PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY - PowerPoint PPT Presentation

slide1 n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY PowerPoint Presentation
Download Presentation
PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY

play fullscreen
1 / 75
PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY
132 Views
Download Presentation
corina
Download Presentation

PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. GENERAL CHEMISTRY Principles and Modern Applications TENTH EDITION PETRUCCI HERRING MADURA BISSONNETTE 7 Thermochemistry PHILIP DUTTON UNIVERSITY OF WINDSOR DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY

  2. Thermochemistry

  3. Thermochemistry: Basic Terms • Thermochemistry is the study of energy changes that occur during chemical reactions. • System: the part of the universe being studied. • Surroundings: the rest of the universe.

  4. 6-1 Getting Started: Some Terminology

  5. Energy, U The capacity to do work (to displace or move matter). Work, w Force acting through a distance. Kinetic Energy, eK (Ek = ½ mv2) The energy of motion. Potential Energy, V The stored energy has potential to do work. Energy has the units of joules (J or kg . m2/s2)

  6. Kinetic Energy 2 1 m kg = J mv2 ek= [ek ] = 2 s • Work w = F x d m m = J [w ] = = mxaxd kg s2

  7. Potential Energy Energy due to condition, position, or composition. Associated with forces of attraction or repulsion between objects. Energy can change from potential to kinetic.

  8. FIGURE 7-2 Potential energy (P.E.) and kinetic energy (K.E.)

  9. 7-2 Heat • Thermal Energy • Kinetic energy associated with random molecular motion. • In general proportional to temperature. • An intensive property. • Heat and Work • q and w. • Energy changes.

  10. Internal Energy (U) • Internal energy (U) is the total energy contained within a system • Part of U is kinetic energy (from molecular motion) • Translational motion, rotational motion, vibrational motion. • Collectively, these are sometimes called thermal energy • Part of U is potential energy • Intermolecular and intramolecular forces of attraction, locations of atoms and of bonds. • Collectively these are sometimes called chemical energy

  11. Heat (q) • Technically speaking, heat is not“energy.” • Heat is energy transfer between a system and its surroundings, caused by a temperature difference. More energetic molecules … … transfer energy to less energetic molecules. • Thermal equilibrium occurs when the system and surroundings reach the same temperature and heat transfer stops. How do the root-mean-square speeds of the Ar atoms and the N2 molecules compare at the point of thermal equilibrium?

  12. Work (w) • Like heat, work is an energy transfer between a system and its surroundings. • Unlike heat, work is caused by a force moving through a distance (heat is caused by a temperature difference). • A negative quantity of work signifies that the system loses energy. • A positive quantity of work signifies that the system gains energy. • There is no such thing as “negative energy” nor “positive energy”; the sign of work (or heat) signifies the direction of energy flow.

  13. Internal energy is transferred between a system and its surroundings as a result of a temperature difference. • Heat “flows” from hotter to colder. • Temperature may change. • Phase may change (an isothermal process).

  14. Calorie (cal) • The quantity of heat required to change the temperature of one gram of water by one degree Celsius. • Joule (J) • SI unit for heat 1 cal = 4.184 J

  15. FIGURE 7-3 Determining the specific heat of lead – Example 7-2 illustrated

  16. q = q = mcT CT Heat Capacity • The quantity of heat required to change the temperature of a system by one degree. • Molar heat capacity. • System is one mole of substance. • Specific heat capacity, c. • System is one gram of substance • Heat capacity, C • (Mass of system) x specific heat. • The heat capacity (C) of a system is the quantity of heat required to change the temperature of the system by 1 OC. • C = q/DT (units are J/oC) • Molar heat capacity is the heat capacity of one mole of a substance. • The specific heat (c) is the heat capacity of one gram of a pure substance (or homogeneous mixture). • c = C/m = q/(mDT)

  17. Law of conservation of energy • In interactions between a system and its surroundings the total energy remains constant— energy is neither created nor destroyed. qsystem + qsurroundings = 0 qsystem = -qsurroundings

  18. TABLE 7.1 Some Specific Heat Values, J g-1 °C-1 Slide 20 of 57

  19. 7-3 Heats of Reaction and Calorimetry • Chemical energy. • Contributes to the internal energy of a system. • Heat of reaction, qrxn. • The quantity of heat exchanged between a system and its surroundings when a chemical reaction occurs within the system, at constant temperature. • qrxnis the quantity of heat exchanged between a reaction system and its surroundings. • An exothermic reaction gives off heat • In an isolated system, the temperature increases. • The system goes from higher to lower energy; qrxn is negative. • An endothermicreaction absorbs heat • In an isolated system, the temperature decreases. • The system goes from lower to higher energy; qrxn is positive.

  20. CaO(s) + H2O(l) -> Ca(OH)2(s) Ba(OH)2 8H2O(s) + 2 NH4Cl(s) -> BaCl2 2H2O(s) + 2NH3(aq) + 8H2O(l) Exothermic and endothermic reactions

  21. Conceptualizing an Exothermic Reaction Surroundings are at 25 °C Typical situation: some heat is released to the surroundings, some heat is absorbed by the solution. 35.4 °C 32.2 °C 25 °C In an isolated system, all heat is absorbed by the solution. Maximum temperature rise. Hypothetical situation: all heat is instantly released to the surroundings. Heat = qrxn

