Chem 101 Exam 3 Study Guide: Chemical Reactions, Energy, Gases, and Phase Changes

Chemical Reactions

Types of Chemical Reactions

Chemical reactions can be classified into several types based on the changes occurring to the reactants and products. Recognizing these types is essential for predicting products and understanding reaction mechanisms.

• Combination (Synthesis) Reaction: Two or more substances combine to form a single product. Example: $\mathrm{A} + \mathrm{B} \rightarrow \mathrm{AB}$

• Decomposition Reaction: A single compound breaks down into two or more simpler substances. Example: $\mathrm{AB} \rightarrow \mathrm{A} + \mathrm{B}$

• Replacement (Single/Double Displacement) Reaction: Atoms or ions are exchanged between compounds. Example (Single): $\mathrm{A} + \mathrm{BC} \rightarrow \mathrm{AC} + \mathrm{B}$ Example (Double): $\mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}$

• Combustion Reaction: A substance reacts with oxygen, releasing energy (often as heat and light). Example: $\mathrm{C}_x\mathrm{H}_y + \mathrm{O}_2 \rightarrow \mathrm{CO}_2 + \mathrm{H}_2\mathrm{O}$

Redox Reactions

Oxidation and Reduction

Redox reactions involve the transfer of electrons between substances. Understanding oxidation and reduction is key to identifying these reactions.

• Oxidation: Loss of electrons by a substance (LEO: Lose Electrons = Oxidation; OIL: Oxidation Is Loss).

• Reduction: Gain of electrons by a substance (GER: Gain Electrons = Reduction; RIG: Reduction Is Gain).

• Redox Reaction: A reaction where one species is oxidized and another is reduced.

• Oxidation States: Numbers assigned to atoms to track electron transfer. Changes in oxidation state indicate redox processes.

• Half-Reactions: Equations showing either oxidation or reduction alone. Used to balance redox equations.

Example: In the reaction $\mathrm{Zn} + \mathrm{Cu}^{2+} \rightarrow \mathrm{Zn}^{2+} + \mathrm{Cu}$, Zn is oxidized (loses electrons), Cu2+ is reduced (gains electrons).

The Mole and Chemical Calculations

The Mole Concept

The mole is a fundamental unit in chemistry for counting particles. It allows conversion between mass, number of particles, and volume.

• Mole: 1 mole = $6.022 \times 10^{23}$ particles (Avogadro's number).

• Moles to Molecules: Multiply moles by Avogadro's number.

• Molar Mass: Mass of 1 mole of a substance (g/mol).

• Mass-Mole Conversions: Use molar mass to convert between mass and moles.

Example: To find molecules in 2 moles of water: $2 \text{ mol} \times 6.022 \times 10^{23} = 1.204 \times 10^{24}$ molecules.

Calculations in Chemical Reactions

Stoichiometry uses balanced equations to relate quantities of reactants and products.

• Molar Ratios: Derived from coefficients in balanced equations.

• Theoretical Yield: Maximum possible product from given reactants.

• Reactant Calculations: Find mass of reactant needed for complete reaction or desired product.

Example: If 2 mol of A reacts with 1 mol of B to produce 2 mol of C, the molar ratio is 2:1:2.

Limiting Reactant and Percent Yield

The limiting reactant determines the maximum amount of product formed. Percent yield compares actual and theoretical yields.

• Limiting Reactant: The reactant used up first, limiting product formation.

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

Energy in Chemical Reactions

Types of Energy

Energy is the capacity to do work or produce heat. The law of conservation of energy states that energy cannot be created or destroyed.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

• Law of Conservation of Energy: Total energy remains constant in a closed system.

Energy Changes, Bond Energy, and Enthalpy

Chemical reactions involve breaking and forming bonds, which require or release energy.

• Bond Energy: Energy required to break a chemical bond.

• Enthalpy ($\Delta H$): Net heat change in a reaction. $\Delta H$ is negative for exothermic (energy released), positive for endothermic (energy absorbed).

• Exothermic Reaction: Releases energy; $\Delta H < 0$.

• Endothermic Reaction: Absorbs energy; $\Delta H > 0$.

• Cold Pack: Endothermic reaction (absorbs heat).

• Hot Pack: Exothermic reaction (releases heat).

• ATP: Adenosine triphosphate; hydrolysis releases energy for cellular processes.

Example: ATP hydrolysis: $\mathrm{ATP} + \mathrm{H}_2\mathrm{O} \rightarrow \mathrm{ADP} + \mathrm{P}_i + \text{energy}$

Energy Diagrams and Activation Energy

Energy diagrams illustrate the energy changes during a reaction.

