Chemistry Chapter 16. Reaction Energy

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1 Chemistry Reaction Energy

2 Section 16.1.I Thermochemistry

3 Objectives Define temperature and state the units in which it is measured. Define heat and state its units. Perform specific-heat calculations. Explain enthalpy change, enthalpy of reaction, and enthalpy of formation. Solve problems involving enthalpies of reaction and enthalpies of formation.

4 Thermochemistry Virtually EVERY chemical reaction is accompanied by a change in energy. Chemical reactions usually either absorb or release energy as heat. Thermochemistry - the study of the transfers of energy as heat that accompany chemical reactions and physical changes

5 Thermochemistry

6 Heat and Temperature Calorimeter - instrument used to measure the energy absorbed or released as heat in a chemical or physical change Known quantities of reactants are sealed in a reaction chamber that is immersed in a known quantity of water. Energy given off by the reaction is absorbed by the water, and the temperature change of the water is measured. From the temperature change of the water, it is possible to calculate the energy as heat given off by the reaction.

7 Heat and Temperature, continued Temperature - a measure of the average kinetic energy of the particles in a sample of matter The greater the kinetic energy of the particles in a sample, the hotter it feels. For calculations in thermochemistry, the Celsius and Kelvin temperature scales are used. Celsius and Kelvin temperatures are related by the following equation. K = C

8 Heat and Temperature, continued The amount of energy transferred as heat is usually measured in joules. Joule - the SI unit of heat and ALL other forms of energy Heat - the energy transferred between samples of matter because of a difference in their temperatures Energy transferred as heat always moves spontaneously from matter at a higher temperature to matter at a lower temperature.

9 Specific Heat The amount of energy transferred as heat during a temperature change depends on the nature of the material and its mass The specific heat of a substance is the amount of energy required to raise the temperature of one gram by one Celsius degree (1 C) or one kelvin (1 K). The temperature difference as measured in either Celsius degrees or kelvins is the same.

10 Specific Heat, continued Values of specific heat are given in units of joules per gram per Celsius degree, J/(g C), or per kelvin, J/(g K).

11 Specific Heat, continued Specific heat is calculated according to the equation given below. c p = c p is the specific heat at a given pressure, q is the energy lost or gained, m is the mass of the sample, and T is the difference between the initial and final temperatures. The above equation can be rearranged to given an equation that can be used to find the quantity of energy gained or lost with a change of temperature. q m T q = m cp T

12 Specific Heat, continued Sample Problem A A 4.0 g sample of glass was heated from 274 K to 314 K, a temperature increase of 40. K, and was found to have absorbed 32 J of energy as heat. a. What is the specific heat of this type of glass? b. How much energy will the same glass sample gain when it is heated from 314 K to 344 K?

13 Specific Heat, continued Sample Problem A Solution Given: m = 4.0 g T = 40. K q = 32 J Unknown: a. c p in J/(g K) b. q for T of 314 K 344 K Solution: a. c p q 32 J m T (4.0 g)(40. K) 0.20 J/(g K)

14 Specific Heat, continued Sample Problem A Solution, continued Solution: b. q cp m T q 0.20 J (4.0 g)(344 K 314 K) (g K) 0.20 J q (4.0 g)(30 K) 24 J (g K)

15 Section 16.1.II Enthalpy

16 Objectives Define temperature and state the units in which it is measured. Define heat and state its units. Perform specific-heat calculations. Explain enthalpy change, enthalpy of reaction, and enthalpy of formation. Solve problems involving enthalpies of reaction and enthalpies of formation.

