Chapter 7: Chemical Reactions and Chemical Quantities – Study Notes

Chapter 7: Chemical Reactions and Chemical Quantities

Introduction to Chemical Reactions

Chemical reactions involve the transformation of substances into new products through the breaking and forming of chemical bonds. Understanding the classification, balancing, and quantitative analysis of chemical reactions is essential in general chemistry.

• Chemical Change: A process where substances are converted into different substances with new properties. Example: Frying an egg.

• Physical Change: A process that does not alter the chemical composition of a substance. Example: Boiling water.

• Indicators of Chemical Change: Color change, gas production, formation of a precipitate, energy change.

• Examples: Food spoilage (chemical change), melting ice (physical change).

Balancing Chemical Equations

Balanced chemical equations ensure the conservation of mass and atoms. Each side of the equation must have the same number of each type of atom.

• Steps to Balance:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms.

  4. Check your work.

• Example: Hydrogen gas reacts with oxygen gas to form water. Unbalanced: $H_2 + O_2 \rightarrow H_2O$ Balanced: $2H_2 + O_2 \rightarrow 2H_2O$

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the reactants and products.

• Combination (Synthesis): Two or more substances combine to form one product. Example: $2H_2 + O_2 \rightarrow 2H_2O$

• Decomposition: A single compound breaks down into two or more simpler substances. Example: $2KClO_3 \rightarrow 2KCl + 3O_2$

• Single Replacement: One element replaces another in a compound. Example: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$

• Double Replacement: Exchange of ions between two compounds. Example: $AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$

• Combustion: A substance reacts with oxygen, releasing energy, usually as heat and light. Example: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Stoichiometry and Chemical Quantities

Stoichiometry involves the calculation of reactants and products in chemical reactions using balanced equations.

• Mole Concept: The mole is a unit for counting particles. $1 \text{ mole} = 6.022 \times 10^{23}$ particles.

• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

• Stoichiometric Calculations: Use the coefficients in balanced equations to relate moles of reactants and products.

• Example: How many grams of oxygen are needed to react completely with 7.80 moles of copper? Reaction: $2Cu + O_2 \rightarrow 2CuO$ From the equation, 2 moles of Cu require 1 mole of $O_2$. Calculate moles of $O_2$ needed for 7.80 moles of Cu: $\frac{1 \text{ mol } O_2}{2 \text{ mol } Cu} \times 7.80 \text{ mol } Cu = 3.90 \text{ mol } O_2$ Convert moles of $O_2$ to grams using molar mass ($32.00 \text{ g/mol}$): $3.90 \text{ mol } \times 32.00 \text{ g/mol} = 124.8 \text{ g}$

Limiting Reactant and Percent Yield

The limiting reactant is the reactant that is completely consumed first, limiting the amount of product formed. Percent yield compares the actual yield to the theoretical yield.

• Limiting Reactant: Identify by comparing the mole ratios of reactants to the balanced equation.

• Theoretical Yield: Maximum amount of product possible from the limiting reactant.

• Percent Yield Formula: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

• Example: If 4.2 g of product is obtained but the theoretical yield is 5.0 g, percent yield is: $\frac{4.2}{5.0} \times 100\% = 84\%$

Balancing and Interpreting Chemical Equations

Balanced equations provide the quantitative relationships between reactants and products. Each atom must be accounted for on both sides.

• Example Table: Balancing $KClO_3 \rightarrow KCl + O_2$

  Atom	On Left Side	On Right Side
  K	1	1
  Cl	1	1
  O	3	2

  To balance, adjust coefficients: $2KClO_3 \rightarrow 2KCl + 3O_2$

Applications and Problem Solving

Applying these concepts allows for solving quantitative problems in chemistry, such as determining the amount of reactant needed, identifying limiting reactants, and calculating percent yield.

• Example: Calculate the mass of sodium chloride that decomposes to form 6.4 g of oxygen. Reaction: $2NaClO_3 \rightarrow 2NaCl + 3O_2$ Use molar mass and stoichiometry to solve.

• Example: When 20.0 g NH3 and 50.0 g O2 are allowed to react, which is the limiting reagent? Reaction: $4NH_3 + 5O_2 \rightarrow 4NO + 6H_2O$ Calculate moles of each and compare to stoichiometric ratios.

Summary Table: Types of Chemical Reactions

Type	General Form	Example
Combination	$A + B \rightarrow AB$	$2H_2 + O_2 \rightarrow 2H_2O$
Decomposition	$AB \rightarrow A + B$	$2KClO_3 \rightarrow 2KCl + 3O_2$
Single Replacement	$A + BC \rightarrow AC + B$	$Zn + 2HCl \rightarrow ZnCl_2 + H_2$
Double Replacement	$AB + CD \rightarrow AD + CB$	$AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$
Combustion	$Fuel + O_2 \rightarrow CO_2 + H_2O$	$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Additional info:

• Practice problems in the file cover classification of changes, balancing equations, stoichiometry, limiting reactant, and percent yield, which are all core topics in General Chemistry Chapter 7.

• Students should be comfortable converting between grams, moles, and molecules using molar mass and Avogadro's number.