Introduction to thermochemistry

1. Introduction to thermochemistry Heat, work, energy and the First Law

2. Learning objectives • Define energy and identify types of energy • Compare and contrast heat and work • Describe internal energy and how it changes during a process • Describe basic properties of state functions • Apply first law of thermodynamics to determine heat flow and work • Define enthalpy

3. Behind it all • Why do things happen in chemistry? • Substances spontaneously move towards a position of greater stability – in energy terms • A high energy state is unstable with respect to a state of lower energy • A simple (but incomplete) analogy is a ball rolling downhill

4. Energy • Is capacity to perform work • Mechanical work is application of force over distance • Heat is energy transferred by virtue of temperature gradient – associated with molecular motion • Joule demonstrated experimentally that heat and work are interchangeable forms of energy

5. Energy: forms • Kinetic energy is the energy of motion • Potential energy is energy stored – by position, within a spring, within a chemical bond, within the particles of a nucleus

6. Energy: units • From the definition of kinetic energy (1/2mv2), we get the units of energy: kg m2/s2 • S.I. unit for energy is the joule (J) = 1Nm • Another common unit is the calorie (cal): the energy required to raise the temperature of 1 g of water by 1ºC 1 cal = 4.184 J • Note the food calorie (Cal) = 1 000 cal

7. Interchange and conservation • Energy in its many forms can be changed from one to another • A stationary ball on a hill has potential energy (P.E.) by virtue of position but no kinetic energy (K.E.). As it rolls down, it gains K.E. at the expense of P.E.

8. Energy conservation • There is no gain or loss: Energy cannot be created or destroyed; it can only be changed from one form to another • Chemical processes involve conversion of chemical potential energy into other forms and vice versa • Energy never goes away, but in some forms it is more useful than others • Efficient energy use means maximizing the useful part and minimizing the useless part

9. Some like it hot • Thermal energy is the kinetic energy of molecular motion • Temperature measures the magnitude of the thermal energy • Heat is the transfer of thermal energy from a hotter to a cooler body • Temperature gradient provides the “pressure” for heat to flow • Chemical energy is the potential energy stored in chemical bonds

10. System and surroundings • Any process can be divided into the SYSTEM contained within the SURROUNDINGS • When energy changes are measured in a chemical reaction, the system is the reaction mixture and the surroundings are the flask, the room, and the rest of the universe.

11. Internal energy • Internal energy is the sum of all of the types of energy (kinetic and potential) of the system. It is the capacity of the system to do work • Typically we don’t know the absolute value of U for the system • (Internal energy usually has symbol U. Other sources use E) • We can measure the change to the internal energy ΔU = Ufinal - Uinitial

12. Work and internal energy • Work done on system increases its internal energy • Work done by system decreases its internal energy ΔU = w

13. Workin’ for a livin’ • Mechanical work is force applied over a distance W = F x d • In chemical process release of gas allows work to be done by system

14. Work done at constant pressure • Gas generated in reaction pushes against the piston with force: P x A • At constant P, volume increases by ΔV and work done by system is: w = -PΔV (ΔV = A x d) • Work done by system is –ve in expansion (ΔV > 0) • ΔU < 0 (ΔV > 0, -PΔV < 0) • Work done by system is +ve in contraction (ΔV < 0) • ΔU > 0 (ΔV < 0, -PΔV > 0)

15. Expansion work • Work done by gas expanding: w = -PexΔV • In expansion the ΔV > 0; w < 0 ΔU < 0 • In contraction, ΔV < 0; w > 0 ΔU > 0

16. Heat and internal energy • Heat is transfer of energy by virtue of temperature gradient ΔU = q • If system is cooler than surroundings q > 0 • If system is hotter than surroundings q < 0

17. Deposits and withdrawals • Process is always viewed from perspective of system • Energy leaving system has negative sign • (decreases internal energy – lowers the chemical bank balance) • Energy entering system has positive sign • (increases internal energy – increases chemical bank balance) • Useful process is one where change is negative • Energy is in the form of heat or work • ΔU = q + w

18. First Law of Thermodynamics Total internal energy of isolated system is constant • Energy change is difference between final and initial states (ΔU = Ufinal – Uinitial) • Energy that flows from system to surroundings has negative sign (Ufinal < Uinitial,) • Energy that flows into system from surroundings has positive sign (Ufinal > Uinitial.)

19. Functions of state • State Function A property that depends only on present state of the system and is independent of pathway to that state • Internal energy is a state function, as are pressure, volume and temperature

20. Significance of state functions • Change in state function between two states is independent of pathway • Given two states of a system: • ΔU is always the same • q and w depend on type of change

21. Heat and work • Any chemical process may have associated with it heat and work terms • The total internal energy change will be the sum of the contributions from each ΔU = q + w = q - P ΔV q = ΔU + P ΔV • In a sealed system ΔV = 0, so q = ΔU

22. Cracked pots and enthalpy • Most reactions are conducted in open vessels where P is constant and ΔV ≠ 0 • The heat change at constant pressure is qP = ΔU + P ΔV • Enthalpy (H) is defined as: H = U + PV

23. Heats of reaction and enthalpy • Absolute enthalpy of system is not known • Enthalpy change is measured • Enthalpy change is known as heat of reaction ΔH = qP = ΔU + P ΔV • If reaction is exothermic and involves expansion: • ΔU < 0, ΔV > 0 ΔH less negative than ΔU • Enthalpy change is portion of internal energy available as heat after work is done • If no work done, all the internal energy change is enthalpy

24. Comparing ΔH and ΔU • In reactions involving volume change at constant P, ΔH and ΔU are different. How big is it? • Consider reaction: • 1 additional mole of gas is produced ΔU = - 2045 kJ, ΔH = - 2043 kJ PΔV = + 2kJ