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AP Chemistry

AP Chemistry

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AP Chemistry

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  1. AP Chemistry Thermodynamics

  2. Bell work • “Imagine that your parents are forcing you to invite someone to go to a party that you are hosting. You don’t actually like the person, but your parents are insistent, so you have no choice but to comply. You give the acquaintance directions to the party, but you intentionally make the directions vague and difficult to use, filled with bland phrases like, ‘Turn left at the light and then turn right a couple of miles down from there.’ Your hope is that you’ve provided the required information, but done so in a manner that will not help much.” • How does this relate to the AP chemistry exam?

  3. A word about course structure… • Things are always subject to change, but for now… • This lecture is very, very detailed again, but after this, I’m going to pick out the parts I think are most important and you can read the other stuff on your own in case the AP test writers thinks that’s important. This will help with pace and to avoid getting bogged down with crazy details

  4. Homework – due next Tuesday • 5.2, 5.3, 5.10 a and b, 5.12, 5.24, 5.25, 5.26, 5.33, 5.35, 5.36, 5.37, 5.43, 5.44, 5.45, 5.46

  5. Outline • Chapter 5 - Thermochemistry • Focus: 5.1, 5.2, 5.6, 5.7 • 5.1 The nature of energy • 5.2 The first law of thermodynamics • 5.6 Hess’s law • 5.7 Enthalpies of formation

  6. Main topicsThermochemistry • Types of energy • Defining system and surroundings • First law of thermodynamics • Internal energy • Enthalpy • Calorimetry • Hess’s Law • Enthalpy changes

  7. Thermodynamics • Thermodynamics: the study of energy and its transformations • Chapter 5 focuses on thermochemistry: relationships between chemical reactions and energy changes involving heat

  8. Energy • The ability to do work or transfer heat. • Work: Energy used to cause an object that has mass to move. • Heat: Energy used to cause the temperature of an object to rise.

  9. Different types of energy • Kinetic energy • Potential energy • Electrostatic potential energy • Chemical energy • Thermal energy

  10. How can matter possess energy and how can that energy be transferred? • Kinetic energy – the energy of motion • Magnitude of kinetic energy (Ekor KE) depends on mass and speed • KE= ½ mv2 • Potential energy • Energy an object possesses by virtue of its position or chemical composition. • PE = mgh • m=mass; h=height; g=gravitational constant (9.8m/s2)

  11. Where is the bicyclist’s potential energy the highest? What happens to his potential energy as he goes down the hill?

  12. Electrostatic potential energy • While gravity is an important force for large objects, chemistry deals with extremely small objects, such as molecules and atoms • Gravitational forces play a negligible role in submicroscopic interactions • Electrostatic potential energy is more important in these interactions • Electrostatic potential energy arises from the interactions between charged particles

  13. Electrostatic potential energy Q1Q2 ĸ • Eel  electrostatic potential energy • ĸ  proportionality constant (8.99 x 109 J-m/C2 ) • Q1and Q2  electrical charges of the two interacting objects Eel= d

  14. Electrostatic potential energy • When Q1and Q2 have the same charge, they repel and Eelis positive • When Q1and Q2 have opposite charges, they attract and Eelis negative • The lower the energy of a system, the more stable it is • Thus, the more strongly opposite charges interact, the more stable the system

  15. Units of energy • SI unit: Joules • A joule is equal to what? • Use KE= ½ mv2 • mkg, vm/s, v2 = m2/s2 • 1 J= kg-m2/s2 • We will often use kJ when discussing energies associated with chemical reactions

  16. Units of energy • Traditionally, calories have been used as the unit (non-SI unit) • calorie: amount of energy required to raise the temperature of 1 g of water from 14.5 ° C to 15.5 ° C • 1 cal = 4.184 J (exactly) • Nutritional Calorie: 1 Cal = 1000 cal = 1 kcal • Nutritional Calorie is capitalized

  17. Systems and surroundings • System: the portion of the universe we single out for study • Surroundings: everything else • Closed system: can exchange energy (in the form of heat and work) but not matter with its surroundings

  18. System and surroundings • Here we have hydrogen and oxygen gases confined in a cylinder with a movable piston. • What is the system? What are the surroundings? • Has the system lost of gained mass? Why or why not? 2 H2 (g) + O2  2 H2O (g) + energy The system includes the molecules we want to study (here, the hydrogen and oxygen molecules). The surroundings are everything else (here, the cylinder and piston).

