Thermochemistry. Thermochemistry 1/25/2010. Reading: Chapter 5 (omit 5.8) As you read ask yourself

Transcription

1 Thermochemistry Reading: Chapter 5 (omit 5.8) As you read ask yourself What is meant by the terms system and surroundings? How are they related to each other? How does energy get transferred between them? What makes up the internal energy of matter? What kind of work do chemical system do? What impact does heat and work transfer have on the internal energy of a system? Does how we effect the transfer of heat and work matter to the value of the internal energy change? Why or why not? How can we measure heat flow? How is temperature change related to the transfer of heat? Does the composition of the substance that undergoes a change in temperature matter, if so why? Is the same thermodynamic quantity measured in all calorimetry experiments? Why or why not? Since we can t measure every reactions enthalpy directly, what methods can we use to figure out the enthalpy of any given reaction? What fundamental property of enthalpy makes Hess s Law work? Why don t we have a value for the heat of formation of O 2 (g) and C(graphite)? Chem Thermochemistry Why do chemical reactions occur? stability! products are more stable than reactants energy is an important factor in determining stability less stable Nature of energy: capacity to do work or transfer heat energy reaction more stable energy when a force causes a mass to move energy to cause an increase in temperature objects possess energy work and heat are ways to transfer energy Chem

2 Energy can be classified as either kinetic or potential energy kinetic energy energy of motion E K = ½ m v 2 thermal energy energy associated with temperature of object thermal energy depends on T and quantity potential energy stored energy arises from position or composition chemical energy energy associated with electrons and nuclei electrostatic energy interaction between charges is the force Chem Units Joules 1 Joule = 1 kg m 2 s -2 e.g. a 4 kg mass moving at 1 m/s E K = ½ m v 2 = ½ (4 kg) (1 m s -1 ) 2 = 2 kg m 2 s -2 = 2 Joules Calories defined as 1 cal = J exactly old definition: energy required to raise temp. of 1 g of water from 14.5 to 15.5 C Nutritional calorie Calorie = 1000 cal = 1 kcal Kilowatt-hour SI unit of power (rate of energy conversion) 1 watt = 1 J s watt incandescent bulb uses 100 J s -1 = 0.1 kw h 1 kw h = 3.6 x 10 6 J CFL 23 watts kw h Chem

3 First Law of Thermodynamics: energy is neither created or destroyed therefore the total energy of the universe is constant (energy is conserved). energy can be converted from one form to another energy can be transferred Chem System and surroundings to understand energy transfers and transformations we have to define the part of the universe we are studying system: part of universe chosen for study surroundings: part of universe outside of the system with which the system interacts isolated Chem

4 most observations are made in the surroundings example: chemical reaction Zn(s) = HCl(aq) Zn 2+ (aq) + H 2 (g) + 2Cl - (aq) once Zn is added system is closed mass is moved so work is done = w heat is given off = q energy is transferred Chem change in temperature energy transferred from hot to cold until equilibrium (both at same T) move an object against a force w = F x d size of force and distance moved determine quantity of work Chem

5 Internal energy, E internal energy is the sum of all the kinetic and potential energy of all components of the system hard to obtain absolute value of E normally determine the change in internal energy ΔE = E final E initial for a chemical reaction: ΔE = E products E reactants Chem Chem

6 How is ΔE changed? depends on heat transfer (q) and work (w) heat added to or liberated by system work done by or done on system ΔE = q + w the magnitude of ΔE depends on the size of q and w and their relative signs note that the values of ΔE, q and w refer to the SYSTEM Chem system loses heat, ΔE so q is negative system uses its energy & does work, ΔE so work is negative in this case E final < E initial, ΔE is < 0 system gains heat, ΔE so q is + work done on system, ΔE so w is + in this case E final > E initial, ΔE is > 0 endothermic exothermic Chem

7 example: A balloon is heated by adding 900 J of heat, it expands and does 422 J of work. What is the change in internal energy? If balloon is heated by adding 1500J it expands and does 800J of work. Suppose the balloon is then cooled by removing 520 J of heat and compressed by doing 298 J of work on the balloon Chem how the energy transfer is divided between work and heat depends on the process the total energy transferred does not depend on the pathway in the example: analogy to altitude: Chem

8 internal energy is a state function A function or property whose value depends only on the present state or condition of the system, not on the path used to arrive at that state a state function Examples of state functions: Chem ΔE in chemical reactions Most chemical changes occur at constant atmospheric pressure Heat and work are exchanged with surroundings Work is typically mechanical (change in volume of gases) or electrical expansion Chem

9 Expansion work (P-V work) work the force that moves an object through a distance pressure = force / area sign convention, when ΔV is positive, work is negative Units: w = pressure x volume = atm x L from Data Sheet: 1 J = L atm Chem Example: If mol of N 2 at constant T is compressed by 15.1 L at P = atm, what is the work involved? What does this mean for our definition of internal energy? Substitute PV work into expression for ΔE If reaction takes place in a sealed container, ΔV = 0 then no PV work is done What if external pressure is zero (expansion against a vacuum)? Chem

10 chemical changes typically take place at constant pressure define a new thermodynamic quantity for the heat change at constant pressure, enthalpy H defined as H = E + PV because E, P and V are state functions Chem Example. ΔE = kj/mol, what is sign of PΔV? What is sign and approximate magnitude of ΔH? Assume that T is constant. Chem

