Chapter 1.3: Physical Foundations

• Small molecules, macromolecules, and supramolecular complexes are continuously synthesized and broken down.
• Living cells maintain themselves in a dynamic steady state distant from equilibrium.
• Maintaining steady state requires the constant investment of energy.
• A system is all the constituent reactants and products, the solvent that contains them, and the immediate atmosphere.
• The universe is a system and its surroundings.
• The types of systems are isolated, closed, and open.
• Isolated systems exchange neither matter nor energy with its surroundings.
• Closed systems exchange energy, but not matter with its surroundings.
• Open systems exchange both energy and matter with its surroundings.
• A living organism is an open system.
• The first law of thermodynamics is that in any physical or chemical change, the total amount of energy in the universe remains constant, although the form of the energy may change.
• Potential energy for cells comes from nutrients in the environment and sunlight.
• In the cell, potential energy is converted into chemical transformations, heat, increased entropy in the surroundings, and decreased randomness in the cell.
• Cellular work done includes chemical synthesis, mechanical work, osmotic and electrical gradients, light production, and genetic information transfer.
• Entropy is increased in the surroundings of the cell when metabolism produces compounds simpler than the initial fuel molecules.
• Entropy is decreased in the cell when simple compounds polymerize to form information rich macromolecules.
• Photoautotrophs undergo the light driven reduction of CO2., as demonstrated by the equation 6CO2 + 6H2O ----> C6H12O6 + 6O2.
• Chemotrophs undergo energy yielding oxidation of glucose, as seen with the equation C6H12O6 + 6O2 ---> 6CO2 + 6H2O + energy.
• Autotrophs and heterotrophs participate in global cycles of O2 and CO2, driven by sunlight, making these two groups interdependent.
• Oxidation reduction reactions is where one reactant is oxidized, or loses electrons, as another is reduced, or gains electrons.
• Redox reactions describe reactions involved in electron flow.
• The second law of thermodynamics is that randomness in the universe is constantly increasing.
• Entropy, S, represents the randomness or disorder of the components of a chemical system.
• Enthalpy, H, is heat content, which is roughly reflected by the number and kinds of bonds.
• Free energy, G, of a closed system is H-TS where H represents the enthalpy, T the absolute temperature, and S the entropy.
• Delta H reflects the total energy change in a chemical reaction, as it is the measure of the number and kinds of bonds that are made and broken.
• Free energy change is delta G = delta H - T delta S where delta H is negative for a reaction that releases heat, and delta S is positive for a reaction that increases the system's randomness.
• Spontaneous reactions occur when delta G is negative.
• Endergonic, energy requiring, reactions are often coupled to exergonic, energy releasing, reactions.
• The breakage of phosphoanhydride bonds in ATP is highly exergonic.
• Breaking one phosphate off of ATP gives inorganic phosphate and ADP.
• Breaking two phosphates off of ATP gives inorganic pyrophosphate and AMP.
• ATP is adenosine triphosphate.
• Free energy change, or delta G, is the amount of energy available to do work.
• Delta G is always less than the theoretical amount of energy released.
• In closed systems, chemical reactions proceed spontaneously until equilibrium is reached.
• If under a given set of conditions, the reaction A --> B occurs with delta G is -14 and the reaction C --> B occurs when delta G is 16. This means that the conversion of A to C is exergonic.
• Equilibrium constant. When [A] is the concentration of A, [B] is the concentration B, and so on, then equilibrium is reached.

  
• Mass action ratio, Q, is the ratio of product concentration to reactant concentrations at a given time.
• Q can be calculated to determine how far the reaction is from equilibrium.