Prentice Hall © 2003Chapter 5 Chapter 5 Thermochemistry

Prentice Hall © 2003Chapter 5 Kinetic Energy and Potential Energy Kinetic energy is the energy of motion Mass and energy are directly proportional Potential energy is the energy an object possesses by virtue of its position Potential energy can be converted into kinetic energy 5.1 The Nature of Energy

Prentice Hall © 2003Chapter 5 Potential energy is energy an object possesses by virtue of its position or chemical composition. –The most important form of potential energy in molecules is electrostatic potential energy, E el :

Prentice Hall © 2003Chapter 5 Units of Energy SI Unit for energy is the joule, J: We sometimes use the calorie instead of the joule: 1 cal = J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal 1J = 1kg-m/s 2

Prentice Hall © 2003Chapter 5 Systems and Surroundings System: part of the universe we are interested in Surroundings: the rest of the universe Example: “System” = reactants “Surroundings” = the beaker they are in and everything beyond it

Prentice Hall © 2003Chapter 5 Transferring Energy: Work and Heat Energy is transferred in two ways: 1.By causing the motion of an object against a force Force is a push or pull on an object Work is the product of force applied to an object over a distance: Energy is the work done to move an object against a force

Prentice Hall © 2003Chapter 5 2. By causing a temperature change Heat is the transfer of energy between two objects Energy is the capacity to do work or transfer heat

Prentice Hall © 2003Chapter 5 Internal Energy Internal Energy: total energy of a system Cannot measure absolute internal energy 5.2 The First Law of Thermodynamics

Prentice Hall © 2003Chapter 5 Relating  E to Heat and Work Energy cannot be created or destroyed Energy of (system + surroundings) is constant

Prentice Hall © 2003Chapter 5

Prentice Hall © 2003Chapter 5 From the first law of thermodynamics: –When a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system (q) plus the work done on or by the system (w): Change in internal energy

Text, P. 173 Remember: q = the heat added to or absorbed by the system w = the work done on or by the system

Text, P. 172

Prentice Hall © 2003Chapter 5 Exothermic and Endothermic Processes Endothermic: absorbs heat from the surroundings Exothermic: transfers heat to the surroundings An endothermic reaction feels cold An exothermic reaction feels hot

Prentice Hall © 2003Chapter 5 State Functions State function: depends only on the initial and final states of system, not on the path by which it arrived at that state 50g H 2 O (l) 25 o C

State Functions: Discharging a battery Text, P. 175 Battery shorted by coil: no work (nothing is moved against a force) All energy lost as heat q < 0 w = 0 Battery turns fan: work! q < 0 (small); w < 0 ΔE < 0 ΔE : state function! Remember: q = the heat added to or absorbed by the system w = the work done on or by the system

Prentice Hall © 2003Chapter 5 Sample Problems, P. 204 # 25, 37

Prentice Hall © 2003Chapter 5 Chemical reactions can absorb or release heat However, they also have the ability to do work Enthalpy, H: Heat transferred between the system & surroundings - constant pressure - no forms of work are performed (other than work from pressure-volume changes)… remember this for a future equation 5.3 Enthalpy

Prentice Hall © 2003Chapter 5 Enthalpy is a state function If the process occurs at constant pressure,

Prentice Hall © 2003Chapter 5 +  H means the system gains heat from the surroundings -  H means the surroundings gain heat from the system at constant pressure Recall, ΔE=q + w Work is done on surroundings (piston moves) We then can write: Text, P. 176 Piston moves to maintain constant pressure

Prentice Hall © 2003Chapter 5

Prentice Hall © 2003Chapter 5 The change in enthalpy equals the heat gained or lost at constant pressure H is more useful than E (heat is easily measurable) For most reactions, the difference in ΔH and ΔE is small because PΔV is small In summary…

Prentice Hall © 2003Chapter 5 Enthalpy is an extensive property (magnitude  H is directly proportional to amount): CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2CH 4 (g) + 4O 2 (g)  2CO 2 (g) + 4H 2 O(g)  H =  1604 kJ 5.4 Enthalpies of Reaction exothermic greatly exothermic “Thermochemical equations” For a reaction: All example equations can be found in the text, P. 180

