🔥Thermodynamics I Unit 15 – Chemical Reactions and Combustion

Key Concepts and Definitions

• Chemical reaction process in which one or more substances (reactants) are converted into one or more different substances (products)
• Reactants starting materials in a chemical reaction
• Products substances formed as a result of a chemical reaction
• Stoichiometry quantitative study of the relative amounts of reactants and products in chemical reactions
  • Balanced chemical equations represent the relative numbers of reactant and product molecules
  • Mole ratios used to determine the quantities of reactants needed or products formed
• Enthalpy (H) measure of the total heat content of a system
  • Change in enthalpy (ΔH) indicates the amount of heat absorbed or released during a reaction at constant pressure
• Entropy (S) measure of the disorder or randomness of a system
  • Second law of thermodynamics states that the total entropy of an isolated system always increases over time

Types of Chemical Reactions

• Synthesis reaction two or more reactants combine to form a single product (2H₂ + O₂ → 2H₂O)
• Decomposition reaction a single compound breaks down into two or more simpler substances (2H₂O → 2H₂ + O₂)
• Single displacement reaction one element replaces another element in a compound (Zn + 2HCl → ZnCl₂ + H₂)
• Double displacement reaction two compounds exchange ions to form two new compounds (NaCl + AgNO₃ → AgCl + NaNO₃)
• Combustion reaction a substance reacts with oxygen, releasing heat and often light (CH₄ + 2O₂ → CO₂ + 2H₂O)
• Acid-base reaction involves the transfer of protons (H⁺) between substances (HCl + NaOH → NaCl + H₂O)
• Redox reaction (oxidation-reduction) involves the transfer of electrons between species
  • Oxidation loss of electrons and increase in oxidation state
  • Reduction gain of electrons and decrease in oxidation state

Combustion Basics

• Combustion exothermic chemical reaction between a fuel and an oxidant, usually atmospheric oxygen, that produces heat and often light
• Hydrocarbon combustion involves the reaction of hydrocarbons (compounds containing only carbon and hydrogen) with oxygen to produce carbon dioxide and water (CH₄ + 2O₂ → CO₂ + 2H₂O)
• Complete combustion occurs when a fuel is burned in sufficient oxygen, resulting in the production of carbon dioxide and water
• Incomplete combustion occurs when there is insufficient oxygen, leading to the formation of carbon monoxide and other byproducts (2CH₄ + 3O₂ → 2CO + 4H₂O)
• Flammability limits range of fuel-to-air ratios within which combustion can occur
  • Lower flammability limit (LFL) minimum concentration of fuel in air required for combustion
  • Upper flammability limit (UFL) maximum concentration of fuel in air that will sustain combustion
• Adiabatic flame temperature theoretical maximum temperature that can be achieved during combustion, assuming no heat loss to the surroundings
• Combustion efficiency ratio of the actual heat released during combustion to the theoretical maximum heat that could be released

Thermochemistry and Energy Transfer

• Thermochemistry study of heat and energy associated with chemical reactions and physical transformations
• Heat (q) energy transferred between systems due to a temperature difference
• Specific heat capacity (c) amount of heat required to raise the temperature of one gram of a substance by one degree Celsius
• Enthalpy of reaction (ΔH) heat absorbed or released during a chemical reaction at constant pressure
  • Exothermic reactions release heat to the surroundings (negative ΔH)
  • Endothermic reactions absorb heat from the surroundings (positive ΔH)
• Hess's law states that the total enthalpy change for a reaction is independent of the route taken
  • Used to calculate enthalpy changes for reactions that cannot be directly measured
• Calorimetry measures the heat transferred during a chemical reaction or physical process
  • Bomb calorimeter used to measure the heat of combustion of a fuel
• Heat of formation (ΔH°f) enthalpy change when one mole of a compound is formed from its constituent elements in their standard states

