Thermodynamics: Energy in Biological Systems

• Energy considerations, also known as thermodynamics, dictate if chemical reactions or biological processes occur.
• Living organisms adhere to thermodynamic laws.
• Thermodynamics is crucial in understanding energy-requiring processes, energy needs, conversion to useful stores, the energy derived from food, and how organisms prevent complex molecules from breaking down.
• Classical thermodynamics studies energy transformations when a system changes states.
• Systems aim for the lowest energy state, reducing free energy in spontaneous processes.

Energy

• Energy lacks a physical form, making it an intrinsic property.
• Energy is the capacity to do work.
• Energy is described as being transformed rather than used in energetic processes.
• The Joule (J) is the SI unit of energy.
• One joule is the work done by a force of one newton moving an object one meter.
• Potential, kinetic, internal, chemical, electrochemical energy, and osmotic work are different types of energy.
• Thermodynamics examines energy conversion, the relationship between energy and heat, and the availability of energy for work.

Thermodynamics Laws

• The First Law states energy cannot be created or destroyed, only transformed, and the total energy of a system and its surroundings remains constant.
• Thermodynamics studies energy transformations.
• A "system" refers to the defined region of interest for tracking energy changes.
• Open, closed, and isolated systems are defined by whether they allow mass and/or heat transfer.
• The Second Law states that for a spontaneous process, the entropy of the thermodynamic universe must increase.

Entropy

• Entropy is the degree of order in a system.
• The First Law of Thermodynamics explains that energy is neither created nor destroyed, but changes form.
• The Second Law of Thermodynamics says that for spontaneous reactions, the entropy of the thermodynamic universe must increase.
• Equilibrium is a stable state where all influences are balanced, resulting in a unchanging system.
• Chemical equilibrium occurs when forward and reverse reactions happen at equal rates maintaining constant reactant and product concentrations.
• Equilibrium is the state of lowest energy and a state where no work can be done.
• A + B <--> AB is a key concept of chemical equilibrium
• Kᴀ = [AB]/[A]x[B] where Kᴀ is the equilibrium constant
• ΔG = -RT In Kᴀ where ΔG is the change in Gibbs free energy, R is the gas constant, T is temperature
• Reactions are favourable if ΔG 0
• ΔG = ΔH - TΔS
• ΔH is the change of enthalpy and ΔS is the change of entropy
• If the Kᴀ value for a chemical reaction is large (greater than 1), then the reaction is favourable.

Work, Energy, Force

• Mechanical work is a form of energy (Potential energy)
• Work done = force x distance moved
• Thermodynamics define the minimum amount of work required
• Force is a vector quantity causing acceleration
• Gravitational, magnetic, electrostatic, and gas pressure are examples of types of force.
• Force = mass x acceleration
• The Newton (N) is the SI unit of force
• One newton accelerates a mass of one kilogram at one metre per second squared.
• When weight falls potential energy is converted to kinetic energy which is converted to heat energy.
• The second law of thermodynamics tells us that not all forms of energy are interchangeable.
• Energy converted to heat cannot be reverted back to potential energy completely.

Molecular Forces

• Like electrostatic charges repel, while opposite charges attract.
• Chemical bonds are essentially electrostatic.
• Van der Waals forces are optimal ways of packing proteins
• Net force is a result of the Leonard-Jones Potential: energy = A/r¹² - B/r⁶
• Enthalpy is chemical bond energy (symbol H), deltaH is changes in enthalpy.
• Energy is released when bonds are formed, stronger bonds release more energy.
• Breaking bonds requires energy; stronger bonds require more energy.
• Bond energy/enthalpy of a reaction = the energy released when a bond is formed, or the energy required to break it.
• The enthalpy change (ΔH) of a chemical reaction is the sum of energy used when bonds are broken during the reaction and the energy released when new bonds are formed.
• Reactions where heath is lost are exothermic (ΔΗ is negative).
• Endothermic reactions are when heat is taken up (ΔΗ is positive).
• Reactions with negative ΔΗ are more likely to occur spontaneously.
• Some reactions with positive ΔΗ do occur.
• ΔG = ΔH - TAS is used to know whether a reaction will occur
• Second law: all processes must increase the entropy of the universe.
• Entropy is the tendency to disorder/increased entropy.
• Reactions where entropy increases (ΔS is positive) are favourable.
• Entropy (S) is the degree of disorder in a system, a measure of the probability of a state, the number of degrees of freedom/sub states, the information content of a system, and a universal parameter which tends to increase
• Units are J mol⁻¹ K⁻¹
• Entropy increases in process such as the melting of a solid, or boiling of a liquid and the expansion of a gas into a larger volume
• Entropy increases during the dilution of a solution– e.g., osmosis.
• A chemical reaction in which the number of molecules increases increases entropy
• Protein denaturation increases entropy.
• S = Kʙ ln N where N is the number of possible states of the system, Kʙ is the Boltzmann constant (1.38x10⁻²³ J K⁻¹)
• The Second Law of Thermodynamics explains the energy of entropy.
• Organisms do not defy the Second Law of Thermodynamics.
• Cells extract energy from food oxidation to maintain order and do work, increasing the universe's entropy by releasing molecules like CO2 and heat into the environment.
• Hydrophobic groups cluster in water due to van der Waals forces.
• Hydrophobic surfaces decreases water entropy by causing organized networks (clathrate structures).
• Clustering minimizes the clathrate water structures, reducing exposed hydrophobic surface area and decreasing unfavorable entropy.