  22. FIGURE 7-4 Conceptualizing a heat of reaction at constant temperature

  23. Define the heat capacity of the calorimeter: qcal = miciT = CcalT heat all i Bomb Calorimetry • Some reactions, such as combustion, cannot be carried out in a coffee-cup calorimeter. • In a bomb calorimeter, a sample of known mass is placed in a heavy-walled “bomb,” which is then pressurized with oxygen. • Since the reaction is carried out at constant volume, –qrxn = qcalorimeter = DU … but in many cases the value of DU is a good approximation of DH. qrxn = -qcal qcal = q bomb + q water + q wires +… FIGURE 7-5 • A bomb calorimeter assembly

  24. Define the heat capacity of the calorimeter: qcal = miciT = CcalT all i The “Coffee-Cup” Calorimeter • A simple calorimeter. • Well insulated and therefore isolated. • Measure temperature change. qrxn = -qcal FIGURE 7-6 • A Styrofoam “coffee-cup” calorimeter

  25. 7-4 Work • In addition to heat effects chemical reactions may also do work. • Gas formed pushes against the atmosphere. • The volume changes. • Pressure-volume work. FIGURE 7-7 • Illustrating work (expansion) during the chemical reaction • 2 KClO3(s) 2 KCl(s) + 3 O2(g)

  26. w = F x d = (m x g) xh (m x g) x A h = A = PV w = -PextV • A negative quantity of work signifies that the system loses energy. • A positive quantity of work signifies that the system gains energy. FIGURE 7-8 Pressure-volume work

  27. First Law of Thermodynamics “Energy cannot be created or destroyed.” Inference: the internal energy change of a system is simply the difference between its final and initial states: DU = Ufinal – Uinitial Additional inference: if energy change occurs only as heat (q) and/or work (w), then: DU = q + w

  28. 7-5 The First Law of Thermodynamics • Internal Energy, U. • Total energy (potential and kinetic) in a system. • Translational kinetic energy. • Molecular rotation. • Bond vibration. • Intermolecular attractions. • Chemical bonds. • Electrons. FIGURE 7-9 • Some contributions to the internal energy of a system

  29. The First Law of Thermodynamics • A system contains only internal energy. • A system does not contain heat or work. • These only occur during a changein the system. • Law of Conservation of Energy • The energy of an isolated system is constant U = q + w

  30. The First Law of Thermodynamics • An isolated system is unable to exchange either heat or work with its surroundings, so that DUisolated system = 0, and we can say: The energy of an isolated system is constant.

  31. Energy entering a system carries a positive sign: • heat absorbed by the system, or • work done on the system • Energy leaving a system carries a negative sign • heat given off by the system • work done by the system

  32. State Functions • Thestate of a system: its exact condition at a fixed instant. • State is determined by the kinds and amounts of matter present, the structure of this matter at the molecular level, and the prevailing pressure and temperature. • A state function is a property that has a unique value that depends only the present state of a system, and does not depend on how the state was reached (does not depend on the history of the system). • Law of Conservation of Energy – in a physical or chemical change, energy can be exchanged between a system and its surroundings, but no energy can be created or destroyed.

  33. Functions of State • Any property that has a unique value for a specified state of a system is said to be a function of state or a state function. • Water at 293.15 K and 1.00 atm is in a specified state. • d = 0.99820 g/mL • This density is a unique function of the state. • It does not matter how the state was established.

  34. Functions of State • U is a function of state. • Not easily measured. • U has a unique value between two states. • Is easily measured.

  35. Path Dependent Functions • Changes in heat and work are not functions of state. • Remember example 7-5, w = -1.24 x 102 J in a one step expansion of gas: • let us consider a two step process from 2.40 atm to 1.80 atm, and finally to 1.20 atm. w = (-1.80 atm)(1.36-1.02)L + (-1.20 atm)(2.04-1.36)L = (-0.61) L atm +(- 0.82) L atm = (-1.43) L atm = -1.44 x 102 J Compared -1.24  102 J for the one stage process

  36. 7-6 Heats of Reaction: U and H Reactants → Products UiUf U = Uf - Ui U = qrxn + w In a system at constant volume (bomb calorimeter): U = qrxn + 0 = qrxn = qv But we live in a constant pressure world! How does qp relate to qv?

  37. FIGURE 7-13 Two different paths leading to the same internal energy change in a system

  38. Heats of Reaction qV = qP + w We know that w = - PV and U = qv, therefore: U = qP - PV qP = U + PV These are all state functions, so define a new function. Let enthalpybe H = U + PV Then H = Hf – Hi = U + PV If we work at constant pressure and temperature: H = U + PV=qP

  39. Comparing Heats of Reaction qV = U = H - PV = -563.5 kJ/mol w = PV = P(Vf – Vi) = RT(nf – ni) = -2.5 kJ qP = H = -566 kJ/mol FIGURE 7-14 • Comparing heats of reaction at constant volume and constant pressure for the reaction 2 CO(g) + 2 O2(g) 2 CO2(g)

  40. Enthalpy Change (DH) Accompanying a Change in State of Matter Molar enthalpy of vaporization: H2O (l) → H2O(g) H = 44.0 kJ at 298 K Molar enthalpy of fusion: H2O (s) → H2O(l) H = 6.01 kJ at 273.15 K

  41. Properties of Enthalpy Two logs on a fire give off twice as much heat as does one log. • Since U, P, and V are all state functions, enthalpy H must be astate functionalso. • Enthalpy changeshave unique values. DH = qP Enthalpy change depends only on the initial and final states. In a chemical reaction we call the initial state the ____ and the final state the ____. Enthalpy is an extensive property. It depends on how much of the substance is present.

  42. Enthalpy Diagrams Values of DH are measured experimentally. Negative values indicate exothermic reactions. Positive values indicate endothermic reactions.