Chapter 3

In the first two chapters we laid the foundation for what is to come in Chapter 3. We built this foundation based on observations in the laboratory and discussed how to interpret, calculate, and manipulated measured quantities. We also analyzed atoms, molecules, and compounds and discussed their properties.

Now it’s time to start observing what happens when we react chemical species together. How do we represent chemical reactions on paper? When two or more reactants are combined together, how can we determine the quantity of products produced? How can we use Stoichiometry to examine the quantity of substances produced and consumed in chemical reactions? This Chapter will answer these questions and then some as we begin to analyze chemical formulas and equations.

Chemical Reactions are represented by chemical equations. These equations list all the reactants and products in a chemical reaction.

3.1 Chemical Reactions

Once the reactants and products are determined for a chemical reaction, those products must be balanced in order to properly determine the amount of a reactant(s) consumed or the amount of the product(s) produced.

3.1 Balancing Chemical Reactions

Chemical reactions don’t just happen randomly. There is an underlying reason for why each reaction takes place and if we are to make a theory based on chemical reactivity we analyze various reactions and look at the trends. This allows a chemist to predict what might happen in a chemical reaction.

The reactions we will familiarize ourselves with in this section are: Combination Reactions, Decomposition Reactions, and Combustion Reactions.

3.2 Patterns of Chemical Reactivity

Our main goal in this chapter is to make quantitative calculations about the amount of reactants and products formed in a chemical reaction. In order to reach this goal, we must first examine the individual atoms and molecules that make up the overall chemical reaction. This starts with determining formula weights and molar mass.

3.3-3.4 Formula Weight and Molar Mass

Section 3.4, which discusses Avagadro’s number and the mole, was covered early in the Chapter 2 material. The mole is a practical unit with which we can analyze and calculate quantities in units that are familiar to us, such as grams.

3.4 Grams to Moles Example Problem

Using Avagadro’s number, we can interconvert between masses, moles, and numbers of particles.

3.4 Grams to Atoms Example

We can also determine the percent by mass of each component in a compound.

3.4 Percent Mass Example

Using the concept of the mole we can determine the empirical formula of a given compound. The empirical formula indicates the atoms in a substance in their smallest whole number ratio.

3.5 Empirical Formula

Once the concept of the empirical formula is established we can use experimental data to perform calculations.

3.5 Empirical Formula Calculation

If we have enough experimental data, we can determine the molecular formula of a substance from its empirical formula. We are often given the molecular weight or molar mass in addition to the empirical formula, allowing us to calculate the molecular formula.

3.5 Molecular Formula

Once the concept of the molecular formula is understood we can perform calculations after collecting experimental data.

3.5 Molecular Formula Example Problem

Using the mole concept, we can convert between the relative numbers of molecules in a reaction to the masses of the reactants and products. Using stoichiometrically equivalent quantities, we can outline a procedure enabling us to calculate the exact amounts of reactants consumed or products formed in a reaction.

3.6 Stoichiometry

Let’s take a closer look at what is happening at the atomic level in chemical reactions.

3.6 Stoichiometry at the Atomic Level

When one reactant is used up before the other reactants, a chemical reaction will not continue forward, leaving excess reactants that go unreacted. We can determine the amount of the excess reactants and the reactant that is used up, which is referred to as the limiting reactant.

3.7 Limiting Reactants

3.7 Excess Reactant Remaining Example

The theoretical calculations we determine on paper operate under the assumption that we maintain 100% efficiency throughout a chemical reaction. In the laboratory, 100% efficiency is rarely maintained. The amount of product recovered in the lab is called the actual yield. The percent yield of a reaction relates the actual and theoretical yields.

3.7 Percent Yield

Now that we have set the foundation for calculations involving chemical reactions, we can give many examples, including the two shown below, to test your understanding of these concepts.

Chapter 3 Stoichiometry Example #1

Chapter 3 Stoichiometry Example #2

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