CHE 107 Exam Study Guide: Chemical Reactions, Solutions, and Acids/Bases (Ch. 9–11)

Chapter 9: Chemical Reactions

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the patterns of reactants and products. Recognizing these types is essential for predicting products and balancing equations.

• Synthesis (Combination): Two or more substances combine to form a single product. General pattern: A + B → AB Example: $2H_2 + O_2 \rightarrow 2H_2O$

• Decomposition: A single compound breaks down into two or more simpler substances. General pattern: AB → A + B Example: $CaCO_3 \rightarrow CaO + CO_2$

• Single Replacement: One element replaces another in a compound. General pattern: A + BC → AC + B Example: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$

• Double Replacement: Ions in two compounds exchange partners to form two new compounds. General pattern: AB + CD → AD + CB Example: $AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$ Tip: If two compounds react, it is likely a double replacement reaction.

Balancing Chemical Equations

Balancing equations ensures the conservation of mass, meaning the same number of atoms of each element are present before and after the reaction.

• Rules: Only change coefficients, never subscripts.

• Steps:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Balance one element at a time.

  4. Save hydrogen and oxygen for last.

• Example: $Na + H_2O \rightarrow NaOH + H_2$ Balanced: $2Na + 2H_2O \rightarrow 2NaOH + H_2$

Conservation of Mass

• The total mass of reactants equals the total mass of products.

• No atoms are lost or created during a chemical reaction.

Energy in Reactions

• Exothermic: Reactions that release heat energy.

• Endothermic: Reactions that absorb heat energy.

Chapter 10: Solutions & Concentration

Basic Definitions

• Solute: The substance dissolved in a solution.

• Solvent: The medium in which the solute is dissolved (often water).

• Solution: A homogeneous mixture of solute and solvent.

Molarity (M)

Molarity is the most common unit of concentration in chemistry, defined as moles of solute per liter of solution.

• Formula: $M = \frac{\text{moles}}{\text{liters}}$

• Rearranged:

  • moles = M × L

  • L = moles / M

• Example: 0.5 mol in 250 mL Convert 250 mL to 0.250 L. $M = \frac{0.5}{0.250} = 2.0$ M

Dilution

Dilution involves adding solvent to decrease the concentration of a solution. The relationship is given by:

• Formula: $M_1 V_1 = M_2 V_2$

• Key idea: Adding water decreases molarity (M) and increases volume (V).

• Example: 2.0 M, 1.0 L diluted to 4.0 L $(2.0)(1.0) = M_2(4.0)$ $M_2 = 0.5$ M

Electrolytes

Electrolytes are substances that produce ions when dissolved in water, affecting electrical conductivity.

Type	Behavior	Example
Strong	Fully dissociate	NaCl
Weak	Partially dissociate	Acetic acid
None	No ions produced	Sugar

Solubility

• Principle: "Like dissolves like"—polar solvents dissolve polar solutes, nonpolar solvents dissolve nonpolar solutes.

• Ionic compounds generally dissolve in water.

Types of Solutions

Type	Meaning
Unsaturated	Can dissolve more solute
Saturated	Maximum amount dissolved
Supersaturated	More than maximum dissolved; unstable

Chapter 11: Acids, Bases, and Buffers

Acids vs Bases

Acids and bases are fundamental chemical species with distinct properties and definitions.

Type	Definition	Example
Acid	H+ donor	HCl
Base	H+ acceptor	NaOH

• Acids typically start with H in their formula.

• Bases typically end with OH.

Strong vs Weak Acids and Bases

• Strong acids/bases: Fully dissociate in water. Examples: HCl, HNO3, NaOH, KOH

• Weak acids/bases: Partially dissociate.

pH Scale

The pH scale measures the acidity or basicity of a solution based on the concentration of hydrogen ions.

• Formula: $pH = -\log[H^+]$

• pH < 7: Acidic

• pH = 7: Neutral

• pH > 7: Basic

• Each unit change in pH represents a tenfold change in [H+].

• Example: If $[H^+] = 1 \times 10^{-3}$, then pH = 3.

pH and pOH Relationship

• Formula: $pH + pOH = 14$

Buffers

Buffers are solutions that resist changes in pH when small amounts of acid or base are added.

• Made from a weak acid and its conjugate base.

Carbonate Buffer System

The carbonate buffer system is crucial for maintaining blood pH.

• Reaction: $H_2CO_3 \rightleftharpoons HCO_3^-$

Titration

Titration is a laboratory technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• Formula used: $M_1 V_1 = M_2 V_2$

Most Important Formulas to Memorize

• $M = \frac{\text{mol}}{L}$

• $M_1 V_1 = M_2 V_2$

• $pH = -\log[H^+]$

• $pH + pOH = 14$

Common Exam Tricks

• Giving volumes in mL instead of L

• Hiding dilution steps in word problems

• Asking for moles instead of molarity

• Mixing acids/bases with molarity calculations

Quick Memory Cheat Sheet

• Acid = gives H+

• Base = takes H+

• Strong = fully dissociates

• M = mol / L

• Dilution = concentration drops

Mini Practice Problems

• Balance: $Fe + O_2 \rightarrow Fe_2O_3$ Balanced: $4Fe + 3O_2 \rightarrow 2Fe_2O_3$

• 0.5 mol in 0.25 L: M = 2.0 M

• Dilution: 3.0 M, 2.0 L → 6.0 L final; Find $M_2$ $M_2 = 1.0$ M

• Is HNO3 strong or weak? Strong

• pH if $[H^+] = 1 \times 10^{-5}$: pH = 5

Final Night Study Strategy

• Rewrite and memorize key formulas

• Practice molarity and dilution problems

• Memorize strong acids and bases

• Review buffer systems, especially carbonate buffer