Thermochemistry and Thermodynamics

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# Thermochemistry and Thermodynamics - PowerPoint PPT Presentation

##### Thermochemistry and Thermodynamics

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1. Chapter 5 Thermochemistry and Thermodynamics

2. 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

3. 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

4. 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

5. 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

6. 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

7. 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

8. 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.

9. 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

10. 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.

11. 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

12. 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

13. 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

14. 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

15. 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)

16. 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.

17. 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

18. 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