17 Enthalpy of Reaction Enthalpy ( H ) - the energy absorbed as heat during a chemical reaction at constant pressure Only changes in enthalpy can be measured. H is read as change in enthalpy. Enthalpy change - the amount of energy absorbed by a system as heat during a process at constant pressure Enthalpy change = the difference between the enthalpies of products and reactants H = H products H reactants

18 Enthalpy of Reaction, continued Exothermic reaction - chemical reaction that releases energy, the energy of the products is less than the energy of the reactants. example: 2H 2 (g) + O 2 (g) 2H 2 O(g) kj

19 Enthalpy of Reaction, continued 2H 2 (g) + O 2 (g) 2H 2 O(g) kj The above is a thermochemical equation, an equation that shows the quantity of energy released or absorbed as heat during the reaction Chemical coefficients in a thermochemical equation must be interpreted as numbers of moles and never as numbers of molecules.

20 Enthalpy of Reaction, continued Quantity of energy released ~ quantity of the reactants & products Producing 2x water in the equation below would require 2x moles of reactants and would release kj of energy as heat 2H 2 (g) + O 2 (g) 2H 2 O(g) kj Producing 4 moles of water releases J of energy

21 Enthalpy of Reaction, continued Endothermic reaction - the products have a higher energy than the reactants, and the reaction absorbs energy example: 2H 2 O(g) kj 2H 2 (g) + O 2 (g)

22 Enthalpy of Reaction, continued Thermochemical equations: a) Include physical states of reactants and products b/c states of matter influence the overall amount of energy exchanged b) Usually written by designating a H value rather than writing the energy as a reactant or product c) Include a + or sign on H to indicate if the system gains or loses energy 2H 2 (g) + O 2 (g) 2H 2 O(g) H = kj

23 Enthalpy of Reaction, continued In an exothermic reaction, energy is evolved, or given off, during the reaction; H is negative.

24 Enthalpy of Reaction, continued In an endothermic reaction, energy is absorbed; in this case, H is designated as positive.

25 Enthalpy of Formation The molar enthalpy of formation is the enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state at 25 C and 1 atm. Why give enthalpies of formation for a standard temperature and pressure? To signify standard states, a 0 sign is added to the enthalpy symbol, and the subscript f indicates a standard enthalpy of formation: H 0 f

26 Enthalpy of Formation, continued Each entry in the table is the enthalpy of formation for the synthesis of one mole of the compound from its elements in their standard states.

27 Stability and Enthalpy of Formation Compounds with a large negative enthalpy of formation are very stable. H f 0 example: the of carbon dioxide is kj per mol of gas produced. Elements in their standard states are defined as 0 having H f = 0. This indicates that carbon dioxide is more stable than the elements from which it was formed.

28 Stability and Enthalpy of Formation, continued Compounds with positive values of enthalpies of formation are typically unstable. 0 example: hydrogen iodide, HI, has a of kj/mol. H f It decomposes at room temperature into violet iodine vapor, I 2, and hydrogen, H 2.

29 Section 16.1.III Calculating Enthalpies of Reaction

30 Objectives Define temperature and state the units in which it is measured. Define heat and state its units. Perform specific-heat calculations. Explain enthalpy change, enthalpy of reaction, and enthalpy of formation. Solve problems involving enthalpies of reaction and enthalpies of formation.

31 Calculating Enthalpies of Reaction a) The basis for calculating enthalpies of reaction is known as Hess s law b) Hess s law - the overall H in a reaction is equal to the sum of H for the individual steps in the process c) The energy difference between reactants and products is independent of the route taken to get from one to the other.

32 Calculating Enthalpies of Reaction, continued If you know the reaction enthalpies of individual steps in an overall reaction, you can calculate the overall enthalpy without having to measure it experimentally. What is the enthalpy of formation for the formation of methane gas, CH 4, from its elements, hydrogen gas and solid carbon: C(s) + 2H 2 (g) CH 4 (g) H 0? f

33 Calculating Enthalpies of Reaction, continued C(s) + 2H 2 (g) CH 4 (g) The component reactions in this case are the combustion reactions of carbon, hydrogen, and methane: a) C(s) + O 2 (g) CO 2 (g) 0 H c kj b) H 2 (g) + ½O 2 (g) H 2 O(l) 0 H c kj c) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) 0 H c kj

