Thermochemistry: The Study of Energy and Chemical Reactions

1. Chapter 5Thermochemistry Prof. Dr. Nizam M. El-Ashgar Chemistry Department, IUG

2. Thermochemistry Thermochemistry:is the study of the relationship between ENERGY and chemical reactions. What is energy? Energy: is the ability to do work or transfer heat. Work:Energy used to cause an object that has mass to move . Heat: Energy used to cause the temperature of an object to rise.

3. Potential Energy Potential energy: is energy of an object possesses by virtue of its position or chemical composition. • chemical energy : • stored within structural units of chemical substances.

4. Potential Energy under the microscope Gravitational forces play little to no role in interactions between atoms and molecules. Predict what plays a much more important role. Electrostatic potential energy, Eel: One of the most important form of P.E. This energy is proportional to electrical charges on two interacting objects, Q1 and Q2, and inversely proportional to the distance, d. Predict an equation.

5. Chemical Energy is Mainly Potential Energy Electrostatic Potential Energy (Eel) • The most important form of potential energy in molecules is electrostatic potential energy, Eel: • κ is a constant of proportionality = 8.99 x 109J.m/C2 • How can we make electrostatic potential energy absolute zero? Answer: if d is  (infinite)

6. At finite separation distances for two charged particles: Eelis positive for like charges and negative for opposite charges. As the particles move farther apart, their electrostatic potential energy approaches zero.

7. Attraction between Ions • Electrostatic attraction is seen between oppositely charged ions. • Energy is released when chemical bonds are formed; energy is consumed when chemical bonds are broken.

8. 1 KE =  mv2 2 Kinetic Energy Kinetic energy is energy an object possesses by virtue of its motion (depends on mass and speed).

9. Thermal energy is kinetic energy in the form of random motion of particles in any sample of matter. As temperature rises, thermal energy increases. • Heat is energy that causes a change in the thermal energy of a sample. Addition of heat to a sample increases its temperature.

10. kg m2 1 J = 1  s2 Units of Energy • The SI unit of energy is the joule (J). A mass of 2 kg moving at a speed of 1m/s possesses a kinetic energy of 1 J: • An older, non-SI unit is still in widespread use: the calorie (cal). Originally defined as: amount of energy require to raise the temperature of 1 g of water from 14.5 C to 15.5 C. 1 cal = 4.184 J exactly In nutrition is the nutritional Calorie (note the capital C): 1 Cal = 1000 cal = 1 kcal

11. Definitions: System and Surroundings The system: The portion of universe under study. The surroundings: Everything else is called (remain of universe) • The system includes the molecules we want to study (here, the hydrogen and oxygen molecules). • The surroundings are everything else (here, the cylinder and piston). Note that this is a CLOSED system.

12. Types of systems Open Closed Isolated Open system: matter and energy can be exchanged with surroundings (ie: boiling water) Closed System: can exchange energy but not matter with surroundings (ie: soft drink bottle) Isolated System: no exchange of matter or energy. (ie: coffee thermos).

13. Definitions: Force , Work & Heat • Force F: is any push or pull exerted on an object. • Work w : Energy used to move an object over some distance is work. w = Fd where w is work, F is the force, and d is the distance over which the force is exerted. • Heat q: Energy transferred (as heat) from warmer objects to cooler objects.

14. Example Problem 5.1 • A bowler lifts a 5.4 kg (12 lb) bowling ball from ground level to a height of 1.6 m (5.2 ft) and then drops it. a) What happens to the potential energy of the ball as it is raised? Because the bowling ball is raised to a greater height above the ground, its potential energy increases. b) What is the quantity of work, in J, used to raise the ball? (Note: The force due to gravity is F = m x g, where m is the mass of the object and g is the gravitational constant; g = 9.8 m/s2.)

