Chapter 5 Thermochemistry and Thermodynamics
water vapor Terms and Definitions • Thermodynamics: Study of the exchange of heat, energy and work between a system and its surroundings. • System: that part of the universe of interest (reaction vessel, etc.) • Surroundings: rest of the universe. open system closed system isolated system energy nothing Exchange mass & energy
State and State Functions • State: that condition in which all variables are fixed and unvarying. • State Functions: variables (properties) whose value depends only on the state of the system. • Thermodynamics is concerned with how the state variables change during a change of state. Work • W is positive when work is done on the system by the surroundings; W is negative when the system does work and expend energy. • Some types of work: a) Mechanical work — exert a force through a distance. b) Work of expansion of a gas under constant pressure (PV work) c) Electrical work, W = charge x voltage
Enthalpy • The first Law of Thermodynamics: the total energy of the universe (or any isolated) system is constant. Energy can neither be created nor destroyed but can be converted from one form to another. • Enthalpy (H): heat transferred between the system & surroundings carried out under constant pressure. • Only the change in enthalpy can be measured: DH = Hfinal - Hinitial
Energy Exchange • Endothermic: system absorbs heat from surroundings (DH = +). • Exothermic: system transfers heat to the surroundings (DH = -). • An endothermic reaction feels cold; an exothermic reaction feels hot. Hproducts < HreactantsDH < 0 (negative) Exothermic Hproducts > HreactantsDH > 0 (positive) Endothermic
H2O (s) H2O (l) Enthalpy and the First Law of Thermodynamics Because DE = q + W At constant pressure, q = DH and W = -PDV DH = DE + PDV Thermochemical Equations 6.01 kJ are absorbed for every 1 mole of ice that melts at 0oC and 1 atm. DH = 6.01 kJ
The Specific Heats of Some Common Substances Calorimetry • Calorimetry: measurement of heat flow. • Calorimeter: apparatus that measures heat flow. • The heat capacity of a substance is the amount of energy required to raise the temperature of a substance by one degree Celsius. qrxn = -Ccal DT where Ccal is the heat capacity of the calorimeter (to be determined experimentally) Bomb Calorimetry: combustion processes
Phase Changes • Fusion: Solid Liquid. • DHfus(Molar Heat of Fusion) =Heat needed to convert one mole of a solid to a liquid at a particular T. • Vaporization: Liquid Vapor. • DHvap(Molar Heat of Vaporization) = Heat needed to convert one mole of a liquid to a gas at a particular T. • Sublimation: Solid Vapor. • DHsub(Molar Heat of Sublimation) = DHfus + DHvap.
Potential energy of hiker 1 and hiker 2 is the same even though they took different paths Hess’ Law • Hess’ law: When reactants are converted to products, DHis the same whether the reaction takes place in one step or in a series of steps. • State functions are properties that are determined by the state of the system, regardless of how that condition was achieved H1 = H2 + H3
Heats of Reactions • The standard enthalpy of reaction (DHqrxn) is the heat change of a reaction carried out at 1 atm. • Standard enthalpy of formation (DHqf) is the heat change that results when one mole of a compound is formed from its elements at a 1 atm. • DHqf of the most stable form of an element is zero.
Standard enthalpy of solution (DHqsoln) is the heat generated or absorbed when a certain amount of solute dissolves in a certain amount of solvent. Step 1 (lattice energy) + Step 2 (heat of hydration) = DHsoln
spontaneous heat + ice water at +10oC Spontaneous, irreversible spontaneous heat + ice water at -10oC Spontaneous, irreversible Reversible; equilibrium heat + ice water at 0oC Spontaneous processes • Spontaneous process is a process that occurs by itself (and the reverse does not occur by itself). • Reversible processes the systems must be at equilibrium Iceand water coexist at 0oC. Either process or can occur at equilibrium
Ice melting Water evaporating Ink drops in water What contributes to spontaneity? • Exothermic processes (heat is evolved). • Any process which increases randomness and disorder – Entropy
Entropy • The thermodynamic quantity which describes randomness and disorder is called entropy (S). • The Second Law of Thermodynamics: the entropy of the universe increases in a spontaneous process and remains unchanged in an equilibrium process. Spontaneous process:DSuniv = DSsys + DSsurr > 0 Equilibrium process:DSuniv = DSsys + DSsurr = 0
Exothermic Process DSsurr > 0 Endothermic Process DSsurr < 0 Entropy Changes in the System (DSsys) • The entropy of reaction is the entropy change for a reaction carried out at 1 atm and 25oC. DSqrxn = SnDSq(products) - SmDSq(reactants) 1) Gas has more entropy than liquid, which has more entropy than solid. 2) Melting or vaporization increases entropy. 3) In reaction, increasing the number of moles of a gas increases the entropy 4) Dissolving or mixing increases entropy; precipitation decreases entropy 5) Increasing the temperature increases entropy Entropy Changes in the Surroundings (DSsurr)
Standard States * The most stale allotropic form at 25oC and 1 atm Gibbs Free Energy (G) • For a constant-T process, the change in Gibbs free energy (DG) is • DG = DHsys - TDSsys 1) If DG is negative, there is a release of usable energy the reaction is spontaneous. 2) If DG is positive, the reaction is not spontaneous. 3) If DG is zero, the reaction is at equilibrium. • The free energy change of reaction(DGqrxn)is the free energy change for a reaction when it occurs under standard-state conditions. • DGqrxn = SnDGqf(products) – SmDGqf(reactants) • The free energy change of reaction(DGqf) is the free energy change that occurs when 1 mole of the compound is formed from its elements in their standard states.
Rate 40.79 kJ 373 K = 109 J/K DS = = Time DH T Gibbs Free Energy and Phase Transitions DGq = 0 = DHq – TDSq H2O(l) H2O(g) Gibbs Free Energy and Chemical Equilibrium DGq = -RT lnK
DGq = +29 kJ K < 1 Alanine + Glycine Alanylglycine ATP + H2O + Alanine + Glycine ADP + H3PO4 + Alanylglycine Adenosine diphosphate Example of the change on Gibbs Free Energy DGq = -2 kJ K > 1 Adenosine triphosphate