Chemistry Energy Concepts

Thermodynamics is the study of energy and its transformations.
Chemical reactions involve energy transfer, either requiring or releasing energy.
Chemical reactions occur because of stability effects, where lower energy correlates to greater stability and lower reactivity.
Higher energy equates to lower stability and greater reactivity.

Energy is the capacity to do work or transfer heat.
Kinetic energy is the energy of motion, calculated as 1/2 * mass * velocity^2.
Potential energy is stored energy, often associated with position or chemical bonds.
The SI unit of energy is the joule (J).

The First Law of Thermodynamics states that energy cannot be created or destroyed, only transformed from one form to another.
This law helps understand energy conversions, such as mechanical to electrical or light to chemical.
Thermal energy is the kinetic energy of molecular motion, measured by temperature.
Heat is the transfer of thermal energy between objects due to temperature differences.
Chemical energy is potential energy stored in chemical bonds.

A thermodynamic system is a specific region of the universe under study.
The surroundings encompass everything outside the system.
A boundary separates the system from its surroundings.
System types include open (exchanging matter and energy), closed (exchanging only energy), and isolated (exchanging neither matter nor energy).

Internal energy (U) is the sum of kinetic and potential energies of all particles within a system.
Change in internal energy (ΔU) is the difference between final and initial internal energy.
Change in internal energy is equal to heat transfer plus work done on or by the system. (ΔU = q + w).
P-V work is work associated with volume changes in a system, under a constant pressure.
(Work = force x distance; w = P x ΔV, for an enclosed system)

Loss of energy from the system: ΔU is negative.
Gain of energy into the system: ΔU is positive.
State functions depend only on the current state; not how it was reached.
Examples of state functions : Internal energy (U) , enthalpy (H), volume and pressure.

If a state function changes, the change for the reverse process is equal but opposite.
The overall change for a complete cycle of a state function is always zero.
Physical work = force x displacement.
In chemical reactions, P-V work often occurs as a result of volume changes.

External pressure acts on the system's boundary.
Work is done as a result of changes in volume.
Work done by a system means the work is done by the system against its external surroundings. A positive volume change corresponds to a negative work done by the system.

Heat transfer (q) is positive when the system gains heat, negative when the system loses heat.
At constant volume: q = ΔU.
At constant pressure: ΔH = ΔU + P ΔV. (Change in Enthalpy).
Work done on the system has a positive value.
Work done by the system has a negative value.

Enthalpy (H) is a thermodynamic property that measures the total heat content of a system at constant pressure.
For reactions, the enthalpy change (ΔH) is important.
A positive ΔH indicates an endothermic process (heat absorbed).
A negative ΔH indicates an exothermic process (heat released).
Enthalpy is a state function and thus its value depends only on the initial and final states of the system, not the path.

Isothermal processes occur at constant temperature. This is usually due to contact with a heat reservoir.
For isothermal processes: ΔT = 0, ΔU = 0 and q = P ΔV
Adiabatic processes occur without heat exchange with the surroundings.
For adiabatic processes: q = 0; ΔU = -P ΔV = W

Standard enthalpy (ΔH°) values are defined under standard state conditions, typically 298K, and 1 bar pressure.
In chemical reactions, standard enthalpies refer to amounts under these specified conditions This allows for consistent comparison across different reactions.

Heat absorbed or released during a reaction is dependent on the quantities of reactants.
Standard enthalpy change of a reaction (ΔH°) is usually reported in kJ/mol.

Enthalpy of reaction is the heat change associated with the transformation that takes place within the reaction.
Different enthalpy changes can be defined, for instance enthalpy of formation (Δ_fH°), enthalpy change of combustion (Δ_combH°), enthalpy change of neutralization (Δ_neutH°), enthalpy change of fusion (Δ_fusH°)

Hess's Law: The overall enthalpy change for a reaction remains the same regardless of the pathway taken.
Enthalpy is a state function. All intermediary steps will have the same value.
This principle allows calculating enthalpy changes for complex reactions through known intermediary steps for enthalpy changes of easier reactions.

Standard heats of formation of reactants and products are used to determine overall enthalpy changes of reactions.

Enthalpy change (AH) of a reaction is found by subtracting the sum of the enthalpy changes of formation of the reactants from the sum of the enthalpy changes of formation of the products. A_rxnH°= Σn_p Δ_fH°(products) - Σn_r Δ_fH°(reactants).

Bond dissociation energy is the energy required to break a particular bond in a compound.
Bond dissociation energy can measure enthalpy changes in reactions.
Bond dissociation energies are used to estimate reaction enthalpy changes when standard thermodynamic data is unavailable.

Determining enthalpy changes of a reaction from bond dissociation energies involve finding bond energies for bonds broken and formed during the reaction, by subtracting broken bonds from formed.

Opposite charges attract each other with a force following Coulomb's law.
Lattice energy is the sum of the attractive electrostatic interaction energies between ions in a crystal.
Lattice energy describes the force needed to break ionic solids.

Lattice energy values describe the energy required for breaking up ionic solids and converting ions into gases.

The Born-Haber cycle is a method for calculating lattice energy in ionic compounds. It follows the energy changes involved in the formation of the crystal of ions.