Thermodynamics: Energy and Systems

Thermodynamics Overview

• Study of energy and its transformations in physical and chemical processes.
• Theory is used to describe energy-related changes in reactions
• Impact the pharmaceutical applications

Foundations of Thermodynamics

• Deals with quantitative relationships of energy interconversions.
• Energy can be mechanical, chemical, electric, or radiant.
• Built on three "laws" based on empirical observations, not direct proofs.
• Expressed mathematically, conclusions align with observations.
• Laws apply to systems with numerous molecules.

Thermodynamic Systems

• A system is the specific part of the universe under study.
• Types of systems: open, closed, and isolated.
• Open System: Exchanges both energy and matter.
• Closed System: Exchanges only energy but not matter; mass remains constant.
• Isolated System: Exchanges neither energy nor matter.

Thermodynamics terms

• System involves directing attention to a region of the universe.
• Surroundings encompass everything external to the system.
• Boundary is the wall separating a system from its surroundings.

Basic Definitions

• A system is a well-defined portion of the universe.
• A boundary is a physical barrier separating the system from its surroundings.
• Work (W) is the transfer of energy to change the height of a weight
• Heat is the transfer of energy due to temperature differences between a system and its surroundings.
• Work and heat are only present at system boundaries during energy transfer.

Work

• Work is calculated as Force x Distance, or FΔd
• Force is calculated as Pressure x Area
• Work also equals P(AΔd) or PΔV
• PV work is also termed expansion work

Energy Transfer

• Systems exchange energy through work and heat.
• Heat involves transfering thermal energy.
• Work transfers energy and can create heat.
• Energy may enter the system as heat or work.
• Energy can be stored as potential energy (PE) and kinetic energy (KE) and withdrawn as work or heat.

First Law of Thermodynamics

• States energy is conserved; it can transform but not be created or destroyed.
• Total energy of a system and surroundings remains constant.
• Einstein's equation is Energy = (mass change) x (velocity of light)².
• Matter is energy; 1g equivalent to 9 x 10^13 units of energy.
• Heat (Q) and work (W) affect a system's internal energy.
• Change in internal energy (ΔE) equals Q + W, where E2 and E1 are the final and initial internal energies, respectively.

Isothermal and Adiabatic Processes

• Isothermal process maintains constant temperature during a reaction.
• This is achieved by a constant-temperature bath.
• Adiabatic process involves no heat exchange (dq = 0).
• Reactions in sealed Dewar flasks (vacuum bottles) are adiabatic.

Thermodynamic State

• Condition where measurable system properties have definite values.
• Defined by specifying temperature (T), pressure (P), and volume (V).
• The change in internal energy only depends on initial and final states
• Q and W are path-dependent, not state variables.

Thermochemistry

• Chemical and physical processes at atmospheric pressure are considered ideally constant.
• Heat exchange equals enthalpy change (Qp = ΔH).
• Enthalpy change remains a function of temperature.
• Exothermic reactions release heat (negative ΔH/Qp).
• Endothermic reactions absorb heat (positive ΔH/Qp).
• Heat exchange equals internal energy change in a container (Qv = ΔE).
• Thermochemistry studies heat changes in isothermal reactions to obtain ΔH and ΔE.
• In solution reactions, PΔV terms are negligible, ΔH ≈ ΔE.
• Gas reactions are not included in this approximation.

Applications of Thermochemistry

• Measuring reaction heats, especially using Hess's Law.
• Determining heats of reaction using bond energies.
• Analyzing heat-of-mixing processes.
• Calculating calories in foods.

Third Law of Thermodynamics

• States a pure, crystalline substance's entropy is zero at absolute zero.
• At absolute zero, crystal arrangement is perfectly ordered.
• A perfect crystal only has one configuration; entropy is zero [S = k ln(1) = 0].
• Reaching absolute zero temperature for any system in a finite number of steps is impossible.

Laws of Thermodynamics

• First law observes energy conservation during conversion.
• Second law relates to process probability based on system tendency toward energy equilibrium.
• Entropy explains process tendencies, originating from steam engine efficiency studies.
• Steam engines produce work only with temperature decrease and heat flow.

Efficiency of Heat Engine

• Converts thermal energy to mechanical energy.
• Heat is unavailable for isothermal work, and cannot fully converted to work
• Operates using two heat reservoirs at different temperatures.
• Part of the heat at the source is converted into work, the remainder is returned to the sink
• Efficiency equals W/Q, where W is the work, and Q is the heat

Temperature and heat flow in engines

• Carnot showed the engine efficiency depends on the equation W/Q_hot = (Q_hot - Q_cold) / Q_hot for cyclic processes.
• Heat flow follows temperature gradient; heat absorbed/rejected is related to temperatures.
• Lord Kelvin linked Q_hot/Q_cold to T_hot/T_cold to establish the Kelvin scale.
• In equation 1, Efficiency is quantified as Efficiency= (Qhot – Qcold) / Qhot = (Thot – Tcold) / Thot
• Engine efficiency increases as Thot increases and Tcold decreases.
• Reaching absolute zero on the Kelvin scale converts heat completely to work, but is impossible, meaning Heat not completely converted to work

Entropy

• Measure of disorder increases spontaneously.
• Increases with softer, less rigid solids, larger atoms, and complex molecular structures.

Efficiency of Heat Engine

• Increases with softer, less rigid solids, larger atoms, and complex molecular structures.
• Heat engine converts thermal energy into mechanical energy
• Heat is isothermally unavailable for work, it can never be converted
• Operates using two heat reservoirs at different temperatures.
• Part of heat at the source is converted into work, and rest is returned to the sink
• Efficiency equals W/Q, where W is the work, and Q is the heat

Predicting Entropy Changes

• In gasses, increased moles (∆S > 0) and decreased moles (∆S < 0) affect the level of entropy
• Phase changes can affect entropy levels AS > 0 and ΔΗ > 0
• Increasing volume leads to increase of entropy

Gibbs Free Energy

• Used in biological thermodynamics and chemical thermos systems.
• Determines process spontaneity.
• Knowledge of entropy changes is essential.
• ΔG = ΔH - TΔS
• ΔH relates to heat change of a reaction with positive values indicating endothermic characteristics
• ΔG measures change in a system's free energy.

Gibbs Free Energy and Spontaneity

• Spontaneous process occurs when ΔG < 0.
• Equilibrium occurs when ΔG = 0.
• Non-spontaneous process (as written) when ΔG > 0, but it is a spontaneous process in the reverse direction.