# Chemistry CH101 - Unit 4: Reaction Kinetics, Phase Rule, Electrochemistry & Corrosion --- ## Reaction Kinetics ### Order and Molecularity of Reactions * **Order of a Reaction:** * Definition: The sum of the exponents of the concentration terms in the experimentally determined rate law equation. * Determined **experimentally**. * Can be **integer, fractional, or zero**. * Examples: * Rate = k\[A]^a\[B]^b. Order = a + b. * **Molecularity of a Reaction:** * Definition: The number of **reactant molecules** involved in the **elementary step** of a reaction that must collide simultaneously for the reaction to occur. * Determined from the **reaction mechanism**. * Must be a **positive integer** (1, 2, 3). * Elementary steps only. * **Bimolecular:** 2 molecules collide (e.g., A + B → Products). * **Unimolecular:** 1 molecule reacts (e.g., A → Products). * **Termolecular:** 3 molecules collide (rare). ### First and Second Order Reactions * **First Order Reaction:** * Rate depends on the concentration of **one reactant** raised to the power of one. * Rate Law: Rate = k\[A] * Integrated Rate Law: **ln([A]_t / [A]₀) = -kt** or **k = (1/t) ln([A]₀ / [A]_t)** * Half-life (t_1/2): **t_1/2 = ln(2) / k** (Independent of initial concentration). * Units of k: **s^-1** or **time^-1**. * **Second Order Reaction:** * Rate depends on the concentration of **one reactant** raised to the power of two, OR the concentration of **two reactants** each raised to the power of one. * **Type 1:** Rate = k\[A]² * Integrated Rate Law: **(1/[A]_t) - (1/[A]₀) = kt** or **k = (1/t) * [(1/[A]_t) - (1/[A]₀)]** * Half-life (t_1/2): **t_1/2 = 1 / (k[A]₀)** (Dependent on initial concentration). * Units of k: **L mol^-1 s^-1** or **M^-1 time^-1**. * **Type 2:** Rate = k\[A]\[B] * Complex integrated rate law (often simplified based on initial concentrations). ### Energy of Activation (E_a) * Definition: The **minimum amount of energy** that reactant molecules must possess to undergo a chemical reaction. * **Arrhenius Equation:** Relates rate constant (k) to temperature (T) and activation energy (E_a). * **k = A e^-Ea/RT** * k = rate constant * A = pre-exponential factor (frequency factor) * E_a = activation energy (J/mol) * R = ideal gas constant (8.314 J/mol·K) * T = absolute temperature (K) * **Graphical Method:** Plotting ln(k) vs. 1/T gives a straight line with slope **-E_a/R**. **[Figure: Arrhenius Plot]** This plot graphically represents the relationship between the rate constant and temperature, allowing for the experimental determination of the activation energy. A steeper slope indicates a higher activation energy. * **Key Points:** * Higher E_a means a slower reaction rate at a given temperature. * Increasing temperature increases the rate constant (k) and thus the reaction rate. --- ## Phase Rule * **Gibbs Phase Rule:** **F = C - P + 2** * **F**: Degrees of freedom (minimum number of independent variables, e.g., temperature, pressure, concentration, that must be fixed to define the system completely). * **C**: Number of components (minimum number of independent chemical species needed to describe the system). * **P**: Number of phases (a homogeneous, physically distinct, and mechanically separable portion of a system). * **Important Concepts:** * A **phase** is uniform in chemical composition and physical state. * **Components** are independent chemically. * **Degrees of freedom** indicate the variability of the system. ### Application to One Component System (Water) * Here, **C = 1** (only H₂O). * Phase Rule becomes: **F = 1 - P + 2 = 3 - P** * **Diagram:** **[Figure: Water Phase Diagram]** This diagram illustrates the stable phases of water (solid, liquid, gas) under different temperature and pressure conditions and shows the curves and points where phases coexist. It's crucial for understanding phase transitions. * **Regions:** * **One Phase (e.g., liquid water):** P = 1. F = 3 - 1 = 2. (T and P can vary independently). * **Two Phases (e.g., liquid-gas coexistence, boiling curve):** P = 2. F = 3 - 2 = 1. (If T is fixed, P is determined, or