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

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

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

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

5. Energy: forms Kinetic energy is the energy of motionKinetic 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 nucleusPotential 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 (1/2mv 2 ), we get the units of energy:From the definition of kinetic energy (1/2mv 2 ), we get the units of energy: kg m 2 /s 2 S.I. unit for energy is the joule (J) = 1NmS.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 1ºCAnother 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 calNote the food calorie (Cal) = 1 000 cal

7. Interchange and conservation Energy in its many forms can be changed from one to anotherEnergy 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: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 motionThermal 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 bodyHeat 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 bondsChemical energy is the potential energy stored in chemical bonds

10. System and surroundings Any process can be divided into the SYSTEM contained within the SURROUNDINGSAny 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 workInternal 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 systemTypically 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 energyWe can measure the change to the internal energy ΔU = U final - U initial

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

13. Workin’ for a livin’ Mechanical work is force applied over a distanceMechanical work is force applied over a distance W = F x d In chemical process release of gas allows work to be done by systemIn 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 AGas 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: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, -PΔV 0, -PΔV < 0) –Work done by system is +ve in contraction (ΔV < 0) ΔU > 0 (ΔV 0)ΔU > 0 (ΔV 0)

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

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

17. Deposits and withdrawals Process is always viewed from perspective of systemProcess is always viewed from perspective of system Energy leaving system has negative signEnergy leaving system has negative sign –(decreases internal energy – lowers the chemical bank balance) Energy entering system has positive signEnergy entering system has positive sign –(increases internal energy – increases chemical bank balance) Useful process is one where change is negativeUseful process is one where change is negative Energy is in the form of heat or workEnergy 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 = U final – U initial ) –Energy that flows from system to surroundings has negative sign (U final < U initial,) –Energy that flows into system from surroundings has positive sign (U final > U initial.)

19. Functions of state State FunctionState 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 temperatureInternal 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 pathwayChange in state function between two states is independent of pathway Given two states of a system: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 termsAny chemical process may have associated with it heat and work terms The total internal energy change will be the sum of the contributions from eachThe 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 = ΔUIn 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 ≠ 0Most reactions are conducted in open vessels where P is constant and ΔV ≠ 0 The heat change at constant pressure isThe heat change at constant pressure is q P = ΔU + P ΔV Enthalpy (H) is defined as:Enthalpy (H) is defined as: H = U + PV

23. Heats of reaction and enthalpy Absolute enthalpy of system is not knownAbsolute enthalpy of system is not known Enthalpy change is measuredEnthalpy change is measured Enthalpy change is known as heat of reactionEnthalpy change is known as heat of reaction ΔH = q P = ΔU + P ΔV –If reaction is exothermic and involves expansion: ΔU 0 ΔH less negative than ΔU ΔU 0 ΔH less negative than ΔU Enthalpy change is portion of internal energy available as heat after work is doneEnthalpy change is portion of internal energy available as heat after work is done If no work done, all the internal energy change is enthalpyIf 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?In reactions involving volume change at constant P, ΔH and ΔU are different. How big is it? Consider reaction:Consider reaction: 1 additional mole of gas is produced1 additional mole of gas is produced ΔU = - 2045 kJ, ΔH = - 2043 kJ PΔV = + 2kJ