  19. Transfer of energy • Energy is transferred between systems and surroundings in two general ways: as work or heat • Work: • Heat: • Energy used to cause an object that has mass to move. • Energy transferred when an object is moved by a force • Energy used to cause the temperature of an object to rise.

  20. Work • w = F x d • Work equals the product of the force, F, and the distance, d, that the object is moved • We perform work when we lift an object against the force of gravity or when we bring two like charges together • In these two situations, what is the system and what are the surroundings? • What is the transfer of energy?

  21. Heat • Heat: energy transferred from a hotter object to a colder one • Energy transferred between a system and its surroundings as a result of the difference in temperature • Example: Burning of natural gas combustion reaction • What is the system and what are the surroundings? • What is the transfer of energy?

  22. Example • The potential energy of this ball of clay is increased when it is moved from the ground to the top of the wall. • As the ball falls, its potential energy is converted to kinetic energy. • When it hits the ground, its kinetic energy falls to zero (since it is no longer moving); some of the energy does work on the ball, the rest is dissipated as heat.

  23. Example 2 • A bowler lifts a 5.4-kg (12-lb) bowling ball from ground level to a height of 1.6 m (5.2 feet) and then drops the ball back to the ground. • What happens to the potential energy of the bowling ball as it is raised from the ground? • What quantity of work, in J, is used to raise the ball? • After the ball is dropped, it gains kinetic energy. If we assume that all of the work done in part (b) has been converted to kinetic energy by the time the ball strikes the ground, what is the speed of the ball at the instant just before it hits the ground? (Note: The force due to gravity is F = mg, where m is the mass of the object and g is the gravitational constant; g = 9.8 m/s2.)

  24. What happens to the potential energy of the bowling ball as it is raised from the ground? • Because the bowling ball is raised to a greater height above the ground, its potential energy increases.

  25. What quantity of work, in J, is used to raise the ball? The ball has a mass of 5.4 kg, and it is lifted a distance of 1.6 m. To calculate the work performed to raise the ball, we use both Equation 5.3 and F = mg for the force that is due to gravity: w = F x d w = m x g x d w = (5.4 kg) (9.8 m/s2) (1.6 m) w = 85 kg-m/s2 w = 85 J

  26. After the ball is dropped, it gains kinetic energy. If we assume that all of the work done in part (b) has been converted to kinetic energy by the time the ball strikes the ground, what is the speed of the ball at the instant just before it hits the ground? (Note: The force due to gravity is F = mg, where m is the mass of the object and g is the gravitational constant; g = 9.8 m/s2.) • When the ball is dropped, its potential energy is converted to kinetic energy. At the instant just before the ball hits the ground, we assume that the kinetic energy is equal to the work done in part (b), 85 J: • How can we find the speed?

  27. We can now solve this equation for v: Check: Work must be done in part (b) to increase the potential energy of the ball, which is in accord with our experiences. The units are appropriate in both parts (b) and (c). The work is in units of J and the speed in units of m/s. In part (c) we have carried an additional digit in the intermediate calculation involving the square root, but we report the final value to only two significant figures, as appropriate.

  28. First law of thermodynamics • Energy cannot be created or destroyed • It is also the law of conservation of energy: simply put, energy is conserved • The total energy of the universe is constant • Any energy lost by the system must be gained by the surroundings and vice versa • This can be expressed quantitatively

  29. Internal Energy • Internal energy: sum of all the kinetic and potential energies of all its components • Denoted as U or E (we’ll use E) • It is very difficult to know precisely the internal energy of a system (includes the motion of molecules through space, their rotations, and internal vibrations) • Rather, we hope to know the change in E • By definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system: E = Efinal−Einitial

  30. Values of ΔE • Negative Δ E: Efinal < Einitial • System has lost energy to surroundings • Positive Δ E: Efinal > Einitial • System has gained energy from surroundings Negative Δ E Positive Δ E

  31. Changes in Internal Energy Energy diagram • If E > 0, Efinal > Einitial • Therefore, the system absorbed energy from the surroundings. • This energy change is called endergonic.

