7.1 Laws of thermodynamics

Thermodynamics is all about energy and how it moves around. The laws of thermodynamics are like the rulebook for energy in chemical reactions. They tell us what's possible and what's not when it comes to energy changes.

These laws are super important for understanding how reactions work. They help us figure out if a reaction will happen on its own, how much energy it'll release or absorb, and what the end result will be. It's like a roadmap for chemical processes.

Thermodynamics: First Law and Energy Conservation

Energy Conservation in Chemical Systems

• The first law of thermodynamics states that the total energy of an isolated system is constant, and energy can neither be created nor destroyed, only converted from one form to another (potential energy to kinetic energy)
• In a chemical reaction, the change in internal energy (ΔU) is equal to the heat (q) added to the system plus the work (w) done on the system: $ΔU = q + w$
• For a closed system at constant volume, the change in internal energy is equal to the heat added to the system: $ΔU = q_v$
• For a closed system at constant pressure, the change in enthalpy (ΔH) is equal to the heat added to the system: $ΔH = q_p$

Implications of the First Law

• The first law implies that energy is conserved in chemical reactions, and the total energy of the reactants must equal the total energy of the products (bond breaking and formation)
• Energy can be transferred between a system and its surroundings in the form of heat or work, but the total energy remains constant
• The first law provides a framework for understanding energy changes in chemical processes, such as exothermic and endothermic reactions
• The law also helps explain the relationship between internal energy, enthalpy, and the heat and work exchanged with the surroundings

Entropy and Spontaneity: The Second Law

Entropy and the Second Law of Thermodynamics

• The second law of thermodynamics states that the total entropy of an isolated system always increases over time, and the entropy of the universe tends to a maximum
• Entropy (S) is a measure of the disorder or randomness of a system, and the change in entropy (ΔS) is equal to the heat transferred (q) divided by the absolute temperature (T): $ΔS = \frac{q}{T}$
• For a spontaneous process, the change in entropy of the universe (system + surroundings) must be greater than zero: $ΔS_{universe} > 0$
• The second law explains why certain processes occur spontaneously in nature, such as the diffusion of gases or the mixing of liquids

Gibbs Free Energy and Spontaneity

• The Gibbs free energy (G) is a thermodynamic quantity that combines the effects of enthalpy and entropy to determine the spontaneity of a process at constant temperature and pressure: $ΔG = ΔH - TΔS$
• A process is spontaneous if ΔG is negative, non-spontaneous if ΔG is positive, and at equilibrium if ΔG is zero
• The sign of ΔG determines the direction of a chemical reaction: negative ΔG favors the forward reaction, while positive ΔG favors the reverse reaction
• The second law and Gibbs free energy provide a means to predict the spontaneity and equilibrium state of chemical reactions (synthesis of ammonia)

Thermal Equilibrium and Temperature: The Zeroth Law

Thermal Equilibrium and the Zeroth Law

• The zeroth law of thermodynamics states that if two systems are in thermal equilibrium with a third system, then they are in thermal equilibrium with each other
• Thermal equilibrium occurs when two systems have the same temperature and there is no net transfer of heat between them (a hot object and a cold object in contact)
• The zeroth law establishes the concept of temperature as a property that determines thermal equilibrium
• The law also provides a basis for comparing temperatures and constructing temperature scales (Celsius, Kelvin)

Temperature Measurement

• Temperature is a measure of the average kinetic energy of the particles in a system, and it determines the direction of heat flow between systems
• The zeroth law provides the basis for temperature measurement using thermometers, which are devices that establish thermal equilibrium with the system being measured (mercury thermometer, thermocouple)
• Different types of thermometers use various thermometric properties, such as the expansion of liquids or gases, or the change in electrical resistance with temperature
• Accurate temperature measurement is essential for studying thermodynamic processes and determining the direction of heat transfer

State Functions vs Path Functions

State Functions

• A state function is a property of a system that depends only on the current state of the system, not on the path taken to reach that state
• Examples of state functions include internal energy (U), enthalpy (H), entropy (S), and Gibbs free energy (G)
• The change in a state function (e.g., ΔU, ΔH, ΔS, ΔG) is independent of the path and can be calculated using only the initial and final states of the system
• State functions are useful for describing the thermodynamic properties of a system and predicting the outcome of processes (calculating enthalpy change using Hess's law)

Path Functions

• A path function is a property of a system that depends on the path taken to reach the final state, not just the initial and final states
• Examples of path functions include heat (q) and work (w)
• The change in a path function (e.g., q, w) depends on the specific process or path taken by the system and requires knowledge of the entire path to calculate
• Path functions are important for understanding the mechanisms of energy transfer between a system and its surroundings (calculating work done by gas expansion)
• The values of path functions can vary for different processes connecting the same initial and final states, while state functions remain constant