Thermodynamics Notes PDF

Thermodynamics Farah Mustafa Department of Biochemistry Rm: IE-179; Office: 03-713-7509 Mobile: 050-760-5516 Definition of Thermodynamics • Greek: thermos = heat; dynamic = change • Study of the relationship between energy, heat, and work. Thermodynamics • All living organisms require energy En...

Thermodynamics Farah Mustafa Department of Biochemistry Rm: IE-179; Office: 03-713-7509 Mobile: 050-760-5516 Definition of Thermodynamics • Greek: thermos = heat; dynamic = change • Study of the relationship between energy, heat, and work. Thermodynamics • All living organisms require energy Energy • Energy is defined as the ability to do work • Heat is energy in transfer between a system and its surroundings other than by work • Thermodynamics is the study of:  laws that govern the conversion of energy from one form to another  The direction in which heat will flow and  The availability of energy to do work Work Heat Lecture Outline • • • • • Definition of Thermodynamics Forms of Energy: Potential and Kinetic Energy Conversion Thermodynamic Parameters and Systems Laws of Thermodynamics  First Law: Enthalpy and Hess’s Law Second Law: Entropy and Order • Gibb’s Free Energy • Exothermic and Endothermic Reactions • Favorable and Unfavorable Reactions Forms of Energy • Found in different forms, such as light, heat, sound and motion, etc. • There are many forms of energy, but they can all be put into two categories: 1. Kinetic 2. Potential Kinetic verses Potential Energy • Kinetic energy. Energy in action, energy of motion––of waves, electrons, atoms, molecules, substances, objects, etc. • Potential energy is stored energy and the energy of position. Forms of Kinetic and Potential Energy Form of Energy Kinetic/Potential Object in Motion /Stored Energy Heat Kinetic Atoms/Molecules Chemical Potential Chemical Bonds Electrical Kinetic Electrons/Ions Gravitational Potential Gravity Optical (Radiant) Kinetic Photons Electrostatic Potential Coulomb Wind Kinetic Atoms/Molecules Nuclear Potential Nuclear binding force http://chemistry.osu.edu/~woodward/ch121/ch5_work.htm Energy Conversion • Potential energy can be converted to kinetic energy • It is kinetic energy that does work Thermodynamics Systems • A Thermodynamic System is that part of the universe that is under consideration • A real or imaginary boundary separates the system from the rest of the universe, which is referred to as the surrounding  System: the process/reaction whose thermodynamic change is being studied  Surrounding: the part of the universe that interacts with the system. Thermodynamic Potential • The Thermodynamic Parameters define the thermodynamic potential of a system.    Enthalpy Entropy Free Energy • The Laws of Thermodynamics deal with these parameters Thermodynamic Parameters • Some commonly considered parameters are: Mechanical parameters: Pressure: p Volume: V Thermodynamic parameters: Temperature: T Enthalpy: H Entropy: S First Law of Thermodynamics • Conservation of Energy Energy cannot be created or destroyed, merely transformed from one form to another • Therefore, for any physical or chemical change, the total amount of energy in the universe should remains constant • Energy may change form or it may be transferred from one region to another, but it cannot be created or destroyed Enthalpy (H) • Enthalpy = Heat content (H) • Is property of a system that reflects its capacity to exchange heat with the surrounding • The total enthalpy of a system cannot be measured directly • However, the enthalpy change of a system can be measured • Enthalpy change is defined by the following equation: ΔH = H final –H initial If Hf > Hi ΔH + If Hf < Hi ΔH - Directionality of Heat Transfer • Heat always transfers from hotter object to cooler one. • EXOthermic: heat transfers from SYSTEM to SURROUNDINGS T(system) goes down T(surr) goes up • Heat always transfers from hotter object to cooler one. • ENDOthermic: heat transfers from SURROUNDINGS to the SYSTEM. T(system) goes up T (surr) goes down Joule: Unit of Energy, Work, or Amount of Heat • 1 calorie = Heat required to raise the temperature of 1 gm of H2O by 1 oC • 1000 cal = 1 kilocalorie = 1 kcal • 1 kcal = 1 Calorie (a food “calorie”) • The SI unit for energy, work, or heat is Joule 1 cal = 4.184 joules James P. Joule 1818-1889 Hess's Law • If a reaction is carried out in a series of steps, ΔH for the reaction will be equal to the sum of the enthalpy changes for the individual steps. • It is often possible to calculate ΔH of a reaction from listed ΔH values of other reactions. CH4(g) + 2O2(g) 2H2O(g) CH4(g) + 2O2 (g) CO2(g) + 2H2O(g) 2H2O(l) CO2 (g) + 2H2O(l) ΔH = -802 kJ ΔH = -88 kJ ΔH = -890 kJ Second Law of Thermodynamics The Entropy Lab