• Reactant Energy: Initial energy level.

• Product Energy: Final energy level.

• Activation Energy ($E_a$): Minimum energy required to start a reaction.

• Exothermic Diagram: Products lower in energy than reactants.

• Endothermic Diagram: Products higher in energy than reactants.

Reaction Rates

The rate of reaction measures how quickly reactants are converted to products.

• Fast Rate: Rapid conversion.

• Slow Rate: Gradual conversion.

• Factors Affecting Rate: Concentration (higher = faster), temperature (higher = faster), catalyst (increases rate).

Chemical Equilibrium

Equilibrium Reactions

Some reactions are reversible and reach equilibrium, where forward and reverse rates are equal.

• Reversible Reaction: Can proceed in both directions.

• Equilibrium Symbol: $\rightleftharpoons$

• Equilibrium: Concentrations remain constant; rates are equal.

• M: Molarity, mol/L.

• Equilibrium Constant ($K$): Indicates extent of reaction; large $K$ favors products, small $K$ favors reactants.

• Le Chatelier's Principle: System shifts to counteract changes in concentration or temperature.

Calculating Equilibrium Constants

Equilibrium constant ($K$) is calculated using concentrations of reactants and products.

• Formula: $K = \frac{[\text{products}]}{[\text{reactants}]}$ (raised to their coefficients)

• Solving for $K$: Use known concentrations.

• Solving for Concentrations: Rearrange formula if $K$ and some concentrations are known.

Properties and Behavior of Gases

Kinetic Molecular Theory

The kinetic molecular theory explains gas behavior based on particle motion.

• Assumptions: Gas particles are in constant, random motion; collisions are elastic; volume of particles is negligible.

Pressure and Units

Pressure is the force exerted by gas particles per unit area.

• Pressure Formula: $\text{Pressure} = \frac{\text{Force}}{\text{Area}}$

• Units: atm, mmHg, torr, Pa.

• Barometer: Measures atmospheric pressure; high pressure = fair weather, low pressure = stormy weather.

• Elevation: Atmospheric pressure decreases with elevation due to fewer air molecules.

Greenhouse Gases and Effect

• Major Greenhouse Gases: CO2, CH4, N2O.

• Greenhouse Effect: Trapping of heat in Earth's atmosphere by greenhouse gases.

Gas Laws

Gas laws relate pressure, volume, temperature, and amount of gas.

• Boyle's Law: $P \propto \frac{1}{V}$ (Pressure inversely proportional to volume; $PV = \text{constant}$)

• Charles' Law: $V \propto T$ (Volume directly proportional to temperature; $\frac{V}{T} = \text{constant}$)

• Gay-Lussac's Law: $P \propto T$ (Pressure directly proportional to temperature; $\frac{P}{T} = \text{constant}$)

• Avogadro's Law: $V \propto n$ (Volume directly proportional to moles; $\frac{V}{n} = \text{constant}$)

Example: Boyle's Law explains why a syringe's volume decreases as pressure increases.

STP and Molar Volume

• STP: Standard Temperature and Pressure (0°C, 1 atm).

• Molar Volume: At STP, 1 mole of gas occupies 22.4 L.

• Calculations: Use molar volume to find volume or moles at STP.

Ideal Gas Law

• Formula: $PV = nRT$

• Variables: P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K).

• Use: Solve for any variable given the others.

Dalton's Law of Partial Pressures

• Law: Total pressure is the sum of partial pressures of each gas.

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

• Calculations: Find partial or total pressure given other values.

Phase Changes and Intermolecular Forces

Types of Phase Changes

Phase changes are transitions between solid, liquid, and gas states.

• Melting: Solid to liquid (endothermic).

• Freezing: Liquid to solid (exothermic).

• Vaporization: Liquid to gas (endothermic).

• Condensation: Gas to liquid (exothermic).

• Sublimation: Solid to gas (endothermic).

• Deposition: Gas to solid (exothermic).

Intermolecular Forces and Phase Changes

• Intermolecular Forces: Stronger forces = higher boiling/melting points.

• Types: Hydrogen bonding, dipole-dipole, London dispersion.

Example: Water (H2O) has high boiling point due to hydrogen bonding.

Heating Curves

Heating curves show temperature changes during phase transitions.

• Interpretation: Plateaus indicate phase changes; slopes indicate temperature increase.

Phase Change	Direction	Energy
Melting	Solid → Liquid	Endothermic
Freezing	Liquid → Solid	Exothermic
Vaporization	Liquid → Gas	Endothermic
Condensation	Gas → Liquid	Exothermic
Sublimation	Solid → Gas	Endothermic
Deposition	Gas → Solid	Exothermic