11 Enthalpy and chemical reactions ΔH = H final H initial so in chemical reactions 1 (ΔH is often written chemical reaction and ΔH together are thermochemical equation provides relationship between amounts of chemicals and heat involved in reaction 2H 2H 2 O 2 (l) 2H 2 O(l) + O 2 (g) ΔH = 196 kj 2 O 2 (l) kj +196 kj 2H 2 O 2 (l) 2H 2 O(g) + O 2 (g) ΔH = 108 kj 2H Chem O(l) + O 2 (g) 21 Example: Sulfuric acid is produced by reacting sulfur trioxide with water according to the equation: SO 3 (g) + H 2 O (l) H 2 SO 4 (l) ΔH = kj/mol How much heat is evolved when 75.0 g of SO 3 reacts with a stoichiometric amount of H 2 O(l)? Chem

12 Heat transfer define system and surroundings Effectiveness of heat transfer depends on nature of substances Chem Heat capacity and specific heat how much the temperature changes per amount of heat added depends on the nature of the substance Heat capacity amount of heat required to raise the temperature of an object by 1 K (or 1 C) heat capacity is proportionality constant, C heat capacity depends on Chem

13 Heat capacity for pure substances intrinsic property of substance Specific heat capacity defined for 1 gram of substance and a temperature increase of 1 C or 1 degree Kelvin amount of heat transferred specific heat = (grams of substance)(change in temperature) Molar heat capacity defined for 1 mole of substance and a temperature increase of 1 C or 1 degree Kelvin Chem Specific heat capacity (molar heat capacity) depends on bonding complexity physical state Chem

14 determining i specific heat: in step (a)the lead dis heated dto 100 C, then added to water in step(b) and the final temperature is measured in step(c) Chem Measuring heat transfer calorimetry: measure the magnitude of the temperature change as heat flows main idea: heat change in water in device heat change in the reaction measure change in T calorimeter Chem

15 Constant pressure calorimetry usually reactions are in solution what is the system? what are the surroundings? When 10.0 ml of a 1.00 M AgNO 3 solution is added to 10.0 ml of 1.00 M NaCl solution at 25.0 C in a constant pressure calorimeter, a white precipitate of AgCl is formed and the temperature of the aqueous mixture increases to 32.6 C. Assuming that the specific heat of the aqueous mixture it is 418J 4.18 g -1 C -1, that t the density of the mixture is 1.00 g ml -1, and that the calorimeter absorbs no heat, calculate ΔH in kj for the reaction. Ag + (aq) + Cl (aq) AgCl(s) Chem Constant volume calorimetry device often used for combustion reactions e.g. C 6 H 6 (l) + 15 / 2 O 2 (g) 6 CO 2 (g) + 3 H 2 O(g) heat flows out of reaction chamber into water and heats calorimeter what is the system? what are the surroundings? q rxn = ΔE why? Chem

16 Example: Combustion of a liquid rocket fuel, methylhydrazine CH 6 N 2 (s), produces CO 2 (g), N 2 (g) and H 2 O(l). When 4.00g of methyhydrazine is burned in a bomb calorimeter, the temperature increased from 25.0 to 39.5 C. The heat capacity of the calorimeter is kj/ C. What is the heat of reaction for combustion of 1.0 mole of CH 6 N 2 in the calorimeter? Chem Reaction enthalpies the enthalpy change for every reaction can not be easily measured Enthalpy is a state function, so can use known enthalpies for stepwise processes that have the initial and final states of interest N 2 (g) + 3H 2 (g) 2 NH 3 (g) ΔH rxn =? N 2 (g) + 3H 2 (g) N 2 H 4 (g) ΔH = 95.4 kj N 2 H 4 (g) + H 2 (g) 2 NH 3 (g) ΔH = kj Chem

17 Hess s Law The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps Procedure: combine the individual id reactions so their sums give the desired reaction Arrange reactions so all reactants appear on the left and all products appear on the right All intermediates must occur on both the right and the left so they cancel Any reaction that is reversed must have the sign of its H changed A reaction can be multiplied by a coefficient as necessary, but H for that reaction must be multiplied by the same coefficient. Chem Example: Calculate ΔH for the reaction S(s) + 3 / 2 O 2 (g) SO 3 (g) From the data: S(s) + O 2 (g) SO 2 (g) ΔH 1 = kj 2 SO 2 (g) + O 2 (g) 2 SO 3 (g) ΔH 2 = kj Chem

18 Hess s law requires tabulated data about the enthalpy change of a reaction many types of physical and chemical changes are tabulated enthalpy of vapourization, ΔH vap &, enthalpy of fusion, ΔH fus enthalpy of combustion, ΔH comb enthalpy of formation, ΔH f Must define the physical state, temperature and pressure to use tabulated values Chem Thermodynamic Standard state pure substance in its most stable form 1 atm pressure 25 C (chosen as reference state) 1 M concentration for all substances in solution Standard Enthalpy change, ΔH a standard enthalpy of a reaction is the enthalpy change for a reaction with reactants and products in their standard states Chem

19 Standard enthalpy of formation: the enthalpy change ΔH f for the (hypothetical) formation of 1 mol of substance in its standard state from its constituent elements in their standard states example: for ethanol: C 2 H 5 OH(l) by definition, ΔH f = 0 for The ΔH f values can be used in Hess s law Chem Example: What is the enthalpy change for this rxn? NaHCO 3 (s) Na 2 CO 3 (s) + CO 2 (g) + H 2 O(l) Chem

20 Hess s law using standard heats of formation can be generalized for the example: The value of the ΔH f for a substance can be determined in a calorimeter by measuring the heat evolved in a combustion experiment and using the known ΔH f for CO 2 and H 2 O Chem Example: Combustion of 1 gram of 2,3,4-trimethylpentane (C 8 H 18, mol. mass = 114 g/mol) in a bomb calorimeter raises the temperature of the calorimeter plus contents by 3.8 C. The calorimeter s total heat capacity is kj/ C. What is the heat of formation of 2,3,4-trimethylpentane? Chem