vice-versa). * **Three Phases (Triple Point):** P = 3. F = 3 - 3 = 0. (T and P are fixed; unique point for water). * **Triple Point of Water:** Occurs at **0.01 °C (273.16 K)** and **611.657 Pa**. All three phases (ice, liquid water, water vapor) coexist in equilibrium. --- ## Electrochemistry ### Equilibrium Potential (E⁰) * Definition: The **potential difference** of an electrode when it is in equilibrium with its ions in solution, with all species at **unit activity**. * Also known as **Standard Electrode Potential**. * Measured relative to the **Standard Hydrogen Electrode (SHE)**, which is assigned E⁰ = 0.00 V. * **Nernst Equation:** Relates the electrode potential (E) at non-standard conditions to the standard electrode potential (E⁰) and concentrations/activities. * **E = E⁰ - (RT / nF) ln(Q)** * R = ideal gas constant * T = temperature (K) * n = number of electrons transferred * F = Faraday constant * Q = reaction quotient ### Electrochemical Cells * **Galvanic Cell (Voltaic Cell):** * Converts **chemical energy into electrical energy**. * Occurs via a **spontaneous redox reaction**. * Consists of two half-cells: **anode** (oxidation) and **cathode** (reduction). * **Anode:** Negative electrode (where electrons are released). * **Cathode:** Positive electrode (where electrons are consumed). * **Cell Potential (E_cell):** E_cell = E_cathode - E_anode (using standard reduction potentials). * **Concentration Cell:** * A type of galvanic cell where the **two half-cells have the same electrodes** but **different concentrations** of the electrolyte. * The difference in concentration drives the reaction. * The electrode with the **lower concentration** acts as the **anode** (oxidation occurs to increase ion concentration). * The electrode with the **higher concentration** acts as the **cathode** (reduction occurs to decrease ion concentration). * E_cell = (RT / nF) ln(C_cathode / C_anode) --- ## Corrosion ### Electrochemical Theory of Corrosion * Definition: Corrosion is the **gradual destruction of materials** (usually metals) by chemical or electrochemical reaction with their environment. * **Electrochemical Process:** Corrosion, especially of metals, is an **electrochemical process** involving: * **Anodic Area:** Where oxidation occurs (metal loses electrons and dissolves as ions). Ex: Fe → Fe²⁺ + 2e^- * **Cathodic Area:** Where reduction occurs, consuming the electrons from the anode. * In neutral/alkaline solutions: O₂ + 2H₂O + 4e^- → 4OH^- * In acidic solutions: 2H⁺ + 2e^- → H₂ * **Electrolyte:** A conductive medium (e.g., moisture, acidic rain) is necessary to complete the circuit. * **Formation of Corrosion Products:** Metal ions react with other species (like oxygen and water) to form oxides, hydroxides, or salts (rust for iron). ### Protection of Corrosion 1. **Protective Coatings:** * **Barrier Coatings:** Paint, plastic, enamel. Prevent contact with the corrosive environment. * **Metallic Coatings:** Plating with a less reactive metal (e.g., tin plating on steel). If scratched, the more reactive metal corrodes preferentially. 2. **Cathodic Protection:** * Making the metal to be protected the **cathode** in an electrochemical cell. * **Sacrificial Anodes:** Connecting a more reactive metal (e.g., Mg, Zn, Al) to the metal to be protected. The sacrificial anode corrodes preferentially. **[Figure: Sacrificial Anode Protection]** This diagram shows a more reactive metal (sacrificial anode) attached to the metal to be protected (e.g., ship hull). The sacrificial anode corrodes instead of the hull, extending its lifespan. * **Impressed Current Cathodic Protection:** Using an external DC power source to force electrons onto the metal structure, making it the cathode. 3. **Anodic Protection:** * For metals that form a stable, passive oxide layer (e.g., stainless steel, titanium). * Making the metal the **anode** under controlled potential to promote the formation of a protective passive film. 4. **Alloying:** * Forming alloys that are inherently more resistant to corrosion (e.g., stainless steel contains chromium).