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Thermodynamics

Thermodynamics. Antoine Lavoisier [1743-94]. Julius Robert Meyer. James Joule : 1818~1889. Hermann von Helmholtz (1821-1894). Rudolf Clausius (1822--1888),. Fuel is burnt to produce energy - combustion (e.g. when fossil fuels are burnt)

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Thermodynamics

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  1. Thermodynamics

  2. Antoine Lavoisier [1743-94]

  3. Julius Robert Meyer

  4. James Joule : 1818~1889

  5. Hermann von Helmholtz (1821-1894)

  6. Rudolf Clausius (1822--1888),

  7. Fuel is burnt to produce energy - combustion (e.g. when fossil fuels are burnt) CH4(g) + 2O2(g) CO2(g) + 2H2O(l) + energy THERMOCHEMISTRY The study of heat released or required by chemical reactions

  8. Energy Kinetic energy (EK) Potential energy (EP) Energy due to motion Energy due to position (stored energy) What is Energy?

  9. Total Energy = Kinetic Energy + Potential Energy E = EK + EP Kinetic energy & potential energy are interchangeable Ball thrown upwards slows & loses kinetic energy but gains potential energy The reverse happens as it falls back to the ground

  10. Law of Conservation of Energy: the total energy of the universe is constant and can neither be created nor destroyed; it can only be transformed. The internal energy, U, of a sample is the sum of all the kinetic and potential energies of all the atoms and molecules in a sample i.e. it is the total energy of all the atoms and molecules in a sample

  11. SYSTEM OPEN ISOLATED CLOSED Systems & Surroundings In thermodynamics, the world is divided into a system and its surroundings A system is the part of the world we want to study (e.g. a reaction mixture in a flask) The surroundings consist of everything else outside the system

  12. SURROUNDINGS SYSTEM Thermochemistry is the study of heat change in chemical reactions. The system is the specific part of the universe that is of interest in the study. open closed isolated energy nothing Exchange: mass & energy

  13. OPEN SYSTEM: can exchange both matter and energy with the surroundings (e.g. open reaction flask, rocket engine) CLOSED SYSTEM: can exchange only energy with the surroundings (matter remains fixed) e.g. a sealed reaction flask ISOLATED SYSTEM: can exchange neither energy nor matter with its surroundings (e.g. a thermos flask)

  14. HEAT and WORK HEAT is the energy that transfers from one object to another when the two things are at different temperatures and in some kind of contact e.g. kettle heats on a gas flame cup of tea cools down (loses energy as heat) Thermal motion (random molecular motion) is increased by heat energy i.e. heat stimulates thermal motion

  15. Work is the transfer of energy that takes place when an object is moved against an opposing force i.e. a system does work when it expands against an external pressure Car engine: petrol burns & produces gases which push out pistons in the engine and transfer energy to the wheels of car • Work stimulates uniform motion • Heat and work can be considered as energy in transit

  16. UNITS OF ENERGY S.I. unit of energy is the joule (J) Heat and work ( energy in transit) also measured in joules 1 kJ (kilojoule) = 103 J Calorie (cal): 1 cal is the energy needed to raise the temperature of 1g of water by 1oC 1 cal = 4.184 J

  17. U = Ufinal - Uinitial U change in the internal energy INTERNAL ENERGY (U) Internal energy changes when energy enters or leaves a system Heat and work are 2 equivalent ways of changing the internal energy of a system

  18. Energy supplied to system as heat Energy supplied to system as work Change in internal energy = + U like reserves of a bank: bank accepts deposits or withdrawals in two currencies (q & w) but stores them as common fund, U. q q w w U U = q (heat) + w (work)

  19. First Law of Thermodynamics: the internal energy of an isolated system is constant Signs (+/-) will tell you if energy is entering or leaving a system + indicates energy enters a system - indicates energy leaves a system

  20. WORK • An important form of work is EXPANSION WORK i.e. the work done when a system changes size and pushes against an external force • e.g. the work done by hot gases in an engine as they push back the pistons HEAT In a system that can’t expand, no work is done (w = 0) U = q + w when w = 0, U = q (at constant volume)

