Thermochemistry & Thermodynamics

1. Thermochemistry & Thermodynamics AP Chemistry

2. Brief Overview • Chemical reactions involve changes in energy • Thermochemistry • Study of relationships of chemical reactions and energy changes • Thermodynamics • Study of energy and its transformations

3. Energy • Energy is the ability to do work or produce heat and is the sum of all potential and kinetic energy in a system • Potential energy is energy possessed due to relative position to other objects • Commonly in chemistry this is energy stored in bonds • Kinetic energy is the energy of motion, usually of particles

4. Work & Heat • Heat (q) transfer of energy from object with higher KE to object with lower KE • Hence heat flows from hot to cold • Heat transfers because of temperature (KE) difference, but temperature is nota measure of heat directly • Work (w) is a force acting over a distance • Usually used when discussing gases • When related to gases, work is a function of pressure and volume:

5. Energy, Work, and Heat • Energy is a state function while heat and work are not • State function is a property that is independent of past or future behavior, only on current conditions (temperature, pressure, etc) • Example: Many roads from your home to school, but you have the same starting point and ending point

6. Signs of Work & Heat • Signs of q • +q means heat is absorbed (endothermic) • -q means heat is released (exothermic) • Signs of w • +w if work done on the system (i.e. compression) • -w if work done by the system (i.e. expansion)

7. First Law of Thermodynamics • Energy cannot be created or destroyed, only conserved • Law of Conservation of Energy • Energy of universe is constant • Energy may only transfers between system (the experiment) and surroundings (universe)

8. Units of Energy • SI unit of energy is Joule • 1 Joule = 1 kg*m2/s2 • Commonly reported in kilojoules • Non-SI unit of energy is calories • 1 cal = 4.184 J • Nutrional Calorie (Cal) = 1 kcal = 1000 calories

9. Enthalpy (∆H) • Describes heat gained or lost at constant pressure H = qp • Following the derivation, E = H + P∆V or H = E - P∆V • Enthalpy is a state function • Can be calculated from multiple sources: • Calorimetry • Stoichiometry • Heats of formation (tables of standard values) • Hess’ Law • Bond Energies

10. Enthalpy of Reaction • Enthalpy of a reaction can be utilized stoichiometrically based on coefficients • For endothermic processes ∆H is positive and should be included as a reactant in the rxn • For exothermic processes ∆H is negative and should be included as a product in the rxn

11. Enthalpy of Reaction • Example Upon adding solid potassium hydroxide pellets to water the following reaction takes place: Answer the following questions regarding the addition of 14.0 g of KOH to water: • Does the beaker get warmer or colder? • Is the reaction endothermic or exothermic? • What is the enthalpy change for the dissolution of the 14.0 g of KOH?

12. Thermochemical Equations • Guidelines • Enthalpy is an extensive property • Enthalpy of the reverse reaction is equal in magnitude but opposite in sign • Enthalpy change depends on physical states of reactants and products • If the thermochemical equation is multiplied by a factor of n, then ∆H must also change by the same factor

13. Calorimetry • Process of measuring heat based on observing the temperature change when a body absorbs or discharges energy as heat • Assumption is that no heat is lost to surroundings (aka closed system) • 2 types of Calorimetry: • Constant pressure(coffee-cup setup) • Constant volume (bomb calorimeter) • Used in industry for determining food calories

14. Calorimetry • Commonly used in lab to determine specific heats of metals • Can also measure the heat released/absorbed of many types of reactions: • Neutralizations • Ionization (Dissolution/Dissolving) • Reaction (ppt, combustion, etc.)

15. Terms to know (and love!) • Heat capacity– energy required to raise the temperature of an object by 1 degree C (J/°C) • Specific heat capacity– same as above but specific to 1 gram of the substance (J/ g °C) • Molar heat capacity – Same as specifc heat capacity, but specific to one mole of a substance (J/mol K or J/mol °C)

16. Relationships of Heat • Heat capacities are extensive properties, while specific heat is an intensive property • Heat capacity can be determined by multiplying a substances mass by its specific heat: J/°C = J/g°C x g • Specific heat of water is 4.184 J/g°C and is the same for all dilute aqueous solutions • Heat of substance + heat of solution = 0 • Therefore, qsubstance = -qsoln

17. Calculating Specific Heat q=mCp∆T • q= heat transferred (J) • Recall: at constant pressure q=∆H • m = mass (g) • Cp = specific heat of material at constant pressure • ∆T = Tf – Ti (final – initial)

18. Calculating Specific Heat • Example How much heat is needed to raise 10.0 grams of aluminum from 22.0 °C to 42.0 °C? (Specific heat of aluminum is 0.90 J/ g K) Also, what is the molar heat capacity of aluminum?

