General Chemistry Exam Review: Chemical Reactions, Solutions, Acids & Bases, and Equilibrium

Chemical Reactions and Thermochemistry

Role of Buffers in Biological Systems

Buffers are essential in biological systems to maintain stable pH levels, which is crucial for proper cellular function and enzyme activity.

• Definition: A buffer is a solution that resists changes in pH when small amounts of acid or base are added.

• Importance: Buffers help maintain homeostasis in living organisms by stabilizing pH levels.

• Example: The bicarbonate buffer system in blood helps regulate pH.

Colligative Properties

Colligative properties depend on the number of solute particles in a solution, not their identity.

• Key Properties: Boiling point elevation, freezing point depression, vapor pressure lowering, and osmotic pressure.

• Example: Adding salt to water increases its boiling point.

Endothermic and Exothermic Reactions

Chemical reactions can absorb or release energy, affecting temperature and reaction progress.

• Endothermic: Absorbs heat; temperature of surroundings decreases.

• Exothermic: Releases heat; temperature of surroundings increases.

• Equation: $\Delta H > 0$ for endothermic, $\Delta H < 0$ for exothermic reactions.

Solutions and Solubility

Concentration and Solubility

Concentration measures the amount of solute dissolved in a given quantity of solvent.

• Definition: Molarity (M) is moles of solute per liter of solution.

• Equation: $M = \frac{n}{V}$, where $n$ is moles of solute and $V$ is volume in liters.

• Example: Dissolving 25.6 g of KCl (molar mass = 74.6 g/mol) in 500 mL water: $n = \frac{25.6}{74.6} = 0.344$ mol, $M = \frac{0.344}{0.500} = 0.688$ M.

Solubility and Temperature

Solubility of solids in liquids generally increases with temperature, while solubility of gases decreases.

• Key Point: Higher temperature usually increases solubility of solids, but decreases solubility of gases.

• Example: More sugar dissolves in hot water than cold water.

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water.

• Strong Electrolytes: Completely dissociate in solution (e.g., NaCl, HCl).

• Weak Electrolytes: Partially dissociate (e.g., acetic acid).

• Nonelectrolytes: Do not dissociate (e.g., sugar).

Acids, Bases, and pH

Arrhenius Acids and Bases

The Arrhenius definition classifies acids and bases based on their behavior in water.

• Arrhenius Acid: Produces H+ ions in water (e.g., HCl, HNO3).

• Arrhenius Base: Produces OH- ions in water (e.g., NaOH, KOH).

Conjugate Acid-Base Pairs

Acids and bases exist in pairs that differ by one proton (H+).

• Example: In the reaction $\mathrm{HCO_3^-} + \mathrm{H_2O} \rightarrow \mathrm{H_2CO_3} + \mathrm{OH^-}$, the conjugate acid-base pairs are HCO3-/H2CO3 and H2O/OH-.

pH and Acidity

pH measures the acidity or basicity of a solution.

• Definition: $\mathrm{pH} = -\log[\mathrm{H^+}]$

• Low pH: More acidic; High pH: More basic.

• Example Table:

Substance	pH Value
Tomato Juice	4.2
Baking Soda	8.3
Lemonade	2.5
Vinegar	3.0

Most acidic: Lemonade (lowest pH).

Chemical Equilibrium

Dynamic Equilibrium

At equilibrium, the rates of the forward and reverse reactions are equal, and concentrations of reactants and products remain constant.

• Key Point: Equilibrium does not mean reactants and products are equal, but that their concentrations do not change over time.

• Example: In a closed system, the amount of reactants and products remains steady once equilibrium is reached.

Le Châtelier's Principle

Le Châtelier's Principle predicts how a system at equilibrium responds to changes in concentration, temperature, or pressure.

• Key Point: If a change is imposed on a system at equilibrium, the system shifts to counteract the change.

• Example: Adding H2O to $\mathrm{CrO_4^{2-} + 2H^+ \rightleftharpoons Cr_2O_7^{2-} + H_2O}$ shifts equilibrium to the left.

Lab Techniques and Calculations

Preparing Solutions

Calculating molarity and preparing solutions are essential laboratory skills.

• Equation: $M_1V_1 = M_2V_2$ (dilution equation)

• Example: Mixing 0.800 M HCl with water to make 1.5 L of solution: $M_1V_1 = M_2V_2$

Distinguishing Solutions and Pure Substances

It is important to differentiate between mixtures (solutions) and pure substances in chemistry.

• Solution: Homogeneous mixture of two or more substances.

• Pure Substance: Contains only one type of particle (element or compound).

• Example: Lemonade is a solution; distilled water is a pure substance.

Summary Table: Electrolyte Strength

Type	Examples	Electrolyte Strength
Strong Acid/Base/Salt	HCl, NaOH, NaCl	Strong Electrolyte
Weak Acid/Base/Salt	CH3COOH, NH3	Weak Electrolyte
Nonelectrolyte	Sugar, Ethanol	Nonelectrolyte

Additional info:

• Some explanations and context have been expanded for clarity and completeness.

• Tables have been recreated and summarized for study purposes.