  32. Changes in Internal Energy • If E < 0, Efinal < Einitial • Therefore, the system released energy to the surroundings. • This energy change is called exergonic.

  33. Relating ΔE to Heat and Work • When a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or liberated from the system plus the work done on or by the system • Any energy entering the system has a positive sign • Therefore, q or w would have a positive value Δ E = q + w heat work

  34. Example • Two gases, A(g) and B(g), are confined in a cylinder-and-piston arrangement. Substances A and B react to form a solid product: As the reactions occurs, the system loses 1150 J of heat to the surrounding. The piston moves downward as the gases react to form a solid. As the volume of the gas decreases under the constant pressure of the atmosphere, the surroundings do 480 J of work on the system. What is the change in the internal energy of the system? A (g) + B (g)  C (s)

  35. We first determine the signs of q and w (Table) and then use Equation 5.5, E = q + w, to calculate E. • Heat is transferred from the system to the surroundings, and work is done on the system by the surroundings, so q is negative and w is positive: q –1150 J and w  480 kJ. Thus, E is • Did the system lose or gain energy?

  36. Practice • Calculate the change in the internal energy of the system for a process in which the system absorbs 140 J of heat from the surroundings and does 85 J of work on the surroundings.

  37. Endothermic vs Exothermic • When heat is released by the system to the surroundings, the process is exothermic. • When heat is absorbed by the system from the surroundings, the process is endothermic.

  38. State functions • The internal energy of a system is independent of the path by which the system achieved that state • Because of this, internal energy is a state function • It depends only on the present state of the system, not on the path by which the system arrived at that state. • And so, E depends only on Einitial and Efinal.

  39. Suppose you are traveling between Chicago and Denver. Chicago is 596 ft above sea level; Denver is 5280 ft above sea level. No matter what route you take, the altitude will change 4684 feet. The distance you travel, however, will depend on your route. Altitude is analogous to a state function because the change in altitude is independent of the path taken. Distance traveled is not a state function. • In what ways is the balance in your checkbook a state function?

  40. A battery in a flashlight can be discharged by producing heat and light. The same battery in a toy car is used to produce heat and work. • The change in internal energy of the battery is the same (ΔE – a state function) • But q and w are different in the two cases (q and w are not state functions)

  41. Bell work • What is the first law of thermodynamics? • How can we quantify the change in internal energy of a system (ie: what equation do we use)?

  42. Agenda • Bell work • Enthalpy • Endothermic and exothermic investigation • Time to work on homework • For tonight! • Look over Hess’s law (in Chapter 5)

  43. Work • When a process occurs in an open container, commonly the only work done is a change in volume of a gas pushing on the surroundings (or being pushed on by the surroundings). • We can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston. This is called pressure-volume work. w = −PV

  44. Enthalpy • If a process takes place at constant pressure (as the majority of processes we study do) and the only work done is this pressure-volume work, we can account for heat flow during the process by measuring the enthalpy of the system. • Enthalpy is the internal energy plus the product of pressure and volume: H = E + PV

  45. Enthalpy • When the system changes at constant pressure, the change in enthalpy, H, is H = (E + PV) • This can be written H = E + PV • Since E = q + w and w = −PV, we can substitute these into the enthalpy expression: H = E + PV H = (q+w) −w H = q • So, at constant pressure the change in enthalpy is the heat gained or lost.

  46. Enthalpy • A process is endothermic, then, when H is positive. • A process is exothermic when H is negative.

  47. Practice • Indicate the sign of the enthalpy change, H, in each of the following processes carried out under atmospheric pressure, and indicate whether the process is endothermic or exothermic: • An ice cube melts • 1 g of butane (C4H10)is combusted in sufficient oxygen to give complete combustion to CO2 and H2O.

  48. We must predict whether heat is absorbed or released by the system in each process. Processes in which heat is absorbed are endothermic and have a positive sign for H; those in which heat is evolved are exothermic and have a negative sign for H • 1.) The water that makes up the ice cube is the system. The ice cube absorbs heat from the surroundings as it melts, so H is positive and the process is endothermic. • 2.) The system is the 1 g of butane and the oxygen required to combust it. The combustion of butane in oxygen gives off heat, so H is negative and the process is exothermic.