19. Constant Pressure Calorimetry • Example (Specific Heat of Metal) A lead (Pb) pellet having a mass of 26.47 g at 89.98C was placed in a constant-pressure calorimeter of negligible heat capacity containing 100.0 mL of water. The water temperature rose from 22.50 °C to 23.17 °C. What is the specific heat of the lead pellet? 0.158 J/g˚C

20. Constant Pressure Calorimetry • Example (Molar Heat of Neutralization) A quantity of 1.00 x 102mL of 0.500 M HCl was mixed with 1.00 x 102mL of 0.500 M NaOH in a constant-pressure calorimeter of negligible heat capacity. The initial temperature of the HCl and NaOH solutions was the same, 22.50 °C, and the final temperature of the mixed solution was 25.86 °C. Calculate the heat change for the neutralization reaction on a molar basis (that is, the molar heat of neutralizaton) -56.2 kJ/mol

21. Constant Volume Calorimetry • In order to perform calculations using a constant volume calorimeter, the bomb calorimeter must be calibrated so that the heat capacity of the calorimeter is known qcal = Ccal∆T (where Ccal is the heat capacity of the calorimeter) • The Ccal is determined by burning a substance with an accurately known heat of combustion • This is a constant that is generally given! Finally, qcal= -qrxn

22. Constant Volume Calorimetry • Example (Molar Heat of Combustion) A quantity of 1.435 g of naphthalene (C10H8) was burned in a bomb calorimeter. The temperature of the water rose from 20.28 °C to 25.95 °C. If the heat capacity of the bomb plus water is 10.17 kJ/°C, calculate the heat of combustion of naphthalene on a molar basis. -5.151x103 kJ/mol

23. Constant Volume Calorimetry • Example (Heat Capacity of Bomb Calorimeter) Camphor (C10H17O) has a heat of combustion of 5903 kJ/mol. When a sample of camphor with mass of 0.1204 g is burned in a bomb calorimeter, the temperature increases by 2.28 °C. Calculate the heat capacity of the calorimeter

24. Standard Enthalpy of Formation & Reaction • Heat (or enthalpy) of formation, ∆Hf, is the heat required to form the elements • Standard enthalpies of formation (∆H°f) can be used as reference points for determining the standard enthalpy of a reaction (∆H°rxn) • Standard state conditions are 1 atm and 25 °C • Standard enthalpies of formation of any element in its most stable form is zero • The greater the heat of formation of a molecule, the less stable the molecule • Higher ∆H°fimplies more energy to form the bonds

25. Standard Enthalpy of Formation & Reaction • Change of enthalpy that occurs in a chemical reaction can be given by: • Where n and m represent coefficients • Two methods exist for determining the ∆H°rxnfor a reaction • Direct Method (using ∆H°f) • Indirect Method (using Hess’ Law)

26. Standard Enthalpy of Formation & Reaction • When to use Direct Method? • If all of the standard heats of formation are known for each participant in a chemical equation, then plug them into the equation on the last slide • When to use Indirect Method? • The heat of a reaction is independent of the steps it takes to get there (enthalpy is a state function) • If only heats of reaction for a series of reactions is known, then these can be manipulated to determine the heat of the reaction • This is known as Hess’s Law

27. Direct Method • Example (Enthalpy of Reaction) • Calculate the ∆H°rxn for the following: Given the following values: • 3 Al (s) + 3 NH4ClO4 (s) → Al2O3 (s) + AlCl3 (s) + 3 NO (g) + 6 H2O (g) Answer: -2680 kJ/mol (exo)