http://www.cartoonstock.com/directory/t/thermodynamics.asp Second Law of Thermodynamics • The entropy, S, of the universe increases (real, spontaneous processes). • Thus, highly ordered systems have low entropy and with time, they move towards more disorder (more entropy) • The more random the distribution of molecules, the greater the entropy • Which has the greatest entropy? Solids, liquids, or gasses S gas> S liquid> S solid Entropy • Entropy can be thought of as a measure of the randomness of a system. • It is related to the various modes of motion in molecules. • Like total energy, E, and enthalpy, H, entropy is a state function. • Therefore, S = Sfinal  Sinitial Entropy on the Molecular Scale Implications: • more particles -> more states • higher T -> more energy states • less structure (gas vs solid) -> more states -> more entropy -> more entropy -> more entropy Entropy on the Molecular Scale • What happens to entropy when you increase:    Temperature Volume (gases) The number of independently moving molecules • Increase in these factors increases entropy Entropy Changes • In general, entropy increases with the freedom of motion of molecules. For e.g., when – the number of gas molecules increase – The number of moles of a substance increase Predict the Sign of ΔS Predict the sign of S PROCESS ice -> liquid water water vapor -> liquid water cold water -> hot water 3 H2(g) + N2(g) -> 2 NH3(g) pressure of a gas increases volume of a gas increases S ? ? ? ? ? ? Entropy Change in the Universe • The universe is composed of the system and the surroundings. Suniverse = Ssystem + Ssurroundings • Entropy change for a process: Suniv = Ssys + Ssurr > 0  process is spontaneous Suniv = Ssys + Ssurr = 0  process is at equilibrium Suniv = Ssys + Ssurr < 0  process is non-spontaneous Thermodynamic Free Energy • It is the maximum amount of chemical energy derived from a spontaneous reaction that can be utilized to do work or to drive a nonspontaneous process. Or • It is the minimum amount of energy that must be supplied to make a non-spontaneous reaction occur. Gibbs Free Energy • The amount of energy a system has available to do useful work is called the Gibbs Free Energy (G) • This change is symbolized as ΔG: ΔG = ΔH - TΔS • Predicts the direction of a spontaneous reaction • By knowing the sign (+ or -) of S and H, one can get the sign of G and determine if a reaction is spontaneous Reactions Exothermic or Negative ΔG Negative lies G Equilibrium towards completion Endothermic or Positive ΔG Positivewill G Reaction not occur without energy input Gibbs Free Energy 1. If G is negative, the forward reaction is spontaneous. 2. If G is 0, the system is at equilibrium. 3. If G is positive, the reaction is spontaneous in the reverse direction. Characteristics of Reactions Exothermis Reactions Endothermic Reactions • ΔG is negative • ΔG is positive • The final state has less energy than the starting state • The final state has more energy than the starting state • Energy is released  Enthalpy Decreases • Energy is absorbed  Enthalpy Increases • The process occurs spontaneously, although it may occur slowly • Non-spontaneous. Without energy input, the process will not occur • Often the starting state (reactants) is more ordered and complex than the final state (products) Entropy Increases • Often the final state (products) is more ordered and complex than the starting state (reactants)  Entropy Decreases • In living systems, exergonic processes often decrease order or release potential energy • Living systems generally use endergonic processes to do work and to maintain order • Some energy is lost as heat • Some energy is lost as heat Types of Work Done by Cells • Chemical Work – Anabolism – The formation of many molecules needed by the cell requires energy. • Mechanical Work – Change physical locations of organisms, cells and internal structures • Transport Work – Movement of molecules and ions into and out of cells against electrical and concentration gradients Thermodynamics Laws Apply to Biological Systems Energy Used by Living Things • Energy is stored in the chemical bonds of biological molecules. E.g., glucose catabolism (exothermic) • Stored energy is released so work can be accomplished. E.g., glucose anabolism (endothermic) 2nd Law Application • Increase in Entropy, i.e. – Useful energy is lost when energy changes forms Energetically Favorable and Unfavorable Reactions An Unfavorable Metabolic Reaction is Driven by a Favorable One A B G = +10kJ/mol (Endergonic reaction) C D G = -20kJ/mol (Exergonic reaction) A+C B + D G = -10kJ/mol (coupling reaction) Example Glucose + Pi ATP + H2O Glucose + ATP Glucose-6-phosphate ADP+ Pi G = +14kJ/mol G = -31kJ/mol Glucose-6-phosphate G = -17kJ/mol