Chemistry Energy Concepts

Thermodynamics

• 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

• 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).

Energy Changes and the Laws of Thermodynamics

• 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.

Important Concepts and Terms

• 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 and P-V Work

• 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)

Internal Energy and P-V Work

• 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.

Internal Energy and P-V Work

• 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.

Internal Energy and P-V Work

• 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.

Internal Energy and P-V Work

• 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

• 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 and Adiabatic Change

• 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

Thermodynamic Standard State

• Standard state conditions are crucial for comparing different reactions.
• Standard state conditions are usually 25°C (298K) and 1 bar pressure.
• Substance concentration is typically 1M.

Thermodynamic Standard State

• 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.

Thermodynamic Standard State

• 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 Reactions

• 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

• 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.

Calculating Enthalpy Changes from Heats of Formation

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

Calculating Enthalpy Changes from Heats of Formation

• 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).

Calculating Enthalpy Changes from Bond Dissociation Energy

• 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.

Calculates Enthalpy from Bond Dissociation

• 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.

Coulombic Forces and Lattice Energy

• 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.

Determining Lattice Energy

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

Born Haber Cycle

• 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.