Thermochemistry

1. Chapter Six Thermochemistry

2. Energy • Energyis the capacity to do work (to displace or move matter). • Energy literally means “work within”; however, an object does not contain work. • Potential energy is energy of position or composition. • Kinetic energy is the energy of motion. Ek = ½ mv2 Energy has the units of joules (J or kg . m2/s2)

3. Potential Energy and Kinetic Energy At what point in each bounce is the potential energy of the ball at a maximum?

4. 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.

5. Types of Systems • Open: energy and matter can be exchanged with the surroundings. • Closed: energy can be exchanged with the surroundings, matter cannot. • Isolated: neither energy nor matter can be exchanged with the surroundings. After the lid of the jar is unscrewed, which kind of system is it? A closed system; energy (not matter) can be exchanged.

6. 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

7. 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?

8. 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.

9. Pressure-Volume Work For now we will consider only pressure-volume work. work (w) = –PDV How would the magnitude of DV compare to the original gas volume if the two weights (initial and final) were identical?

10. 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.

11. 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

12. First Law: Sign Convention • 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

13. Example 6.1 A gas does 135 J of work while expanding, and at the same time it absorbs 156 J of heat. What is the change in internal energy? Example 6.2: A Conceptual Example The internal energy of a fixed quantity of an ideal gas depends only on its temperature. If a sample of an ideal gas is allowed to expand against a constant pressure at a constant temperature, (a) what is DUfor the gas? (b) Does the gas do work? (c) Is any heat exchanged with the surroundings?

14. Heats of Reaction (qrxn) • 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.

15. 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

16. Internal Energy Change at Constant Volume • For a system where the reaction is carried out at constant volume, DV = 0 and DU = qV. • All the thermal energy produced by conversion from chemical energy is released as heat; no P-V work is done.