General Chemistry Study Guide: Chemical Kinetics, Equilibrium, Thermodynamics, Acids/Bases, and Electrochemistry

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Chemical Kinetics

Reaction Rates and Monitoring Progress

Chemical kinetics studies the speed at which chemical reactions occur and the factors affecting them. The rate of a reaction is defined as the change in concentration of reactants or products per unit time.

Reaction Rate: for disappearance of reactant A; for appearance of product B.
Monitoring Progress: Rates can be measured by tracking concentration changes over time.
Example: For the reaction , A is consumed at twice the rate of C formation.

Runner's speed analogy for reaction rate

Rate Laws and Experimental Determination

Rate laws express the relationship between reaction rate and reactant concentrations.

General Form: For , .
Experimental Determination: Rate laws must be determined experimentally, as they may not directly match reaction stoichiometry for complex mechanisms.
Example: Iodine clock reaction: .

Iodine clock reaction color change

Reaction Order and Integrated Rate Laws

The order of a reaction describes how the rate depends on reactant concentrations.

Zero Order: ;
First Order: ;
Second Order: ;
Half-life: For first order,

Rate laws for zero, first, and second order reactions Integrated rate laws: concentration vs time

Enzyme Kinetics and Michaelis-Menten Model

Enzyme-catalyzed reactions often display saturation behavior, leading to zero-order kinetics at high substrate concentrations.

Michaelis-Menten Equation:
Zero Order: At high substrate, rate is independent of [S].
Example: Ethanol metabolism in the liver follows zero-order kinetics at high concentrations.

Michaelis-Menten kinetics: rate vs substrate concentration Ethanol metabolism pathway in the liver

Chemical Equilibrium

Dynamic Equilibrium and the Equilibrium Constant

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products.

Equilibrium Constant (K): For ,
Law of Mass Action: The equilibrium expression depends on the balanced equation.
ICE Tables: Used to calculate equilibrium concentrations.

Le Châtelier’s Principle

Le Châtelier’s Principle states that a system at equilibrium will shift to counteract any imposed change.

Adding Reactant: Shifts equilibrium toward products.
Adding Product: Shifts equilibrium toward reactants.
Changing Temperature: Endothermic reactions: increasing temperature shifts equilibrium toward products; exothermic: toward reactants.

Thermodynamics

Spontaneity, Enthalpy, and Entropy

Thermodynamics determines whether a reaction is spontaneous and how much energy is exchanged.

Enthalpy (ΔH): Heat exchanged at constant pressure.
Entropy (ΔS): Measure of disorder;
Spontaneous Process: Occurs naturally;

Reaction coordinate diagram: ΔG and Ea

Gibbs Free Energy

Gibbs free energy combines enthalpy and entropy to predict spontaneity.

Equation:
Standard State:
Relationship to K:

Acids, Bases, and Buffers

Acid/Base Definitions and pH

Acids and bases are defined by their ability to donate or accept protons or electron pairs.

Arrhenius: Acids produce H+, bases produce OH- in water.
Brønsted-Lowry: Acids donate protons, bases accept protons.
Lewis: Acids accept electron pairs, bases donate electron pairs.
pH Calculation:

Strong and Weak Acids/Bases

Strong Acids: Completely dissociate (e.g., HCl, HNO3)
Weak Acids: Partially dissociate (e.g., CH3COOH)
Ka and Kb: Acid and base dissociation constants; relate to strength.

Buffer Solutions and Henderson-Hasselbalch Equation

Buffers resist changes in pH by neutralizing added acid or base.

Buffer Criteria: Mixture of weak acid and its conjugate base.
Henderson-Hasselbalch:

Electrochemistry

Redox Reactions and Oxidation States

Redox reactions involve electron transfer.

Oxidation: Loss of electrons; increase in oxidation state.
Reduction: Gain of electrons; decrease in oxidation state.
Assigning Oxidation States: Based on electronegativity and bonding.

Electrochemical Cells

Electrochemical cells convert chemical energy to electrical energy (voltaic/galvanic) or vice versa (electrolytic).

Anode: Site of oxidation.
Cathode: Site of reduction.
Cell Potential (Ecell):
Standard Electrode Potentials: Measured against the Standard Hydrogen Electrode (SHE).

Relationship Between Gibbs Energy and Cell Potential

Equation:
Nernst Equation:

Applications: Batteries, Fuel Cells, and Electrolysis

Batteries: Engineered voltaic cells (e.g., lead-acid, Li-ion).
Fuel Cells: Convert fuel (e.g., H2, CH4) to electricity.
Electrolysis: Uses electrical energy to drive non-spontaneous reactions.

Radioactivity and Nuclear Chemistry

Types of Radioactive Decay

Alpha Decay: Emission of He nucleus.
Beta Decay: Emission of electron or positron.
Gamma Decay: Emission of high-energy photon.

Kinetics of Radioactive Decay

Radioactive decay follows first-order kinetics.

Rate Law:
Half-life:
Example: Carbon-14 dating, medical tracers.

First order decay: concentration vs time and half-lives

Tables

Summary Table: Rate Laws for Zero, First, and Second Order Reactions

Reaction Order	Rate Law	Integrated Form
Zero	Rate = k	(A) – (A0) = -kt
First	Rate = k(A)	ln[(A)/(A0)] = -kt
Second	Rate = k(A)2	(1/A) – (1/A0) = kt

Summary Table: Common Strong Acids and Bases

Strong Acids	Strong Bases
HI, HBr, HCl, HClO4, H2SO4, HNO3	LiOH, NaOH, KOH, Ba(OH)2, Ca(OH)2, Sr(OH)2

Summary Table: Standard Electrode Potentials (Selected)

Half-Reaction	Eo (V)
2 H+ + 2 e- → H2	0.00
F2 + 2 e- → 2 F-	+2.87
Cu2+ + 2 e- → Cu	+0.34
Zn2+ + 2 e- → Zn	-0.76

Additional info:

Some context and examples were inferred for completeness, such as the use of ICE tables and the application of the Nernst equation.
Images included are directly relevant to the explanation of the adjacent paragraphs, such as reaction coordinate diagrams, enzyme kinetics, and radioactive decay.