Reaction Rates and Kinetics

• Reaction rate speed at which a chemical reaction proceeds, typically expressed as the change in concentration of a reactant or product per unit time
• Factors affecting reaction rates include temperature, concentration, pressure, surface area, and the presence of catalysts
  • Increasing temperature typically increases reaction rates by providing more kinetic energy for collisions
  • Higher concentrations of reactants lead to more frequent collisions and faster reaction rates
• Activation energy (Ea) minimum energy required for reactants to collide and form products
  • Catalysts lower the activation energy, increasing the reaction rate without being consumed
• Rate law mathematical expression relating the reaction rate to the concentrations of reactants
  • Rate constant (k) proportionality constant in the rate law, dependent on temperature
• Arrhenius equation describes the relationship between the rate constant and temperature: $k = Ae^{-Ea/RT}$
  • A pre-exponential factor related to the frequency of collisions
  • R gas constant (8.314 J/mol·K)
• Reaction mechanisms series of elementary steps that describe the detailed molecular pathway of a reaction
  • Rate-determining step slowest step in a reaction mechanism, which determines the overall reaction rate

Equilibrium in Chemical Reactions

• Chemical equilibrium state in which the forward and reverse reactions proceed at equal rates, resulting in no net change in the concentrations of reactants and products
• Law of mass action states that the rate of a reaction is proportional to the product of the concentrations of the reactants, each raised to a power equal to its stoichiometric coefficient
• Equilibrium constant (K) ratio of the product of the concentrations of the products to the product of the concentrations of the reactants, each raised to their stoichiometric coefficients
  • For the general reaction aA + bB ⇌ cC + dD, the equilibrium constant is expressed as: $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
• Le Chatelier's principle states that when a system at equilibrium is disturbed, the system will shift to counteract the disturbance and re-establish equilibrium
  • Changes in concentration, pressure, volume, or temperature can shift the equilibrium position
• Equilibrium shifts increasing the concentration of a reactant or decreasing the concentration of a product will shift the equilibrium to the right (towards products)
  • Decreasing reactant concentration or increasing product concentration shifts equilibrium to the left (towards reactants)
• Temperature changes affect the equilibrium constant and the position of equilibrium
  • Increasing temperature favors the endothermic direction, while decreasing temperature favors the exothermic direction

Applications in Engineering

• Combustion engines (internal combustion engines, gas turbines) rely on the controlled combustion of fuels for power generation
  • Efficiency of combustion engines depends on factors such as fuel type, air-fuel ratio, and operating conditions
• Catalytic converters used in automobiles to reduce harmful emissions by promoting the complete combustion of unburned hydrocarbons and carbon monoxide
• Chemical reactors vessels designed to contain and control chemical reactions on an industrial scale
  • Batch reactors single vessel where reactants are added, mixed, and allowed to react for a specific time
  • Continuous stirred-tank reactors (CSTRs) reactants are continuously added and mixed, while products are continuously removed
  • Plug flow reactors (PFRs) reactants flow through a tubular reactor in a plug-like manner, with no mixing in the axial direction
• Process optimization involves selecting the most efficient and economical reaction conditions, such as temperature, pressure, and reactant concentrations, to maximize product yield and minimize waste
• Fuel cells electrochemical devices that convert the chemical energy of a fuel (hydrogen, methanol) directly into electrical energy through redox reactions
  • Proton exchange membrane fuel cells (PEMFCs) commonly used in vehicle applications due to their low operating temperature and high power density

Problem-Solving Techniques

• Dimensional analysis method of converting between different units by multiplying or dividing by conversion factors
  • Useful for solving problems involving mass, volume, density, and concentration
• Stoichiometry calculations involve determining the quantities of reactants needed or products formed in a chemical reaction
  • Mole ratios from balanced chemical equations are used to convert between the amounts of different substances
• Limiting reactant determines the maximum amount of product that can be formed in a reaction
  • Reactant that is completely consumed first, limiting the extent of the reaction
• Theoretical yield maximum amount of product that can be obtained based on the limiting reactant and the balanced chemical equation
• Percent yield ratio of the actual yield to the theoretical yield, expressed as a percentage
  • $\text{Percent yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$
• Equilibrium calculations involve determining the concentrations of reactants and products at equilibrium, given the initial concentrations and the equilibrium constant
  • ICE tables (Initial, Change, Equilibrium) used to organize information and solve for equilibrium concentrations
• Thermochemical calculations determine the heat absorbed or released during a reaction or process
  • Specific heat capacity, heat of formation, and Hess's law are used to calculate enthalpy changes
• Kinetics calculations involve determining reaction rates, rate constants, and activation energies from experimental data
  • Graphical methods, such as plotting the natural logarithm of the rate constant vs. the inverse of temperature (Arrhenius plot), can be used to determine the activation energy