Thermodynamics II: Entropy and Enthalpy Quiz

Thermodynamics: Fundamentals

• Thermodynamics describes the macroscopic state of a complex system using a small number of macroscopic variables, such as pressure and temperature, also known as state variables. Thermodynamic potentials are also used.
• This subject encompasses a broad range of phenomena, including the efficiency of heat engines and heat pumps, along with chemical processes and biological processes of life.

System, Universe, and Surroundings

• The universe comprises the system, the surroundings, and the universe as a whole.
• The system is the subject of interest (e.g., a block of iron, a beaker of water, an engine, a human body).
• The surroundings are the rest of the universe outside the system.

Types of Systems

• Open Systems: Both matter and energy are exchanged between the system and its surroundings. (e.g., an open flask)
• Closed Systems: Energy can be exchanged between the system and surroundings, but matter cannot. (e.g., sealed bottle)
• Isolated Systems: Neither matter nor energy are exchanged between the system and surroundings. (e.g., a stoppered vacuum flask).

Extensive and Intensive Properties

• Extensive Properties: Depend on the amount of matter in the system, e.g., mass, volume.
• Intensive Properties: Independent of the amount of matter in the system; e.g., temperature, density.
• The density of a substance is an example of an intensive property. For example the density of iron is 8.9kg/cm³ regardless of the mass of the iron block.

State and Path Functions

• State Functions: The value depends only on the current state of the substance.
• Path Functions: The value depends on the path taken to reach the final state.
• Examples of state functions are internal energy (U), enthalpy (H), and entropy (S).
• Examples of path functions are heat (q) and work (w).

Laws of Thermodynamics

• Zeroth Law: All parts of a system in thermodynamic equilibrium have the same temperature.
• First Law: Energy is conserved; it can not be created nor destroyed, it only changes form. The change in the internal energy of a system can only be altered by heat addition (or subtraction) or work done on (or by) the system. (∆U=q+w)
• Second Law: Processes tend to proceed in a direction that increases the total entropy of the universe involved. (∆Suniverse > 0)
• Third Law: The entropy of a perfect crystal approaches zero as the temperature approaches absolute zero (0K).

Thermodynamic Equilibrium

• A state of equilibrium is reached when the internal pressure and external pressure are equal.
• Heat is an example of energy transfer that is proportional to temperature difference.
• Work is a transfer of energy that causes a change of motion or position.

Other Concepts

• Internal Energy (U): The total of all microscopic energies within a system.
• Heat (q): Energy transfer due to a temperature difference.
• Work (w): Form of energy transfer associated with a change of position or motion.
• Temperature: A measure of the average kinetic energy of the molecules in a substance.

State Equation of an Ideal Gas

• pV = nRT, where p is pressure, V is volume, n is the number of moles, R is the ideal gas constant, and T is temperature.

Temperature

• Measured in Kelvin (K).
• The boiling point of pure water at standard pressure is 373.15K.
• The freezing point of pure water at standard pressure is 273.15K.

Thermochemical properties of fluids

• The properties of enthalpy and entropy depend on the pressure (not in the case of ideal gases), the volume of the substance, and the pathway.

Specific Heat Capacity

• The amount of heat required to change the temperature of a substance by a given amount.

Calculating Enthalpy Changes

• Enthalpy changes in chemical reactions can be derived when reactants and products are at a known temperature and pressure.

Hess's Law

• The enthalpy change for a reaction is independent of the pathway.

Reversible and Irreversible Processes

• A reversible process can be reversed by an infinitesimal change in conditions, as in equilibrium.
• Irreversible processes proceed in a specific direction and cannot easily be reversed.

Entropy

• A thermodynamic measure of disorder in a system.
• ∆S is the Change in entropy.
• Entropy changes in chemical processes can be determined quantitatively at given temperatures and pressures.

The Gibbs Function

• A useful function to assess whether a process can occur spontaneously at a constant temperature and pressure. -For the reaction to proceed spontaneously at a given temperature, ∆G < 0,∆G= ∆H − T∆S, where H is Enthalpy change and S entropy change.

Phase Diagrams

• Illustrate conditions of temperature and pressure under which various phases of a substance are stable.

Vapour Pressure

• The pressure exerted by a vapour in equilibrium with a liquid or solid at a given temperature.

Raoult's Law

• Ratio of the partial vapor pressure of a liquid in a mixture relative to the vapor pressure as a pure substance.

Henry's Law

• The partial vapor pressure of a solute in a mixture is proportional to its mole fraction where the proportionality constant is the vapor pressure of the pure solute.

Activities

• Activity is defined as the effective concentration of a substance in solution.
• The activity coefficient is the correction for interactions between solute, or different molecules.

The Free Energy of Mixing

• The Gibbs energy change when two or more components are mixed.

Reaction Rates

• Rate of a reaction can be determined experimentally.
• Rate laws show the relationship between reaction rate and concentrations of reactants and/or products.

Rate Order

• The exponent in a rate law equation tells which reactant it depends on.

Collision Theory & Concentration

• For a reaction to occur, collisions between molecules need to occur with sufficient energy and proper orientation.

Activation Energy

• The activation energy is minimum energy necessary to convert reactants to products in a reaction.

Arrhenius Equation

• Shows the relationship between reaction rate or constant, activation energy and temperature -The Arrhenius equation allows calculation of rate constants for reaction at different temperatures.

Reaction Mechanisms

• A series of steps that depict a chemical reaction occurring.
• Elementary steps in a reaction mechanism may be single or multiple molecules.

Rate-Determining Step

-In a reaction mechanism, the slowest step (with the highest activation energy) is the rate-determining step.

Catalysis

• A catalyst changes the rate of a reaction, without being consumed in the overall reaction.
• Catalysts lower the activation energy.

Ionic Equilibria

• Proton transfer equilibria: Reactions describing proton transfer between different molecules in aqueous solutions.
• H3O+ and OH−: The hydronium and hydroxide ions play crucial roles in determining acidity and basicity.
• pH: A measure of the concentration of H3O+ ions which is expressed in terms of the pH.
• pOH: A measure of the concentration of OH− ions which is expressed in terms of the pOH.

Test your understanding of the second law of thermodynamics and its implications for spontaneous processes. This quiz covers key concepts such as entropy, enthalpy changes, and the characteristics of reversible processes. Assess your knowledge of thermodynamic principles relevant to biochemical reactions and temperature effects.