CHE 103 Exam #3 Study Guide: Acids, Bases, Solutions, Organic Chemistry, and Chemical Quantities

Acids and Bases

Bronsted-Lowry Acid-Base Theory

The Bronsted-Lowry theory defines acids as proton (H+) donors and bases as proton acceptors. Conjugate acid-base pairs differ by one proton.

• Acid: Substance that donates an H+ ion.

• Base: Substance that accepts an H+ ion.

• Conjugate Acid: Formed when a base gains a proton.

• Conjugate Base: Formed when an acid loses a proton.

• Example: In the reaction HSO4- + OH- → SO42- + H2O, HSO4- is the acid, OH- is the base, SO42- is the conjugate base, and H2O is the conjugate acid.

pH and Acid Strength

pH is a measure of the hydrogen ion concentration in a solution. Strong acids and bases dissociate completely in water, while weak acids and bases do not.

• pH Formula: $pH = -\log[H^+]$

• Strong Acids: HCl, HNO3, H2SO4

• Strong Bases: NaOH, KOH

• Example Table:

[H3O+]	[OH-]	pH	Acidic/Basic/Natural
5.6 × 10-7	1.8 × 10-8	6.25	Acidic
1.0 × 10-10	8.2 × 10-5	9.92	Basic

Le Chatelier's Principle

Shifts in Chemical Equilibrium

Le Chatelier's Principle states that if a system at equilibrium is disturbed, the system will shift to counteract the disturbance and restore equilibrium.

• Adding Reactant: Shifts equilibrium toward products.

• Removing Reactant: Shifts equilibrium toward reactants.

• Adding Product: Shifts equilibrium toward reactants.

• Removing Product: Shifts equilibrium toward products.

• Example: CH3COOH(aq) + NH3(aq) ⇌ CH3COO-(aq) + NH4+(aq)

Electrolytes and Solubility

Types of Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water. They are classified as strong, weak, or non-electrolytes.

• Strong Electrolytes: Dissociate completely (e.g., NaCl, HCl).

• Weak Electrolytes: Partially dissociate (e.g., CH3COOH).

• Non-electrolytes: Do not dissociate (e.g., sugar).

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water.

Positive Ion	Negative Ion	Solubility
Na+, K+, NH4+	Any	Soluble
Any	NO3-, CH3COO-	Soluble
Any	SO42-	Soluble except with Ba2+, Pb2+, Ca2+
Any	Cl-, Br-, I-	Soluble except with Ag+, Pb2+

Solutions and Concentration

Molarity and Dilution

Molarity (M) is the concentration of a solution expressed as moles of solute per liter of solution. Dilution involves adding solvent to decrease concentration.

• Molarity Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Dilution Formula: $C_1V_1 = C_2V_2$

• Example: To prepare 375 mL of 1.8% (w/v) solution from 6.2% (w/v) stock: $V_1 = \frac{C_2V_2}{C_1} = \frac{(1.8)(375)}{6.2} = 109 \text{ mL}$

Colligative Properties

Colligative properties depend on the number of solute particles in solution, not their identity. Examples include boiling point elevation and osmotic pressure.

• Isotonic Solution: Same osmotic pressure as cells; no net movement of water.

• Hypotonic Solution: Lower osmotic pressure; water enters cells, causing hemolysis.

• Hypertonic Solution: Higher osmotic pressure; water leaves cells, causing crenation.

• Example: 0.9% NaCl and 5.0% glucose are isotonic to human cells.

Gas Laws

Ideal Gas Law and Related Equations

Gas laws describe the behavior of gases in terms of pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$

• Standard Molar Volume: 1 mol gas at STP = 22.4 L

• Standard Pressure: 1 atm = 760 mmHg

• Partial Pressure: $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

• Example: For a mixture of gases, add the individual partial pressures to get the total pressure.

Chemical Reactions and Stoichiometry

Balancing Chemical Equations

Balancing equations ensures the same number of atoms of each element on both sides of the reaction.

• Example: C9H20 + 14 O2 → 9 CO2 + 10 H2O

Stoichiometric Calculations

Stoichiometry involves using balanced equations to calculate amounts of reactants and products.

• Mole-Mass Conversion: $\text{mass} = \text{moles} \times \text{molar mass}$

• Example: 345 mL of H2O at STP contains $\frac{345 \text{ mL}}{22,400 \text{ mL/mol}} = 0.0154 \text{ mol}$

Organic Chemistry: Nomenclature and Structures

IUPAC Naming of Alkanes

Alkanes are named according to the number of carbon atoms and the position of substituents.

• Parent Chain: Longest continuous chain of carbon atoms.

• Substituents: Groups attached to the parent chain; named and numbered for lowest possible locants.

• Example: 6-bromo-2,5-dimethylnonane

Condensed and Line-Angle Formulas

Organic molecules can be represented in different ways for clarity and simplicity.

• Condensed Formula: Shows all atoms in a linear format (e.g., CH3CH2CH2CH3).

• Line-Angle Formula: Each vertex represents a carbon atom; hydrogens are implied.

Isomers

Isomers are compounds with the same molecular formula but different structures.

• Structural Isomers: Same formula, different connectivity.

• Example: CH3CH2CH2CH3 vs. CH3CH(CH3)CH3

Periodic Table of the Elements

Element Classification

The periodic table organizes elements by atomic number and properties. Groups and periods indicate similar chemical behavior.

• Groups: Vertical columns; elements have similar valence electron configurations.

• Periods: Horizontal rows; elements have the same number of electron shells.

Additional Info

• All equations and constants are provided in standard LaTeX format for clarity.

• Solubility rules and electrolyte classification are essential for predicting solution behavior.

• Organic nomenclature and isomerism are foundational for understanding molecular diversity.