28. Direct Method • Example (Enthalpy of Formation) • Sometimes all values are not found in the table of thermodynamic data. For most substances it is impossible to go into a lab and directly synthesize a compound from its free elements. The heat of formation for the substance must be found by working backwards from its heat of combustion. Find the ∆Hf of C6H12O6 (s) from the following information: C6H12O6 (s) + 6 O2 (g) → 6 CO2 (g) + 6 H2O (l) + 2800 kJ Answer: -1276 kJ.mol for glucose

29. Hess’s Law (of heat summation) • Guidelines • First decide how to rearrange equations so reactants and products are on appropriate sides of the arrows • Manipulate the equations using the thermochemical equation guidelines (mentioned previously) • Check to ensure that everything cancels out to give you the exact equation you want • Note: It is often helpful to begin your work backwards from the answer that you want!

30. Hess’s Law • Given the following equations: H3BO3 (aq) → HBO2 (aq) + H2O (l) ∆Hrxn= -0.02 kJ/mol H2B4O7 (aq) + H2O (l) → 4 HBO2 (aq) ∆Hrxn= -11.3 kJ/mol H2B4O7 (aq) → 2 B2O3 (s) + H2O (l) ∆Hrxn= 17.5 kJ/mol Find the ∆Hfor this overall reaction: 2 H3BO3 (aq) → B2O3 (s) + 3 H2O (l) Answer: 14.4 kJ/mol (endothermic)

31. Thermodynamics! • Study of energy changes in chemistry • Involves three major players: ∆H, ∆S, ∆G • One of the major objectives of thermodynamics is to predict whether or not a reaction will occur when reactants are brought together (spontaneous vs. nonspontaneous) • First Law of Thermodynamics (stated previously) is that energy cannot be created or destroyed, only conserved

32. Thermodynamics! • Enthalpy can be used to predict spontaneity of a reaction • In general, lower energy states are preferred, meaning exothermic reactions are favorable (-∆H)

33. Entropy (∆S) • Enthalpy (∆H) is not the only predictor of spontaneity • Entropy (∆S) can be used, and is the measurement of disorder in a system • Second Law of Thermodynamics states that the entropy (or disorder) of the universe is constantly increasing • Examples: Ice cube melting, your room at the end of a week. Nature tends towards chaos!

34. Entropy (∆S) • + ∆S = more disorder in the system • Increase in disorder is favorable, indicates higher chance to be spont. • - ∆S = less disorder in the system • Decrease in disorder is not favorable, indicates lower chance to be spont. • Thermodynamics can predict spontaneity, but does NOT predict the rate (speed) of the reaction! • Talking state functions: we do not look at pathways, only beginning and end states!

35. Twice is extra nice! • Entropy, like enthalpy, is a state function • Entropy can be calculated for reactions exactly like enthalpy: • Units are J/K*mol (where ∆H is kJ/mol)

36. Changes in Entropy • Entropy increases as one goes from solid to a liquid Ssolid  Sliquid  Sgas

37. Changes in Entropy • Entropy increases as a substance divides into parts • Dissolving solids into liquids (exception: carbonates!)

38. Changes In Entropy • Entropy will increase in reactions in which number of product molecules are greater • Entropy increases with temperature

39. Third Law of Thermodynamics • States that the entropy of a perfect crystal at 0 K is zero • This implies no movement, and no randomness exists • Not many perfect crystals out there, so entropy values rarely ever zero • For determination of standard values of entropy (as in ∆S°), this allows an absolute standard to exist which all other values can be based on

40. Free Energy • To ultimately decide the spontaneity of a reaction, we use Gibb’s free energy (G) • For a constant temperature process, change in free energy can be given by: • The most important thermodynamic equation, and one of the most beneficial equations in chemistry!

41. Free Energy & Spontaneity • Free energy is the energy available to do work • - ∆G means forward reaction is spontaneous • + ∆G means forward reaction is non-spontaneous • Reverse reaction would be spontaneous! • ∆G = 0 means reaction is at equilibrium (more on this later)

42. Third Time’s the Charm! • Standard free enthalpies of reaction can be calculated using the following equation: • ∆G has units of energy (J or kJ) • ∆G is a state function (like ∆H and ∆S)

43. Predicting Reaction Spontaneity • Both enthalpy and entropy must be known to predict spontaneous reactions