Thermodynamics

1. Thermodynamics Energy and Heat

2. Energy = the ability to do work or to produce heat • Kinetic energy: energy of motion • Potential energy: stored energy • Chemical potential energy = the energy stored in a substance because of its composition • Law of Conservation of Energy: in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed

3. Heat = the energy that is in the process of flowing from a warmer object to a cooler object • Measuring heat: • Calorie (cal) = the amount of heat required to raise the temperature of one gram of pure water by 1°C • Joule (J) = SI unit of heat and energy 1 calorie = 4.184 Joules

4. Heat Unit Conversions Convert the following: 1 calorie = 4.184 Joules • 230 calories = _____________ joules • 5.87 joules = ____________ calories • 52 kilocalories = ______________ joules • 4.56 kilojoules = ______________ calories

5. Specific Heat Specific heat (J/g x oC)= the amount of heat required to raise the temperature of one gram of that substance by one degree Celsius q = cm∆T where: q = energy c = specific heat capacity m = mass of sample in grams ∆T = temperature change in Celsius

6. Calorimetry • Calorimeter = an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process • You can use a calorimeter to determine the specific heat of an unknown metal • Measured mass of water has an initial temperature • Piece of hot metal is added • The metal transfers heat to the water until the metal and water attain the same temperature • Final temperature of the water is measured • Specific heat can be calculated

7. Thermochemistry Thermochemistry = the study of heat changes that accompany chemical reactions and phase changes • Universe = system + surroundings • System = the specific part of the universe that contains the reaction or process you wish to study. • Surroundings = everything in the universe other than the system

8. Thermochemistry Enthalpy (H) = the heat content of a system at constant pressure • Enthalpy (heat) of reaction (∆Hrxn) is the change in enthalpy for a reaction ΔHrxn = Hproducts – Hreactants • If positive, energy has gone INTO the reaction • If negative, energy has been RELEASED by the reaction

9. Endothermic vs. Exothermic • Endothermic reaction = chemical reaction that absorbs heat • A greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the product molecules • Exothermic reaction = chemical reaction that gives off heat • More energy is released forming new bonds than is required to break bonds in the initial reactants

10. Thermochemical Equations Thermochemical Equation = a balanced chemical equation that includes the physical states of all reactants and products and the energy change, usually expressed as change in enthalpy, ΔH. Heat pack equation 4Fe(s) + 3O2(g)  2Fe2O3(s) ΔH = -1625 kJ Cold pack equation NH4NO3(s)  NH4+(aq) + NO3-(aq) ΔH = 27kJ

11. Changes of State • Molar enthalpy (heat) of vaporization(ΔHvap) – the heat required to vaporize one mole of a liquid H2O(l)  H2O(g) ΔHvap = 40.7 kJ • Molar enthalpy (heat) of fusion (ΔHfus) • the heat required to melt one mole of a solid substance H2O(s)  H2O(l) ΔHfus = 6.01 kJ

12. Combustion ReactionsA Review Combustion reaction= a substance rapidly combines with oxygen to form either carbon dioxide and water or metal oxides • Hydrocarbons burning carbon dioxide and water • Example: CH4 (g) + 2O2 (g)  CO2(g) + 2H2O (g) • Metals burning  metal oxides • Example: 2Mg (s) + O2 (g)  2MgO (s)

13. Combustion and Welding When welding is done with an acetylene torch, acetylene combines with oxygen to form carbon dioxide and water. This reaction is exothermic and enough energy is released to melt metal 2C2H2(g) + 5O2(g)  4CO2(g) + 2H2O(g) + energy

14. Combustion and Challenger • Hydrogen gas and oxygen gas react when hydrogen is heated  forming water and releasing a large amount of energy 2H2(g) + O2 (g)  2 H2O (g) • Hydrogen tank propelled into the oxygen tank by leak • Combustion reaction above took place, thus destroying the shuttle

15. Calculating Enthalpy Change • Hess’s Lawstates that if you can add two or more thermochemical equations to produce a final equation for a reaction, then the sum of the enthalpy changes for the individual reactions is the enthalpy change for the final reaction.