34 Calculating Enthalpies of Reaction, continued C(s) + 2H 2 (g) CH 4 (g) The overall reaction involves the formation rather than the combustion of methane, so the combustion equation for methane is reversed, and its enthalpy changed from negative to positive: c) CO 2 (g) + 2H 2 O(l) CH 4 (g) + 2O 2 (g) H 0 = kj

35 Calculating Enthalpies of Reaction, continued CO 2 (g) + 2H 2 O(l) CH 4 (g) + 2O 2 (g) Because 2 moles of water are used as a reactant in reversed reaction (c), 2 moles of water will be needed as a product. H 2 (g) + ½O 2 (g) H 2 O(l) Therefore, the coefficients for the formation of water reaction, as well as its enthalpy, need to be multiplied by 2: 2H 2 (g) + O 2 (g) 2H 2 O(l) 2( kj) H 0 c

36 Calculating Enthalpies of Reaction, continued We are now ready to add the three equations together using Hess s law to give the enthalpy of formation for methane and the balanced equation. C(s) + O 2 (g) CO 2 (g) 0 H c kj 2H 2 (g) + O 2 (g) 2H 2 O(l) CO 2 (g) + 2H 2 O(l) CH 4 (g) + 2O 2 (g) 2( kj) 0 H c H kj C(s) + 2H 2 (g) CH 4 (g) 0 H f 74.3 kj

37 Calculating Enthalpies of Reaction from H f 0 Using Hess s law, any enthalpy of reaction may be calculated using enthalpies of formation for all the substances in the reaction of interest, without knowing anything else about how the reaction occurs. Mathematically, the overall equation for enthalpy change will be in the form of the following equation: H 0 = sum of [( H 0 of products) (mol of products)] f sum of [( H 0 of reactants) (mol of reactants)] f

38 Calculating Enthalpies of Reaction, continued Sample Problem B Calculate the enthalpy of reaction for the combustion of nitrogen monoxide gas, NO, to form nitrogen dioxide gas, NO 2, as given in the following equation. NO(g) + ½O 2 (g) NO 2 (g) Use the enthalpy-of-formation data in the appendix on page 862. Solve by combining the known thermochemical equations.

39 Calculating Enthalpies of Reaction, continued Sample Problem B Solution Given: N ( g ) + O ( g ) NO( g ) H = kj N ( g) + O ( g) NO ( g) ΔH =+33.2 kj f f Unknown: 0 1 H g + 2 g 2 g 2 for NO( ) O ( ) NO ( ) Solution: Using Hess s law, combine the given thermochemical equations in such a way as to obtain the unknown equation, and its H 0 value.

40 Calculating Enthalpies of Reaction, continued Sample Problem B Solution, continued The desired equation is: 1 NO( g ) + O ( g ) NO ( g) Reversing the first given reaction and its sign yields the following thermochemical equation: NO( g ) N ( g ) + O ( g ) H = kj f The other equation should have NO 2 as a product, so we can use the second given equation as is: 0 N ( g) + O ( g) NO ( g) ΔH =+33.2 kj f

41 Calculating Enthalpies of Reaction, continued Sample Problem B Solution, continued We can now add the equations and their H 0 values to obtain the unknown H 0 value NO( g ) N ( g ) + O ( g ) H = kj f 0 N ( g) + O ( g) NO ( g) H =+33.2 kj f 1 NO( g ) + O ( g ) NO ( g) H kj

42 Section 16.2 Driving Force of Reactions

43 Objectives Explain the relationship between enthalpy change and the tendency of a reaction to occur. Explain the relationship between entropy change and the tendency of a reaction to occur. Discuss the concept of free energy, and explain how the value of this quantity is calculated and interpreted. Describe the use of free energy change to determine the tendency of a reaction to occur.

