Thermodynamics II

🧊Thermodynamics II Unit 9 – Chemical Reactions and Combustion

Chemical reactions and combustion are fundamental processes in thermodynamics. They involve the transformation of substances, energy exchange, and the study of reaction rates. Understanding these concepts is crucial for engineers and scientists working with energy systems and chemical processes. This unit covers key concepts like stoichiometry, reaction types, and thermodynamic principles. It also explores combustion processes, reaction kinetics, and practical applications in engineering. Mastering these topics provides a solid foundation for analyzing and optimizing chemical and energy-related systems.

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 or substances that take part 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
  • Thermodynamics branch of physics that deals with heat, work, and energy in a system, especially in relation to chemical reactions
  • Enthalpy (H) measure of the total heat content of a system at constant pressure
  • Entropy (S) measure of the disorder or randomness in a system
  • Gibbs free energy (G) thermodynamic potential that measures the capacity of a system to do maximum or reversible work at a constant temperature and pressure
  • Activation energy minimum energy required to initiate a chemical reaction
  • Catalyst substance that increases the rate of a chemical reaction without being consumed in the process

Types of Chemical Reactions

  • Synthesis or combination reaction two or more reactants combine to form a single product (2H2 + O2 -> 2H2O)
  • Decomposition reaction 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 substance reacts with oxygen, releasing energy in the form of heat and light (CH4 + 2O2 -> CO2 + 2H2O)
  • Acid-base reaction proton (H+) is transferred from an acid to a base (HCl + NaOH -> NaCl + H2O)
  • Redox reaction involves the transfer of electrons between species, resulting in changes in oxidation states (2Na + Cl2 -> 2NaCl)

Stoichiometry and Reaction Balancing

  • Balanced chemical equation reactants and products are represented using chemical formulas, with coefficients to ensure conservation of mass
  • Mole unit of measurement in chemistry, equal to 6.022 × 10^23 particles (atoms, molecules, or ions)
  • Molar mass mass of one mole of a substance, expressed in grams per mole (g/mol)
  • Limiting reactant reactant that is completely consumed in a reaction and determines the amount of product formed
  • Excess reactant reactant that remains after the limiting reactant is completely consumed
  • Theoretical yield maximum amount of product that can be obtained from the given amounts of reactants
  • Actual yield amount of product actually obtained from a reaction
  • Percent yield (actual yield / theoretical yield) × 100%

Thermodynamics of Reactions

  • Exothermic reaction releases energy to the surroundings, resulting in a negative change in enthalpy (ΔH < 0)
  • Endothermic reaction absorbs energy from the surroundings, resulting in a positive change in enthalpy (ΔH > 0)
  • Hess's law enthalpy change of a reaction is independent of the pathway and can be calculated by summing the enthalpy changes of individual steps
  • Standard enthalpy of formation (ΔH°f) enthalpy change when one mole of a compound is formed from its constituent elements in their standard states at 1 atm and 25°C
  • Entropy change (ΔS) measure of the change in disorder or randomness of a system during a process
  • Gibbs free energy change (ΔG) determines the spontaneity of a reaction at constant temperature and pressure
    • ΔG < 0: reaction is spontaneous
    • ΔG > 0: reaction is non-spontaneous
    • ΔG = 0: system is at equilibrium
  • Relationship between ΔG, ΔH, and ΔS: ΔG = ΔH - TΔS, where T is the absolute temperature in Kelvin

Combustion Processes

  • Complete combustion reaction in which a fuel is burned in excess oxygen, producing carbon dioxide and water (CH4 + 2O2 -> CO2 + 2H2O)
  • Incomplete combustion reaction in which a fuel is burned in insufficient oxygen, producing carbon monoxide and other byproducts (2CH4 + 3O2 -> 2CO + 4H2O)
  • Hydrocarbon combustion burning of compounds composed of hydrogen and carbon (octane: C8H18 + 12.5O2 -> 8CO2 + 9H2O)
  • Adiabatic flame temperature maximum temperature reached during a combustion process, assuming no heat loss to the surroundings
  • Calorific value or heating value measure of the energy released during the combustion of a specific amount of fuel (usually expressed in kJ/kg or kJ/mol)
  • Flue gas mixture of gases produced by the combustion of fuels, primarily containing nitrogen, carbon dioxide, water vapor, and excess oxygen
  • Air-fuel ratio mass ratio of air to fuel in a combustion process, affecting the efficiency and emissions of the reaction
  • Stoichiometric air-fuel ratio ideal ratio at which all the fuel is burned completely with no excess air