  21. A change in internal energy can be identified with the heat supplied at constant volume ENTHALPY (H) (comes from Greek for “heat inside”) • the change in internal energy is not equal to the heat supplied when the system is free to change its volume • some of the energy can return to the surroundings as expansion work •  U < q

  22. The heat supplied is equal to the change in another thermodynamic property called enthalpy (H) • i.e. H = q • this relation is only valid at constant pressure As most reactions in chemistry take place at constant pressure we can say that: A change in enthalpy = heat supplied

  23. Burning fossil fuels is an exothermic reaction EXOTHERMIC & ENDOTHERMIC REACTIONS Exothermic process: a change (e.g. a chemical reaction) that releases heat. A release of heat corresponds to a decrease in enthalpy Exothermic process: H < 0 (at constant pressure)

  24. Endothermic process: a change (e.g. a chemical reaction) that requires (or absorbs) heat. An input of heat corresponds to an increase in enthalpy Endothermic process: H > 0 (at constant pressure) Forming Na+ and Cl- ions from NaCl is an endothermic process Photosynthesis is an endothermic reaction (requires energy input from sun)

  25. reaction reaction Measuring Heat Exothermic reaction, heat given off & temperature of water rises Endothermic reaction, heat taken in & temperature of water drops

  26. How do we relate change in temp. to the energy transferred? Heat capacity (J/oC) = heat supplied (J) temperature (oC) Heat Capacity = heat required to raise temp. of an object by 1oC • more heat is required to raise the temp. of a large sample of a substance by 1oC than is needed for a smaller sample

  27. b. Specific heat is the amount of heat required to raise the temperature of 1 kg of a material by one degree (C or K). 1) C water = 4184 J / kg C 2) C sand = 664 J / kg C This is why land heats up quickly during the day and cools quickly at night and why water takes longer.

  28. Why does water have such a high specific heat? Water molecules form strong bonds with each other; therefore it takes more heat energy to break them. Metals have weak bonds and do not need as much energy to break them.

  29. How to calculate changes in thermal energy Q = m x T x Cp Q = change in thermal energy m = mass of substance T = change in temperature (Tf – Ti) Cp = specific heat of substance

  30. c. A calorimeter is used to help measure the specific heat of a substance. First, mass and temperature of water are measured Knowing its Q value, its mass, and its T, its Cp can be calculated Then heated sample is put inside and heat flows into water This gives the heat lost by the substance T is measured for water to help get its heat gain

  31. Specific Heat Capacity (Cs) J / oC / g Heat capacity J / oC = = Mass g Specific heat capacity is the quantity of energy required to change the temperature of a 1g sample of something by 1oC

  32. Vaporisation • Energy has to be supplied to a liquid to enable it to overcome forces that hold molecules together • endothermic process (H positive) • Melting • Energy is supplied to a solid to enable it to vibrate more vigorously until molecules can move past each other and flow as a liquid • endothermic process (H positive) • Freezing • Liquid releases energy and allows molecules to settle into a lower energy state and form a solid • exothermic process (H negative) • (we remove heat from water when making ice in freezer)

  33. Endothermic reactions: Reactants + energy as heat products (H +ve) Reaction Enthalpies All chemical reactions either release or absorb heat Exothermic reactions: Reactants products + energy as heat (H -ve) e.g. burning fossil fuels e.g. photosynthesis

  34. Bond Strengths • Bond strengths measured by bond enthalpy HB (+ve values) • bond breaking requires energy (+ve H) • bond making releases energy (-ve H) Lattice Enthalpy • A measure of the attraction between ions (the enthalpy change when a solid is broken up into a gas of its ions) • all lattice enthalpies are positive • I.e. energy is required o break up solids

