
What is the volume of 3.70 mol of N2 at STP?
11.1 Describing Chemical Reactions
Read p. 321-328 and make your own notes by answering the following questions.
In general, what happens during a chemical reaction?
Describe how a word equation is written. Provide an example.
Describe how a chemical equation is written. Provide an example.
Summarize the meaning of the symbols in Table 11.1.
Describe how coefficients are used to balance equations and summarize the rules for writing and balancing
equations. Provide an example.
Do Practice Problems p.327 #3, 4; p.328 #5, 6; and Section Assessment p. 329 #10-12
11.2 Types of Chemical Reactions
Read p. 330-339 and make your own notes by answering the following questions.
Describe the five general types of reactions. Give a general and a specific example of each.
Explain the activity series and describe how and where it is used. Give an example.
Do Section Assessment p. 339 #24-26
12.1 The Arithmetic of Equations
Read p. 356-357 and make your own notes by answering the following question.
In a balanced chemical equation, what information can be gained from the coefficients?
12.2 Chemical Calculations
Read p. 359-366 and make your own notes by answering the following questions.
What is a mole ratio and when is it used?
EX. What mass of Fe2O3 is required to produce 1000.0 g of iron?
Fe2O3(s) + CO(g) → Fe(s) + CO2(g)
(MM of Fe2O3 is 159.70 & MM of Fe is 55.85 g)
Balance the chemical equation:
Solution: Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
mole ratio
1
mass ?
:
3
:
2
:
1000.0g
1000.0 g
55.85 g/mol
= 17.90 mol
3
Proper Format:
1000.0g Fe x 1mol Fe x 1 mol Fe2O3 x 159.70 g Fe2O3
55.85g
2 mol Fe
1 mol
= 1429.32 g of Fe2O3 is required to produce 1000.0 g of iron
17.90 mol
2
= 8.95 mol
= 8.95 mol x 159.70 g/mol
= 1429.32 g of Fe2O3 is required to produce 1000.0 g of iron
MASS-MASS CALCULATIONS: STEPS
1) Write the balanced chemical equation for the reaction.
2) Convert the given mass to moles.
3) Use the mole ratio to relate the known moles to the unknown moles.
4) Convert to mass/particles/volume of the unknown substance.
EX. How many molecules of oxygen gas are produced by the decomposition of 6.54g of KClO3?
12.3 Limiting Reactant and Percent Yield
Read p. 368-375 and make your own notes by answering the following questions.
Explain what determines how much of a product can be made from a chemical reaction.
EX. What mass of HCl is produced when 4.50 g of H2(g) and 140.0 g of Cl2(g) are reacted according to the
following equation? H2(g) + Cl2(g) → 2HCl(g)
Determine the number of moles of each reactant.
From H2
From Cl2
Calculate the number of moles of HCl expected from each reagent.
From H2
From Cl2
Which is the limiting reagent (the one producing the least number of moles)?
Determine the amount of product using limiting reagent.
Cl2 is the limiting reagent and H2 is in excess. Therefore, 288 g of HCl is produced.
EX. If 2.70 mol of C2H4 is reacted with 6.30 mole of O2, how many moles of water can be produced?
C2H4 (g) + 2O2 (g) →2CO (g) + 2H2O (g)
Most reactions do not produce the exact amount of product that is predicted by the balanced chemical equation.
You can calculate how many moles of a product would be expected under a certain set of conditions. However,
unexpected things occur in the real world of the laboratory. Some of the chemicals could be impure, so even
though you put in 5 g of a certain compound, there may actually be only 4.5 g of that compound present (0.5 g
would be impurities). Chemicals are sometimes spilled or inadvertently left on the sides of the containers. Other
times an alternate reaction occurs rather than the desired reaction. These real situations can lead to a lower than
expected yield of product.
The percent yield is the ratio of the actual yield to the theoretical yield. As with all ratios, as long as you are
consistent with your units, it does not matter which ones you use. It measures the efficiency of a reaction.
For example, suppose that you expected 500 g of a product, but actually only saw 250 g. The percent yield is
therefore 50%. Likewise, if 6 moles of product is expected and only 3 are produced, the percent yield is 50% in
this case as well.
In order to calculate the maximum expected amount, calculate how much product would be produced if all of the
limiting reagent were to react perfectly.
Theoretical Yield: the
maximum possible mass of
product that could be produced
in a chemical reaction.
CALCULATED
Actual Yield: the mass of
product actually obtained from a
reaction. MEASURED
Percentage Yield: a method of
expressing how efficiently a
reactant can be converted into a
product in a chemical reaction.
EX. Bromine was made according to the following reaction: HBrO3 + 5HBr → 3Br2 + 3H2O
If 10.0 g of HBrO3 was reacted with an excess of HBr and 26.3 g of Br2 was produced, what was the
percentage yield of the reaction?
Determine number of moles of reactant used.
Determine the theoretical yield of product (mole ratio, convert moles of product into grams).
Determine % Yield