Thermodynamics II Unit 9 Review: 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 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