Energy in Chemical Reactions

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• Chemistry involves the study of matter and its properties as well as how matter changes.
• Energy is closely associated with these changes.
• Energy change refers to the amount of energy that is either released or absorbed during a chemical or physical process.
• In chemical reactions, energy changes are primarily associated with the breaking and forming of chemical bonds.

Forms of Energy

• Kinetic energy is the energy of motion. Examples include the movement of molecules, a car driving, or electrons moving through a wire.
• Potential energy is stored energy. Examples include chemical bonds, a ball held in the air, or water behind a dam.
• Thermal energy is the energy associated with the temperature of an object. It is a form of kinetic energy due to the motion of atoms and molecules.
• Chemical energy is a form of potential energy stored in the bonds of chemical compounds. It is released or absorbed during chemical reactions.
• Electrical energy is the energy associated with the flow of electric charge.
• Radiant energy is the energy of electromagnetic radiation, such as light, microwaves, and radio waves.
• Nuclear energy is the energy stored in the nucleus of an atom. It is released during nuclear reactions like fission and fusion.

Energy Transfer

• Heat (q) is the transfer of thermal energy between two bodies at different temperatures. Heat flows from the hotter object to the cooler one.
• Work (w) is the energy transfer when a force acts over a distance. In chemistry, it often refers to the work done by a gas expanding or contracting.
• Energy can be transferred between a system and its surroundings in the form of heat or work.
• The change in internal energy (ΔU) of a system is the sum of the heat added to the system and the work done on the system: ΔU = q + w.
• This is the first law of thermodynamics: energy cannot be created or destroyed, only converted from one form to another.

Exothermic and Endothermic Processes

• Exothermic processes release energy into the surroundings, usually in the form of heat.
• The temperature of the surroundings increases in exothermic reactions.
• The change in enthalpy (ΔH) for an exothermic reaction is negative (ΔH < 0) .
• Examples of exothermic processes include combustion, freezing, and condensation.
• Endothermic processes absorb energy from the surroundings.
• The temperature of the surroundings decreases in endothermic reactions.
• The change in enthalpy (ΔH) for an endothermic reaction is positive (ΔH > 0).
• Examples of endothermic processes include melting, boiling, and chemical reactions like the decomposition of some compounds.

Enthalpy

• Enthalpy (H) is a thermodynamic property of a system, defined as the sum of the internal energy (U) and the product of pressure (P) and volume (V): H = U + PV.
• Enthalpy change (ΔH) is a measure of the heat absorbed or released during a chemical reaction at constant pressure.
• ΔH is more commonly used than ΔU in chemistry because most reactions occur under constant atmospheric pressure.
• ΔH = H(products) - H(reactants)
• A negative ΔH indicates an exothermic reaction, while a positive ΔH indicates an endothermic reaction.

Calorimetry

• Calorimetry is the experimental measurement of heat absorbed or released during a chemical or physical process.
• A calorimeter is a device used to measure heat flow.
• The heat capacity (C) of a substance is the amount of heat required to raise its temperature by one degree Celsius (or one Kelvin).
• Specific heat capacity (c) is the amount of heat required to raise the temperature of one gram of a substance by one degree Celsius (or one Kelvin).
• The equation q = mcΔT relates heat (q), mass (m), specific heat capacity (c), and temperature change (ΔT).
• In calorimetry experiments, the heat released or absorbed by the reaction (qreaction) is equal in magnitude but opposite in sign to the heat absorbed or released by the calorimeter (qcalorimeter): qreaction = -qcalorimeter.
• Bomb calorimeters are used to measure the heat of combustion at constant volume.
• Coffee-cup calorimeters (simple calorimeters) are used to measure heat changes at constant pressure.

Hess's Law

• Hess's Law states that the enthalpy change for a reaction is independent of the path taken, meaning that if a reaction can occur through multiple steps, the sum of the enthalpy changes for each step will equal the enthalpy change for the overall reaction.
• This law is a consequence of enthalpy being a state function.
• Hess's Law allows the calculation of enthalpy changes for reactions that are difficult or impossible to measure directly.
• By manipulating known thermochemical equations (i.e., reversing them, multiplying them by coefficients) and summing them, the enthalpy change for the target reaction can be determined.