15. c) After the ball is dropped, it gains kinetic energy. If all the work done in part b has been converted to K.E at the point of impact with the ground, what is the speed of the ball at the point of impact? (5.6 m/s) PE converted to KE:

16. THE FIRST LAW OF THERMODYNAMICS • Energy can neither be created nor destroyed • Any energy lost by a system must be gained by the surroundings (vice versa) • In other words, the total energy of the universe is a constant (Energy is conserved) • Energy can be converted from one type to another. Potential energy is converted to kinetic energy.

17. Internal Energy: IE The internal energy of a system is the sum of all kinetic (molecular motion)and Potential energies(attractive/repulsive interactions) of all components of the system; we call it E. Note: Numerical values of a systems IE are not known In thermodynamics, we are mainly concerned with the changes in E of the system (ΔE).

19. Internal Energy By definition, the change in internal energy, E, is the final energy of the system minus the initial energy of the system: E = Efinal – Einitial Efinal & Einitial(can’t be determined exactly) for a practical system we need only Eto apply the law To deal with thermodynamic values (Ex. E ) we need three things: 1- number 2- unit 3- sign

21. Changes in Internal Energy 2 H2(g) + O2(g)  2 H2O(l) • If E > 0, Efinal > Einitial • Therefore, the system absorbed energy from the surroundings. • This energy change is called endothermic • E is +ve

22. Changes in Internal Energy • If E < 0, Efinal < Einitial • Therefore, the system released energy to the surroundings. • This energy change is called exothermic • E is -ve

23. Changes in Internal Energy • When energy is exchanged between the system and the surroundings, it is exchanged as either heat (q) or work (w). We algebraically represent a change in internal energy by: ∆E = q + w

25. E, q, w, and Their Signs

27. Sample Exercise 5.2 • Two gases, A(g) and B(g), are confined in a cylinder-and-piston arrangement like that in the Figure below. Substances A and B react to form a solid product: A(g) + B(g)  C(s). As the reaction occurs, the system loses 1150 J of heat to the surroundings. The piston moves downward as the gases react to form a solid. As the volume of the gas decreases under the constant pressure of the atmosphere, the surroundings do 480 J of work on the system. What is the change in the internal energy of the system? ∆E = q + w q = -1150 J (heat lost from sys, to surr.) W = + 480 J (work done by sys. On surr.) So, ∆E = -1150 J + 480 J = -670 J The system transfers heat to surrounding

28. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic. • When heat is released by the system into the surroundings, the process is exothermic. Ice melting, Evaporation of acetone Combustion of Gasoline

31. State Functions • A state function is a property of a system that is determined only by the present state of the system, not the path that system took. 50 oC 150 oC - 90 oC 30 oC -50 oC 57 oC 39 oC 120 oC Ti Tf • T = Tf – Ti = 120 – 50 = 70 oC Examples of state functions: V, E, P, H, n =

32. State Functions • The internal energy of a system is independent of the path by which the system achieved that state. • In the system below, the water could have reached room temperature from either direction. And so, E depends only on Einitial and Efinal. , (state function) E = Efinal ─ Einitial

33. State Functions q max w max • However, q and w are not state functions. • Whether the battery is shorted out or is discharged by running the fan, its E is the same. • But q and w are different in the two cases. • Path a: E1 = q1+w1 • Path b: E2 = q2+w2 • E1 = E2 • q1  q2 and w1  w2

34. Work In an open container: The only work done is by a gas pushing on the surroundings (or by the surroundings pushing on the gas).

35. Work We can measure the work done by the gas if the reaction is done in a vessel that has been fitted with a piston. W = − PV (Expansion) Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)

36. Enthalpy • If a process takes place at constant pressure (as the majority of processes we study do), ieopen systems. • The only work done is this pressure-volume work. We can account for heat flow during the process by measuring the enthalpy (H) of the system • Enthalpy is the internal energy plus the product of pressure and volume: H = E + PV

37. Pressure – Volume Work • Work involved in expansion or compression of gas • When pressure is constant: W = P∆V ∆V is +ve(expansion) and -ve(compression) • Q: If a system does not change its volume during the course of a process, does it do pressure-volume work? No work done ie w = 0 because ∆V = 0

38. Enthalpy • When the system changes at constant pressure, the change in enthalpy, H, is derived as follows: H = (E + PV) • This can be written H = E + PV------------------(1) • Since E = q + w and w = -PV, we can substitute these into the enthalpy expression: E = qP-PV qp = E + PV---------------- (2) From (1) and (2) : H = qp • So, at constant pressure, the change in enthalpy is the heat gained or lost. H = E + PV

39. Endothermicity and Exothermicity • A process is endothermic when H is positive. • A process is exothermic when H is negative.