16. Hess’s Law Let’s Walk Through a Problem… • Use thermochemical equations “a” and “b” below to determine ΔH for the decomposition of hydrogen peroxide (H2O2) 2 H2O2(l)  2H2O(l) + O2(g) • 2H2(g) + O2(g)  2H2O(l) ΔH = -572 kJ • H2(g) + O2(g)  H2O2(l) ΔH = -188 kJ

17. Heat of Formation Standard enthalpy (heat) of formation (∆Hf) = the change in enthalpy that accompanies the formation of one mole of the compound in its standard state from its constituent elements in their standard states 2S (s) + 3O2 (g)  2SO3 (g) ∆Hf = -396 kJ 396 kJ of heat are given off in this reaction

18. Heat of Reaction Heat of reaction (∆Hrxn)= amount of heat energy given off or absorbed in a particular chemical reaction for a given amount of reactants or products The standard heats of formation equations combine to produce the desired equation and its ∆Hrxn. ∆Hrxn = ∑∆Hf (products) - ∑∆Hf (reactants)

19. Heat of ReactionExample What is the heat of reaction for the following reaction? 2 Mg (s) + O2 (g)  2 MgO (s) Given: ∆Hf(Mg) = 0 kJ/mol ∆Hf(O2) = 0 kJ/mol ∆Hf(MgO) = -602 kJ/mol

20. Heat of ReactionExample (cont’d) ∆Hrxn = ∑∆Hf (products) - ∑∆Hf (reactants) = 2 (∆Hf(MgO) ) – [2(∆Hf(Mg) + ∆Hf(O2)] = 2 mol (-602 kJ/mol) – [0 + 0] = -1204 kJ for 2 moles of MgO 1204 kJ are given off to make 2 moles of MgO

21. 15.5 – Reaction Spontaneity • Spontaneous process – any physical or chemical change that once begun, occurs with no outside intervention (but some outside energy may be necessary to get it started) • Entropy (S) is a measure of the disorder of a system.

22. Second Law of Thermodynamics • “Spontaneous processes always proceed in such a way that the entropy of the universe increases” • Systems prefer to be in disorder.

23. Activation Energy • Activation energy (Ea) = the minimum amount of energy that reacting particles must have to cause a chemical reaction

24. Activation Energy Endothermic Reactions

25. Activation EnergyExothermic Reactions

26. Factors that Influence Reaction Rates Nature of Reactants Concentration Surface Area Temperature Catalysts

27. Nature of Reactants • Some elements are more reactive than others • Example: sodium is more reactive than calcium so the reaction of sodium and water occurs at a faster rate than calcium and water • More reactive elements  faster rate of reaction

28. Concentration Higher concentration means there are more particles around in the reaction mix to collide Reactant A “finds” reactant B more easily if there are more A and B particles near each other Higher concentration of reactants  faster rate of reaction

29. Surface Area • Think about dissolving sugar: • Sugar cube + water  dissolves slowly • Granulated sugar + water  dissolves faster • Sugar cube overall has smaller surface area than sugar granules exposed to water for dissolving reaction to occur • Larger surface area  faster rate of reaction

30. Temperature Think about dissolving cocoa mix – is it easier when the water is cold or hot? Increased temperature increases the kinetic motion of particles  particles can collide easier  more collisions  reaction Increased temperature  faster rate of reaction

31. Catalysts Catalyst = a substance that increases the rate of a chemical reaction without itself being consumed in the reaction Catalysts reduce the activation energy, making it easier for the reaction to occur

32. ***LeChâtelier’s Principle • If a stress is applied to a reversible system at equilibrium, the system shifts in the direction that relieves the stress. (Think of a see-saw) • Example: • Adding more reactants leads to production of more products (shifts the equilibrium to the right) • Adding more of the products leads to an increase in the reverse reaction, to produce more reactants (shifts the equilibrium to the left)