CH107 Exam 2 Study Guide: Gases, Solutions, Acids & Bases, Nuclear Chemistry

Chapter 8: Gases

Units of Measurement for Gases

Gases are measured using several units for pressure, volume, temperature, and amount:

• Pressure: Common units include atmospheres (atm), millimeters of mercury (mmHg), torr, and pascals (Pa).

• Volume: Usually measured in liters (L) or milliliters (mL).

• Temperature: Always measured in Kelvin (K) for gas law calculations.

• Amount: Measured in moles (mol).

Example: 1 atm = 760 mmHg = 101,325 Pa

Boyle’s Law

Boyle’s Law describes the relationship between pressure and volume at constant temperature and amount of gas:

• Formula: $P_1 V_1 = P_2 V_2$

• Key Point: As pressure increases, volume decreases (inverse relationship).

• Example: If a gas at 2.0 L and 1.0 atm is compressed to 0.5 L, the new pressure is $P_2 = \frac{P_1 V_1}{V_2} = \frac{1.0 \times 2.0}{0.5} = 4.0$ atm.

Charles’ Law

Charles’ Law relates volume and temperature at constant pressure and amount:

• Formula: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

• Key Point: Volume increases as temperature increases (direct relationship).

• Example: A balloon expands when heated.

Gay-Lussac’s Law

Gay-Lussac’s Law describes the relationship between pressure and temperature at constant volume and amount:

• Formula: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$

• Key Point: Pressure increases as temperature increases.

Combined Gas Law

The combined gas law incorporates Boyle’s, Charles’, and Gay-Lussac’s laws:

• Formula: $\frac{P_1 V_1}{T_1} = \frac{P_2 V_2}{T_2}$

• Key Point: Used when pressure, volume, and temperature all change.

Avogadro’s Law

Avogadro’s Law relates the volume of a gas to the number of moles at constant pressure and temperature:

• Formula: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$

• Key Point: Volume is directly proportional to the number of moles.

Dalton’s Law of Partial Pressures

Dalton’s Law states that the total pressure of a mixture of gases is the sum of the partial pressures of each gas:

• Formula: $P_{total} = P_1 + P_2 + P_3 + \ldots$

• Example: Used in calculating oxygen and nitrogen pressures in air.

Chapter 9: Solutions

Definitions: Solute, Solvent, Solution

• Solute: The substance dissolved in a solution.

• Solvent: The substance that dissolves the solute (usually present in greater amount).

• Solution: A homogeneous mixture of solute and solvent.

Example: Salt (solute) dissolved in water (solvent) forms a saline solution.

Identifying Solute and Solvent

• Solute is usually the component present in lesser quantity.

• Solvent is the component present in greater quantity.

"Like Dissolves Like" Principle

This principle states that polar solutes dissolve in polar solvents, and nonpolar solutes dissolve in nonpolar solvents.

• Example: Sugar (polar) dissolves in water (polar); oil (nonpolar) does not.

Electrolytes and Non-Electrolytes

• Electrolyte: A substance that conducts electricity when dissolved in water (forms ions).

• Non-Electrolyte: Does not conduct electricity (does not form ions).

Example: NaCl is an electrolyte; glucose is a non-electrolyte.

Electrolyte Concentrations

• Measured in milliequivalents (mEq).

• Important in clinical settings (e.g., IV solutions).

Solution Concentration Calculations

• Mass/Volume Percent: $\text{\% (m/v)} = \frac{\text{grams of solute}}{\text{mL of solution}} \times 100$

• Mass/Mass Percent: $\text{\% (m/m)} = \frac{\text{grams of solute}}{\text{grams of solution}} \times 100$

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

Isotonic, Hypertonic, Hypotonic Solutions

These terms describe the relative concentration of solutes in solutions compared to cells:

• Isotonic: Same solute concentration as cells; cells remain unchanged.

• Hypertonic: Higher solute concentration than cells; cells shrink (water leaves).

• Hypotonic: Lower solute concentration than cells; cells swell (water enters).

Clinical Application: Incorrect IV solutions can cause cell damage.

Dilution Process and Calculations

• Formula: $C_1 V_1 = C_2 V_2$

• Used to calculate new concentration after dilution.

Chapter 10: Acids, Bases, and Equilibrium

Acids, Bases, and Neutralization Reactions

• Acid: Substance that donates a proton (H+).

• Base: Substance that accepts a proton.

• Neutralization: Reaction between acid and base to form water and salt.

Bronsted-Lowry Acids and Bases

• Bronsted-Lowry Acid: Proton donor.

• Bronsted-Lowry Base: Proton acceptor.

Common Acids and Bases

• Sulfuric acid: $H_2SO_4$

• Nitric acid: $HNO_3$

• Hydrochloric acid: $HCl$

• Phosphoric acid: $H_3PO_4$

• Acetic acid: $CH_3COOH$

• Carbonic acid: $H_2CO_3$

• Ammonia: $NH_3$

• Hydroxide bases: $NaOH$, $KOH$

Conjugate Acid-Base Pairs

• When an acid donates a proton, it forms its conjugate base.

• When a base accepts a proton, it forms its conjugate acid.

Strong vs. Weak Acids

• Strong acids: Completely dissociate in water.

• Weak acids: Partially dissociate.

Acid-Base Equilibrium and Reversible Reactions

• Acid-base reactions can be reversible.

• Equilibrium is reached when forward and reverse reaction rates are equal.

Le Chatelier’s Principle

• When reaction conditions change (concentration, temperature, pressure), equilibrium shifts to counteract the change.

• Example: Adding more acid shifts equilibrium to consume excess acid.

Water Dissociation Expression

• Formula: $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

• Used to calculate hydronium and hydroxide ion concentrations.

pH Calculations

• Formula: $pH = -\log [H_3O^+]$

• Formula: $[H_3O^+] = 10^{-pH}$

• pH < 7: Acidic; pH = 7: Neutral; pH > 7: Basic

Acid Reactions with Metals, Carbonates, Bicarbonates, and Bases

• Acids react with metals to produce hydrogen gas.

• Acids react with carbonates and bicarbonates to produce carbon dioxide.

• Acids react with bases to produce water and salt (neutralization).

Buffers and Blood pH Regulation

• Buffers resist changes in pH.

• Buffer Equation: $pH = pK_a + \log \left( \frac{[A^-]}{[HA]} \right)$

• Used to explain pH changes in blood during hyperventilation (alkalosis) or hypoventilation (acidosis).

• Le Chatelier’s principle applies to buffer systems.

Clinical Application: Interpreting arterial blood gas (ABG) values.

Chapter 5: Nuclear Chemistry

Types of Radiation

• Alpha radiation: Helium nuclei ($^4_2He$), low penetration.

• Beta radiation: Electrons ($^0_{-1}e$), moderate penetration.

• Gamma radiation: High-energy photons, high penetration.

Radioisotopes in Medicine

• Used in imaging, cancer treatment, and diagnostic tests.

• Example: Technetium-99m in nuclear medicine scans.

Half-Life Calculations

• Half-life: Time required for half the radioactive atoms to decay.

• Formula: $N = N_0 \left( \frac{1}{2} \right)^{\frac{t}{t_{1/2}}}$

• Used to determine remaining radioactivity over time.

MRI Chemistry

• Magnetic Resonance Imaging uses radioisotopes and nuclear magnetic resonance.

• Clinical applications include imaging soft tissues.

• Additional info: MRI does not use ionizing radiation; relies on nuclear spin properties.