40. Enthalpy of Reaction ,H H is a state function In reactions: The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants: H = Hproducts-Hreactants This quantity, H, is called the enthalpy of reaction, or the heat of reaction.

42. Sample Excercise 5.3 • Indicate the sign of the enthalpy change, ∆H, in these processes carried out under atmospheric pressure and indicate whether each process is endothermic or exothermic: • An ice cube melts: Endothermic • 1 g of butane is combusted in sufficient oxygen to give complete combustion to CO2 and H2O: Exothermic

43. Thermochemical Reactions • A thermochemical equation is a balanced chemical equation that shows the associated enthalpy change but does not specify amount of chemical involved. • Consider the thermochemical equation • The thermochemical equivalencies are • 1 mol C3H8(g) react to give off 2.22×103 kJ • 5 mol O2(g) react to give off 2.22×103 kJ • 3 mol CO2(g) are formed and give off 2.22×103 kJ • 4 mol H2O() are formed and give off 2.22×103 kJ C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O() DH = -2.22×103 kJ Mass of CO2 Moles of CO2 Enthalpy change Molar mass of CO2 Thermochemical equivalent

44. The Truth about Enthalpy 1- Enthalpy is an extensive property. Directly proportional to the amount of material Ex: CH4(g) + 2O2(g) → CO2(g) + 2H2O() DH = -890 kJ 1 mol CH4(g)  -890 kJ 2 mol CH4(g)  -1780 kJ 2- H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction. CO2(g) + 2H2O() → CH4(g) + 2O2(g) DH = +890 kJ

45. 3- H for a reaction depends on the state of the products and the state of the reactants. • CH4(g) + 2O2(g) → CO2(g) + 2H2O() DH = -890 kJ • CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) DH = -802 kJ As: 2H2O() → 2H2O(g) DH = +88 KJ 4- Multiplying or dividing the equation by a factor (n) the same thing should be for the value of DH 2H2O() → 2H2O(g) DH = +88 KJ H2O() → H2O(g DH = +44 KJ

46. Sample Exercise 5.4 • How much heat is released when 4.50 g of methane gas is burned in a constant pressure system? CH4(g) + 2O2(g) → CO2(g) + 2H2O() DH = -890 kJ 4.50 g ?? Heat 1mol CH4(g)  = -890 kJ

47. Example: Given the following equation: C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l) ΔH =  2803 kJ/mol calculate the mass in grams of glucose burned in oxygen for releasing 700.0 KJ of energy.

48. Practice Exercise Hydrogen peroxide can decompose to water and oxygen by the following reaction: 2 H2O2(l)  2 H2O(l) + O2(g) ΔH = –196 kJ Calculate the value of q when 5.00 g of H2O2(l) decomposes at constant pressure. Answer: –14.4 kJ

49. Calorimetry • We cannot know the exact enthalpy of the reactants and products. We can measure H through calorimetry. • Calorimetry is the experimental measurement of heat released or absorbed by a chemical reaction. • Acalorimeter is the device used to measure heat. • The quantity of heat transferred by the reaction causes a change of temperature inside the calorimeter. A simple calorimeter can be made from styrofoam coffee cups.

50. Heat Capacity and Specific Heat Heat capacity C: The amount of energy required to raise the temperature of a substance by 1 K (1C). Specific heat CS: The amount of energy required to raise the temperature of 1 g of a substance by 1 K (1C). Molar heat capacity Cm: The amount of energy required to raise the temperature of 1 mole a substance by 1 K (1C).