Reaction Kinetics and Rates

  • Reaction rate speed at which a chemical reaction proceeds, usually expressed as the change in concentration of a reactant or product per unit time (M/s)
  • Rate law mathematical expression that relates the reaction rate to the concentrations of the reactants, raised to a power (rate = k[A]^m[B]^n)
  • Rate constant (k) proportionality constant in the rate law, dependent on temperature and the nature of the reactants
  • Reaction order exponents (m, n) in the rate law, indicating the dependence of the reaction rate on the concentrations of the reactants
  • First-order reaction reaction in which the rate is directly proportional to the concentration of one reactant (rate = k[A])
  • Second-order reaction reaction in which the rate is proportional to the product of the concentrations of two reactants or the square of the concentration of one reactant (rate = k[A][B] or rate = k[A]^2)
  • Half-life time required for the concentration of a reactant to decrease to half its initial value in a first-order reaction (t1/2 = ln(2) / k)
  • Arrhenius equation relates the rate constant to the activation energy and temperature: k = Ae^(-Ea/RT), where A is the pre-exponential factor, Ea is the activation energy, R is the gas constant, and T is the absolute temperature
  • Collision theory explains the rate of a reaction in terms of the frequency and energy of collisions between reactant molecules
  • Transition state or activated complex unstable intermediate formed during a chemical reaction, with the highest potential energy along the reaction coordinate

Applications in Engineering

  • Combustion engines internal combustion engines (gasoline, diesel) and external combustion engines (steam turbines) that convert chemical energy from fuel into mechanical work
  • Fuel cells electrochemical devices that convert the chemical energy of a fuel (hydrogen, methanol) directly into electrical energy through a redox reaction
  • Catalytic converters devices used in vehicles to reduce harmful emissions by catalyzing the conversion of pollutants (CO, NOx, unburned hydrocarbons) into less harmful substances (CO2, N2, H2O)
  • Chemical reactors vessels designed to contain and control chemical reactions on an industrial scale (batch reactors, continuous stirred-tank reactors, plug flow reactors)
  • Flue gas desulfurization process of removing sulfur dioxide (SO2) from the exhaust flue gases of fossil fuel power plants to reduce air pollution and acid rain
  • Selective catalytic reduction (SCR) process used to reduce nitrogen oxide (NOx) emissions from industrial sources by reacting the NOx with ammonia (NH3) over a catalyst to form nitrogen (N2) and water (H2O)
  • Biofuels fuels derived from biomass, such as ethanol from corn or sugarcane and biodiesel from vegetable oils or animal fats, that can be used as alternatives to fossil fuels
  • Hydrogen economy vision of using hydrogen as a low-carbon energy carrier for transportation, heating, and power generation, with the potential to reduce greenhouse gas emissions

Problem-Solving Strategies

  • Identify the given information and the unknown quantities in the problem
  • Write a balanced chemical equation for the reaction, if applicable
  • Convert given quantities to moles using molar masses or ideal gas law, if necessary
  • Use stoichiometric ratios from the balanced equation to relate the moles of reactants and products
  • Apply relevant thermodynamic or kinetic equations (e.g., Hess's law, Gibbs free energy, rate law, Arrhenius equation) to solve for the desired quantity
  • Substitute given values into the appropriate equations and solve for the unknown variable
  • Check the units and the reasonableness of the answer based on the problem context
  • Consider any limiting reactants, side reactions, or other factors that may affect the outcome of the problem
  • Break down complex problems into smaller, manageable steps and solve them systematically
  • Practice solving a variety of problems to develop a strong understanding of the underlying concepts and problem-solving techniques


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© 2024 Fiveable Inc. All rights reserved.
AP® and SAT® are trademarks registered by the College Board, which is not affiliated with, and does not endorse this website.