  35. Enthalpy of hydration Hhyd • the enthalpy change accompanying the hydration of gas-phase ions • Na+ (g) + Cl- (g) Na+ (aq) + Cl- (aq) • -ve H values (favourable interaction) WHY DO THINGS DISSOLVE? • If dissolves and solution heats up : exothermic • If dissolves and solution cools down: endothermic

  36. Enthalpy of Solution Lattice Enthalpy Enthalpy of Hydration = + Ions associating with water Breaking solid into ions + = Dissolving Substances dissolve because energy and matter tend to disperse (spread out in disorder) 2nd law of Thermodynamics

  37. USING ENTHALPY Making H2O from H2 involves two steps. H2(g) + 1/2 O2(g) ---> H2O(g) + 242 kJ H2O(g) ---> H2O(liq) + 44 kJ ----------------------------------------------------------------------- H2(g) + 1/2 O2(g) --> H2O(liq) + 286 kJ Example of HESS’S LAW— If a rxn. is the sum of 2 or more others, the net ∆H is the sum of the ∆H’s of the other rxns.

  38. Hess’s Law & Energy Level Diagrams Forming H2O can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed.

  39. Hess’s Law & Energy Level Diagrams Forming CO2 can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed.

  40. ∆H along one path = ∆H along another path • This equation is valid because ∆H is a STATE FUNCTION • These depend only on the state of the system and not on how the system got there. • V, T, P, energy — and your bank account! • Unlike V, T, and P, one cannot measure absolute H. Can only measure ∆H.

  41. Standard Enthalpy Values Most ∆H values are labeled ∆Ho Measured under standard conditions P = 1 bar = 105 Pa = 1 atm /1.01325 Concentration = 1 mol/L T = usually 25 oC with all species in standard states e.g., C = graphite and O2 = gas

  42. Enthalpy Values Depend on how the reaction is written and on phases of reactants and products H2(g) + 1/2 O2(g) --> H2O(g) ∆H˚ = -242 kJ 2 H2(g) + O2(g) --> 2 H2O(g) ∆H˚ = -484 kJ H2O(g) ---> H2(g) + 1/2 O2(g) ∆H˚ = +242 kJ H2(g) + 1/2 O2(g) --> H2O(liquid) ∆H˚ = -286 kJ

  43. Standard Enthalpy Values NIST (Nat’l Institute for Standards and Technology) gives values of ∆Hfo = standard molar enthalpy of formation — the enthalpy change when 1 mol of compound is formed from elements under standard conditions. See Table 6.2

  44. ∆Hfo, standard molar enthalpy of formation Enthalpy change when 1 mol of compound is formed from the corresponding elements under standard conditions H2(g) + 1/2 O2(g) --> H2O(g) ∆Hfo (H2O, g)= -241.8 kJ/mol By definition, ∆Hfo= 0 for elements in their standard states.

  45. (product is called “water gas”) Using Standard Enthalpy Values Use ∆H˚’s to calculate enthalpy change for H2O(g) + C(graphite) --> H2(g) + CO(g)

  46. Using Standard Enthalpy Values H2O(g) + C(graphite) --> H2(g) + CO(g) From reference books we find • H2(g) + 1/2 O2(g) --> H2O(g) ∆Hf˚ = - 242 kJ/mol • C(s) + 1/2 O2(g) --> CO(g) ∆Hf˚ = - 111 kJ/mol

  47. Using Standard Enthalpy Values H2O(g) --> H2(g) + 1/2 O2(g) ∆Ho = +242 kJ C(s) + 1/2 O2(g) --> CO(g)∆Ho = -111 kJ -------------------------------------------------------------------------------- H2O(g) + C(graphite) --> H2(g) + CO(g) ∆Honet = +131 kJ To convert 1 mol of water to 1 mol each of H2 and CO requires 131 kJ of energy. The “water gas” reaction is ENDOthermic.

  48. Using Standard Enthalpy Values In general, when ALL enthalpies of formation are known: Calculate ∆H of reaction? ∆Horxn =  ∆Hfo(products) -  ∆Hfo(reactants) Remember that ∆ always = final – initial

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