Standard Enthalpy of Formation

• The standard enthalpy of formation (ΔH°f) is the enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions (298 K and 1 atm).
• The standard state of an element is its most stable form under standard conditions (e.g., O2(g) for oxygen, C(s, graphite) for carbon).
• The standard enthalpy of formation of an element in its standard state is defined as zero.
• The standard enthalpy change (ΔH°rxn) for a reaction can be calculated using the standard enthalpies of formation of the reactants and products: ΔH°rxn = ΣnΔH°f(products) - ΣnΔH°f(reactants), where n represents the stoichiometric coefficients in the balanced chemical equation.

Bond Enthalpy

• Bond enthalpy is the energy required to break one mole of a particular bond in the gaseous phase.
• Breaking bonds is an endothermic process (positive bond enthalpy), while forming bonds is an exothermic process (negative of bond enthalpy).
• Average bond enthalpies can be used to estimate enthalpy changes for reactions, particularly in the gas phase.
• ΔH ≈ Σ(bond enthalpies of bonds broken) - Σ(bond enthalpies of bonds formed)
• Bond enthalpy calculations are approximate because they use average values and do not account for the effects of neighboring bonds in a molecule.

Entropy

• Entropy (S) is a measure of the disorder or randomness of a system.
• The greater the disorder, the higher the entropy.
• Entropy is a state function.
• The change in entropy (ΔS) is positive for processes that increase disorder and negative for processes that decrease disorder.
• Units of entropy are typically J/K.
• Entropy tends to increase in processes where: a gas is formed from a solid or liquid; the number of gas molecules increases; a solid or liquid dissolves.

Second Law of Thermodynamics

• The second law of thermodynamics states that the total entropy of an isolated system always increases or remains constant in irreversible processes.
• For a spontaneous process, the entropy of the universe (system + surroundings) must increase: ΔSuniverse = ΔSsystem + ΔSsurroundings > 0.

Gibbs Free Energy

• Gibbs free energy (G) combines enthalpy (H) and entropy (S) to predict the spontaneity of a process at constant temperature and pressure: G = H - TS, where T is the absolute temperature in Kelvin.
• The change in Gibbs free energy (ΔG) is given by: ΔG = ΔH - TΔS.
• A negative ΔG indicates a spontaneous process (also called exergonic).
• A positive ΔG indicates a non-spontaneous process (also called endergonic). Energy must be supplied to drive the reaction.
• ΔG = 0 indicates that the system is at equilibrium.
• The spontaneity of a reaction can depend on temperature, as the TΔS term in the Gibbs free energy equation can become more or less significant at different temperatures.

Standard Free Energy Change

• The standard free energy change (ΔG°) is the change in Gibbs free energy when a reaction is carried out under standard conditions (298 K and 1 atm, 1M concentrations for solutions).
• Similar to standard enthalpy changes, standard free energy changes can be calculated using standard free energies of formation: ΔG°rxn = ΣnΔG°f(products) - ΣnΔG°f(reactants).

Free Energy and Equilibrium

• The change in Gibbs free energy (ΔG) is related to the standard free energy change (ΔG°) and the reaction quotient (Q) by the equation: ΔG = ΔG° + RTlnQ, where R is the ideal gas constant and T is the absolute temperature.
• At equilibrium, ΔG = 0, and the reaction quotient Q is equal to the equilibrium constant K: ΔG° = -RTlnK.
• This equation relates the standard free energy change to the equilibrium constant, allowing predictions about the extent to which a reaction will proceed to completion under standard conditions.

Energy change refers to the amount of energy released or absorbed during a chemical or physical process. Kinetic energy is the energy of motion. Potential energy is stored energy. Thermal energy is the energy associated with the temperature of an object.