Thermochemistry

1. Thermochemistry

2. Thermochemistry • Thermodynamics is the science of the relationship between heat and other forms of energy. • Thermochemistry is the study of the quantity of heat absorbed or evolved by chemical reactions.

3. Energy • Energy is defined as the capacity to move matter. Energy can be in many forms: • Radiant Energy -Electromagnetic radiation. • Thermal Energy - Associated with random motion of a molecule or atom. • Chemical Energy - Energy stored within the structural limits of a molecule or atom.

4. Energy • There are three basic forms of energy: • Kinetic Energy (K.E.) is the energy associated with an object by virtue of its motion. • Potential Energy(P.E) is the energy an object has by virtue of its position in a field of force. • Internal Energy (U) is the sum of the kinetic and potential energies of the particles making up a substance. We will look at each of these in detail.

5. Energy • Kinetic Energy: An object of mass m and speed or velocity v has kinetic energy Ek equal to • This shows that the kinetic energy of an object depends on both its mass and its speed.

6. A Problem to Consider • Consider the kinetic energy of a person whose mass is 130 lb (59.0 kg) traveling in a car at 60 mph (26.8 m/s). • The SI unit of energy, kg.m2/s2, is given the name Joule (energy is also measured in calories).

7. Energy • Potential Energy: This energy depends on the “position” (such as height) in a “field of force” (such as gravity). (also: Electrical Potential) • For example, water of a given mass m at the top of a dam is at a relatively high “position” h in the “gravitational field” g of the earth.

8. A Problem to Consider • Consider the potential energy of 1000 lb of water (453.6 kg) at the top of a 300 foot dam (91.44 m).

9. Potential Energy • Electrostatic attraction, repulsion, or the energy stored in ionic or covalent bonds E = kc (q+)*(q-) r kc = 8.99 x 109 J m/C2 ; sec. 9.1 ; Energy of Attraction

10. Energy • Internal Energy is the energy of the particles making up a substance. (eg., vib., rot., trans.; use SPARTAN). • The total energy of a system is the sum of its kinetic energy, potential energy, and internal energy, U. Typically;

11. Energy • The Law of Conservation of Energy: Energy may be converted from one form to another, but the total quantities of energy remain constant.

12. Heat of Reaction • In chemical reactions, heat is often transferred from the “system” to its “surroundings,” or vice versa. • Thermodynamic system (or simply system)is thesubstance or mixture of substances under study in which a change occurs. • Thesurroundingsare everything in the vicinity of the thermodynamic system.

13. Heat of Reaction • Heat is defined as the energy that flows into or out of a system because of a difference in temperature between the system and its surroundings. • Heat flows from a region of higher temperature to one of lower temperature; once the temperatures become equal, heat flow stops. • (See Animation: Kinetic Molecular Theory/Heat Transfer)

14. Heat of Reaction • Heat is denoted by the symbol q. • The sign of q is positive if heat is absorbed by the system (endothermic; system gets cool). • The sign of q is negative if heat is evolved by the system (exothermic; system warms up). • Heat of Reaction (q) is the amount of energy gained or lost during a chemical reaction.

15. Endothermicity “into” a system Dq > 0 Heat of Reaction • Exothermicity • “out of” a system Dq < 0 Surroundings Surroundings Energy Energy System System

16. Heat of Reaction • An exothermic process is a chemical reaction or physical change in which heat is evolved (q is negative). • An endothermic process is a chemical reaction or physical change in which heat is absorbed (q is positive).

17. Enthalpy and Enthalpy Change • The heat absorbed or evolved by a reaction depends on the conditions under which it occurs. • Usually, a reaction takes place in an open vessel, and therefore at the constant pressure of the atmosphere. • The heat of this type of reaction is denoted qp, the heat at constant pressure.

18. Enthalpy and Enthalpy Change • We can show that the change in enthalpy is equal to the heat of reaction at constant pressure. Enthalpy is defined as the amount of internal energy (U) of the system plus the PV work done by the system.

19. Enthalpy and Enthalpy Change • The internal energy of a system (U), is defined as the heat (qp; at const. P) in the system + the work (w) done by the system: • Internal Energy • (See Animation: Work vs. Energy Flow) • In chemical systems, work is defined as a change in volume at a given pressure, that is:

20. Work and Pressure • Work is the effect of a system moving a object through a field of force (w = -F x d). • Recall that P = F/A, • rearranging; one obtains: F = P x A • And d = Volume/Area • Thus, w = - F x distance, or = -1 * (PxA) * DV/A w = - P DV

21. w = - F x h = - P x V

22. Enthalpy and Enthalpy Change We Know: and 1st Law of Thermo-d Path dependent; ie., not a state function (pg. 185) since, Substituting for U and w: Enthalpy and internal energy are state functions; ie. path independent

23. Enthalpy and Enthalpy Change • Since the heat at constant pressure, qp, represents DH, then • So DH is the heat released, or absorbed by a reaction.

24. Thermochemical Equations • A thermochemical equation includes a chemical equation, and the energy associated with the reaction it represents. • The enthalpy of reaction is written directly after the chemical equation.

25. Thermochemical Equations • In a thermochemical equation the enthalpy change, (DH) depends on the phase of the substances. See Table 6.2

26. Enthalpy Depends upon the phase of the substance.

27. Applying Stoichiometry and Heats of Reactions • Consider the combustion of methane, CH4, with oxygen at constant pressure. • Given the following equation, how much heat could be obtained by the combustion of 10.0 grams CH4?