Prentice Hall © 2003Chapter 5 When we reverse a reaction, we change the sign of  H: CO 2(g) + 2H 2 O (l)  CH 4(g) + 2O 2(g)  H = +890 kJ 2O 2(g) + CH 4(g)  CO 2(g) + H 2 O (l) ΔH = -890 kJ Change in enthalpy depends on state: Previously: 2O 2(g) + CH 4(g)  CO 2(g) + H 2 O (l) ΔH = -890 kJ If: 2O 2(g) + CH 4(g)  CO 2(g) + H 2 O (g) then ΔH = -802 kJ Because: H 2 O (l)  H 2 O (g) ΔH = +88 kJ Why: Less heat is available for transfer to the surroundings because enthalpy of H 2 O (g) is greater than that of H 2 O (l) (liquid  gas)

Prentice Hall © 2003Chapter 5 Enthalpy as a Guide A spontaneous process is one that is thermodynamically favored to happen –A reaction with a large, negative ΔH will be spontaneous –A reaction with a small, positive ΔH will be spontaneous in the reverse direction

Prentice Hall © 2003Chapter 5 Sample problems P # 39, 43, 45, 47

Prentice Hall © 2003Chapter 5 Heat Capacity and Specific Heat Calorimetry = measurement of heat flow Calorimeter = apparatus that measures heat flow Heat capacity = the amount of energy required to raise the temperature of an object (by one degree) Molar heat capacity = heat capacity of 1 mol of a substance Specific heat = specific heat capacity = heat capacity of 1 g of a substance 5.5 Calorimetry

Prentice Hall © 2003Chapter 5 Constant Pressure Calorimetry Atmospheric pressure is constant!

Prentice Hall © 2003Chapter 5 Read “Chemistry and Life” on P. 186 Sample Problems, P. 206 # 53 and 55

Prentice Hall © 2003Chapter 5 Hess’s law: if a reaction is carried out in a number of steps,  H for the overall reaction is the sum of  H for each individual step. For example: CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(g)  H = -802 kJ 2H 2 O(g)  2H 2 O(l)  H = -88 kJ CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l)  H = -890 kJ 5.6 Hess’s Law Combustion reaction Condensation of water Overall reaction Simplify similar terms in reactants and products!

Prentice Hall © 2003Chapter 5 Note that:  H 1 =  H 2 +  H 3 The quantity of heat generated is independent of the # of steps needed for the reaction to occur

Prentice Hall © 2003Chapter 5 Sample Problems: P. 207 # 63 and 65

Prentice Hall © 2003Chapter 5 Enthalpy of formation,  H f : the enthalpy change that occurs if 1 mol of compound is formed from its constituent elements Standard conditions:1 atm and 25 o C (298 K) Standard enthalpy,  H o, is the enthalpy measured when everything is in its standard state (then we’re able to compare) Standard enthalpy of formation,  H o f : 1 mol of compound is formed from substances in their standard states 5.7 Enthalpies of Formation

Prentice Hall © 2003Chapter 5 If there is more than one state for a substance under standard conditions, the more stable one is used  H o f of the most stable form of an element is zero Refer to chapter 1, periodic table info Diatomics Appendix C, P. 1041

Prentice Hall © 2003Chapter 5 See Appendix C, P. 1100…standard states are included!

Prentice Hall © 2003Chapter 5 Calculating Enthalpies of Reaction We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation: For a reaction Where n and m are stoichiometric coefficients Practice: P # 71, 73, 75, 77

Prentice Hall © 2003Chapter Foods and Fuels Fuel value = energy released when 1 g of substance is burned –Chart, P. 181: Fuel Values (will be given these on the test) 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal 1 kcal = kJ Read P –Calorimeter used to replicate metabolism in the body Carbohydrates yield the same results Proteins yield different results

Prentice Hall © 2003Chapter 5 Fuels Read text, P No problem solving on this topic, test questions on general knowledge will be asked…read carefully!

Prentice Hall © 2003Chapter 5 End of Chapter 5: Thermochemistry