44 Section 2 Driving Force of Reactions Enthalpy and Reaction Tendency The great majority of chemical reactions in nature are exothermic. Everything in nature, including reactions, proceed in a direction that leads to a lower energy state. Some endothermic reactions do occur spontaneously. Something other than enthalpy change can help determine whether a reaction will occur.

45 Section 2 Driving Force of Reactions Entropy and Reaction Tendency Melting is one example of a naturally occurring endothermic process. An ice cube melts spontaneously at room temperature as energy is transferred from the warm air to the ice. Well-ordered arrangement of the ice crystal is lost, and the less-ordered liquid phase of higher energy content is formed. A system that can go from one state to another without an enthalpy change does so with an increase in entropy.

46 Section 2 Driving Force of Reactions Entropy and Reaction Tendency, continued The decomposition of ammonium nitrate: 2NH 4 NO 3 (s) 2N 2 (g) + 4H 2 O(l) + O 2 (g) On the left side are 2 mol of solid ammonium nitrate. The right-hand side of the equation shows 3 mol of gaseous molecules plus 4 mol of a liquid. The arrangement of particles on the right-hand side of the equation is more random than the arrangement on the left side and hence is less ordered.

47 Section 2 Driving Force of Reactions Entropy and Reaction Tendency, continued Nature ALSO tends to proceed in a direction that increases the randomness of a system. A random system is one that lacks a regular arrangement of its parts. This tendency toward randomness is called entropy. Entropy, S - as a measure of the degree of randomness of the particles, such as molecules, in a system.

48 Section 2 Driving Force of Reactions Standard Entropy Changes for Some Reactions

49 Section 2 Driving Force of Reactions Entropy and Reaction Tendency, continued Compare entropy in solids, liquids and gases: State Position Location Determination Randomness Entropy Solid Liquid Fixed, very close Moving, close Easy Low Low Difficult Medium Medium Gas Moving fast, far apart Very difficult High High

50 Section 2 Driving Force of Reactions Entropy and Reaction Tendency, continued Compare entropy in solids, liquids and gases:

51 Section 2 Driving Force of Reactions Entropy and Reaction Tendency, continued Absolute entropy, or standard molar entropy, of substances are recorded in tables and reported in units of kj/(mol K). Entropy change - the difference between the entropy of the products and the reactants Increase in entropy = + S Decrease in entropy = - S

52 Section 2 Driving Force of Reactions Free Energy Processes in nature are driven in two directions: toward least enthalpy and toward largest entropy. How do you know which factor will dominate and whether a reaction will occur? A function has been defined to relate the enthalpy and entropy factors at a given temperature and pressure. This combined enthalpy-entropy function is called the free energy, G, of the system; it is also called Gibbs free energy.

53 Section 2 Driving Force of Reactions Free Energy, continued Only the change in free energy can be measured. It can be defined in terms of enthalpy and entropy at a constant pressure and temperature: G 0 = H 0 T S 0 H = change in enthalpy, T= temp in Kelvins, S = change in entropy

54 Section 2 Driving Force of Reactions Free Energy, continued G 0 = H 0 T S 0 If G < 0, the reaction is spontaneous. H and G can have positive or negative values. This leads to four possible combinations of terms.

55 Section 2 Driving Force of Reactions Free Energy, continued Sample Problem D For the reaction NH 4 Cl(s) NH 3 (g) + HCl(g), at K, H 0 = 176 kj/mol and S 0 = kj/(mol K). Calculate G 0, and tell whether this reaction is spontaneous in the forward direction at K.

56 Section 2 Driving Force of Reactions Free Energy, continued Sample Problem D Solution Given: H 0 = 176 kj/mol at K S 0 = kj/(mol K) at K Unknown: G 0 at K Solution: The value of G 0 can be calculated according to the following equation: G 0 = H 0 T S 0 G 0 = 176 kj/mol 298 K [0.285 kj/(mol K)] G 0 = 176 kj/mol 84.9 kj/mol G 0 = 91 kj/mol

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