28. Measuring Heats of Reaction • To measure the heat of reaction, one can monitor the change in temperature as a reaction occurs, and calculate the amount of heat transferred. • Thus, the thermochemical measurement is based on the relationship between heat and temperature change. • The heat required to raise the temperature of a substance by one degree is itsheat capacity.

29. Measuring Heats of Reaction • Heat Capacity and Specific Heat • The molar heat capacity, C, of a sample of substance is the quantity of heat required to raise 1 mole of substance one degree Celsius. • Changing the temperature of the sample requires heat equal to:

30. A Problem to Consider • Suppose a piece of iron requires 6.70 J of heat to raise its temperature by one degree Celsius (ie., the heat capacity). • What quantity of heat is required to raise the temperature of the piece of iron from 25.0oC to 35.0oC ?

31. Measuring Heats of Reaction • The specific heat capacity (or “specific heat”) is the heat required to raise the temperature of one gram of a substance by one degree Celsius. • To find the heat absorbed or released, multiply the specific heat, s, of the substance by its mass in grams, m, and by the temperature change, DT.

32. A Problem to Consider • Calculate the heat absorbed when the temperature of 15.0 grams of water is raised from 20.0oC to 50.0oC. (The specific heat of water is 4.184 J/g.oC.)

33. Heats of Reaction: Calorimetry • A calorimeter is a device used to measure the heat absorbed or evolved during a physical or chemical change. (See Figure 6.12) • The heat absorbed by the calorimeter and its contents is the negative of the heat of reaction.

34. A Problem to Consider • When 23.6 grams of calcium chloride, CaCl2, was dissolved in water in a calorimeter, the temperature rose from 25.0oC to 38.7oC. If the heat capacity of the solution and the calorimeter is 1258 J/oC, what is the enthalpy change per mole of calcium chloride?

35. Heats of Reaction: Calorimetry • First, let us calculate the heat absorbed by the calorimeter. • Now we must calculate the heat per mole of calcium chloride.

36. Heats of Reaction: Calorimetry • Calcium chloride has a molecular mass of 111.1 g, so • Now we can calculate the heat per mole of calcium chloride.

37. Hess’s Law • Hess’s lawstates relates to a chemical equation that can be written as the sum of two or more reactions. • The enthalpy change for the overall reaction is the sum of the enthalpy changes for the individual reactions. (See Animation: Hess’s Law)

38. Thermochemical Equations • The following are two important rules for manipulating thermochemical equations: • When a thermochemical equation is multiplied by any factor, the value of DH for the new equation is obtained by multiplying the DH in the original equation by that same factor. • When a chemical equation is reversed, the value of DH is reversed in sign.

39. Could you use these data to obtain the enthalpy change for the following reaction? Hess’s Law • For example, suppose you are given the following data:

40. Hess’s Law • If we multiply the first equation by 2 and reverse the second equation, they will sum together to become the third.

41. Enthalpy and Enthalpy Change • Anextensive propertyis one that depends on the quantity of substance. • Enthalpy is a state function, a property of a system that depends only on its present state and is independent of any previous history of the system. • Enthalpy, denoted H, is an extensive property of a substance that can be used to obtain the heat absorbed or evolved in a chemical reaction.

42. Enthalpy and Enthalpy Change • The change in enthalpy for a reaction at a given temperature and pressure (called the enthalpy of reaction) is obtained by subtracting the enthalpy of the reactants from the enthalpy of the products.

43. Standard Enthalpies of Formation • The term standard state refers to the standard thermodynamic conditions chosen for substances when listing or comparing thermodynamic data: 1 atmosphere pressure and the specified temperature (usually 25oC). • The enthalpy change for a reaction in which reactants are in their standard states is denoted DHo(“delta H zero” or “delta H naught”).

44. Standard Enthalpies of Formation • The standard enthalpy of formation of a substance, denotedDHfo, is the enthalpy change for the formation of one mole of a substance in its standard state from its component elements in their standard state. • Note that the standard enthalpy of formation for a pure element in its standard state is zero.

45. Standard Enthalpies of Formation • The law of summation of heats of formation states that the enthalpy of a reaction is equal to the total formation energy of the products minus that of the reactants. • Sis the mathematical symbol meaning “the sum of”, and m and n are the coefficients of the substances in the chemical equation.

46. A Problem to Consider • Large quantities of ammonia are used to prepare nitrous oxide according to the following equation: • What is the standard enthalpy change for this reaction? Use Table 6.2 for data.

47. A Problem to Consider • You record the values of DHfo under the formulas in the equation, multiplying them by the coefficients in the equation. • You can calculateDHoby subtracting the values for the reactants from the values for the products.

48. A Problem to Consider • Using the summation law: • Be careful of arithmetic signs as they are a likely source of mistakes.

49. Example Problem Calculate the DHvap of H2O using standard enthalpies of formation. H2O (l) H2O (g) DHvap = -241.8 kJ – (-285.8 kJ) = 44kJ/mol On page 426, DHvap is given as 40.7 kJ/mol. Why is there a discrepancy between that value and the one calculated here?

50. Fuels • A fuel is any substance that is burned to provide heat or other forms of energy. • In this section we will look at: • Foods as fuels • Fossil fuels • Coal gasification and liquefaction