ENTHALPY ENTROPY AND GIBBS LAW OF FREE ENERGY

ENTHALPY, ENTROPY AND GIBBS LAW OF FREE ENERGY Dr Nadeem Asad 2 -8 -2014

Energy reactions Review the Energy Diagram Endothermic have a high startup energy Exothermic have a low start up energy

3 Heat and Temperature • Heat is energy that is transferred from one object to another due to a difference in temperature • Temperature is a measure of the average kinetic energy of a body • Heat is always transferred from objects at a higher temperature to those at a lower temperature 3

Energy Curve • Activation energy: energy required to get the reaction to move forward • Energy released or absorbed is noted at the end of the curve

Factors that effect the Reaction Rate • Temperature: Endothermic vs Exothermic reactions • Concentration: Increase the reactants will increase the products • Surface Area: Smaller particles have large surface area • Catalysts and inhibitors

Enthalpy Delta H = Hproducts - Hreactants Delta H represents the transfer of heat

Exothermic vs endothermic reactions Exothermic Reaction: Reaction A process that releases heat to its surroundings. Products have less energy than the reactants Endothermic Reaction : A process that absorbs heat from the surroundings. Products have more energy than the reactants.

Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. 2 H 2 (g) + O 2 (g) H 2 O (g) 2 H 2 O (l) + energy Endothermic process is any process in which heat has to be supplied to the system from the surroundings. energy + 2 Hg. O (s) energy + H 2 O (s) 2 Hg (l) + O 2 (g) H 2 O (l) 6. 2

Combustion and neutralization processes Combustion Exothermic reaction General Combustion Reaction Formula: Compound (usually hydrocarbon) + O 2 CO 2 + H 2 O + energy CH 4 + 2 O 2 CO 2 + 2 H 2 O + 890 k. J ∆H = -890 k. J Neutralization Exothermic reaction Acid + Base Salt + Water + energy HCl + Na. OH Na. Cl + H 2 O + 57. 3 k. J ∆H = -57. 3 k. J

Activation energy and enthalpy diagrams Exothermic Reactions Products more stable than reactants (lower energy). ΔH = Hproducts – Hreactants Since the products have less energy than the reactants, the ΔH value is negative. Endothermic Reactions Products less stable than reactants (higher energy) ΔH = Hproducts – Hreactants Since the products have more energy than the reactants, the ΔH value is positive.

Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. H = H (products) – H (reactants) H = heat given off or absorbed during a reaction at constant pressure Hproducts < Hreactants H < 0 Hproducts > Hreactants H > 0 6. 4

Standard enthalpy of reaction Standard Enthalpy Change of Reaction (∆H): The heat energy exchanged with the surroundings when a reaction happens under standard conditions (NOT STP… see below). Since the enthalpy change for any given reaction will vary with the conditions, esp. concentration of chemicals, ΔH are measured under standard conditions: • • pressure = 101. 3 k. Pa temperature = 25ºC = 298 K Concentrations of 1 mol dm-3 The most thermodynamically stable allotrope (which in the case of carbon is graphite) Only ΔH can be measured, not H for the initial or final state of a system.

Terminology of rate of reaction (Δho) Pseudonyms (other names) for H l Heat of Reaction: Hrxn heat produced in a chemical reaction l Heat of Combustion: Hcomb heat produced by a combustion reaction l Heat of Neutralization: heat produced in a neutralization reaction (when an acid and base are mixed to get water, p. H = 7) l Heat of solution: Hsol heat produced by when something dissolves l Heat of Fusion: Hfus heat produced when something melts l Heat of Vaporization: Hvap heat produced when something evaporates l Heat of Sublimation: Hsub heat produced when something sublimes l Heat of formation: Hf change in enthalpy that accompanies the formation of 1 mole of compound from it’s elements (this has special uses in chemistry…)

Spontaneous Reactions • If the reaction moves forward as written without an intervention the reaction is said to be spontaneous. • Remember Delta H (enthalpy) • Negative signs represent exothermic • Always spontaneous • Positive sign represents endothermic • Sometimes spontaneous

Examples H 2 O (s) --> H 2 O (l) delta H = +6 k. J 2 Na(s) + Cl 2 --> 2 Na. Cl (s) delta H = -822 k. J The above reactions are all spontaneous. But most endothermic rxn are Nonspontaneous.

Delta H • Most enthalpy reactions that are endothermic rxn are not spontaneous. This is because the reaction has a high energy barrier. • Only because of a large difference between entropy and enthalpy can the reaction happen.

Average bond enthalpy Enthalpy changes of reactions are the result of bonds breaking and new bonds being formed. Remember… • Breaking bonds requires energy • Forming new bonds releases energy Bond enthalpy is the energy required to break one mole of a certain type of bond in the gaseous state averaged across a variety of compounds. FYI: Bond enthalpies for unlike atoms will be affected by surrounding bonds and will be slightly different in different compounds so average bond enthalpies are used.

• H = ∑ (energy required to break bonds) OR – ∑(energy released when bonds are formed) – ∑(bond enthalpy of products) • H = ∑ (bond enthalpy of reactants)

19 Energy Changes in endothermic and exothermic processes In an endothermic reaction there is more energy required to break bonds than is released when bonds are formed. The opposite is true in an exothermic reaction. 19

Exothermic and endothermic reactions in light of average bond enthalpies If the amount of energy required to break the bonds in the reactants is greater than the amount of energy released when bonds are formed in the products, the reaction is endothermic. average bond enthalpy reactants > average bond enthalpy products If the amount of energy required to break the bonds in the reactants is less than the amount of energy released when bonds are formed in the products, the reaction is exothermic. average bond enthalpy reactants < average bond enthalpy products

Calorimetry • Calorimetry involves the measurement of heat changes that occur in chemical processes or reactions. Determines the ΔH by measuring temp Δ's created from the rxn • The heat change that occurs when a substance absorbs or releases energy is really a function of three quantities: • The mass • The temperature change • The heat capacity of the material 21

Entropy • Entropy is the measurement of disorder of particles in a reaction. Equation: Delta S = SProduct - Sreactant *not easy to measure, but we can compare based on the states of matter

Comparing Entropy • Sproduct > Sreactants delta S is positive • Sproduct< Sreactants delta S is negative • Solids have a very low entropy • Liquids have a slightly higher entropy • Gases have a very high entropy

Example reactions • CO 2 (s) --> CO 2 (g) S > 0 • H 2 O (g) --> H 2 O (s) S<0 • 2 NH 3(g) --> N 2(g) + 2 H 2 (g) S>0

Law of Thermodynamics • States: in any spontaneous process, the overal entropy of the universe always increases. Suniverse = Sreaction + Ssurroundings

Law of Thermodynamics • Sreaction - calculated using the states of matter • Ssurround - calculated based on enthalpy • If the reaction is endothermic it is taking energy away from the surrounding making it negative • If the reaction is exothermic it is placing energy into the system and making it positive.

Gibbs Free Energy • Delta G = Delta H - T delta. S Compares entropy, enthalpy and temperature * If delta G is negative the reaction is spontaneous * If delta G is positive the reaction is nonspontaneous * If delta G is 0 the reaction is at equilibrium