1. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Thermochemistry Virtually every chemical reaction is accompanied by a change in energy. Chemical reactions usually either absorb or release energy as heat. Thermochemistry is the study of the transfers of energy as heat that accompany chemical reactions and physical changes. Section 1 Thermochemistry Chapter 16

2. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Heat and Temperature, continued Temperature is 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. K = 273.15 + °C Section 1 Thermochemistry Chapter 16

3. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Heat and Temperature, continued The amount of energy transferred as heat is usually measured in joules. Heat can be thought of as 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. (hot coffee in a cold room) Section 1 Thermochemistry Chapter 16

4. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Specific Heat (c p ) -The amount of energy required to raise the temperature of 1 gram of a substance by 1°C (or 1 K) -For water (liquid), the specific heat (c p ) is 4.18 J/ (g * K)

5. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Energy gained or lost (q) -Energy Gained or Lost (q) with a change in temperature -q = c p * m * ∆T -A 4.0 g sample of glass was heated from 274K to 314K and was found to have absorbed 32J as heat. What is the specific heat of this type of glass? -32 J = c p * 4.0g * (314K – 274K) -c p = 0.2 J / (g*K)

6. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy of Reaction The energy absorbed as heat during a chemical reaction at constant pressure is represented by ∆H. H is the symbol for a quantity called enthalpy. Only changes in enthalpy can be measured. ∆H is read as “change in enthalpy.” An enthalpy change is the amount of energy absorbed by a system as heat during a process at constant pressure. Section 1 Thermochemistry Chapter 16

7. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy of Reaction, continued Enthalpy change is always the difference between the enthalpies of products and reactants. ∆H = H products – H reactants Section 1 Thermochemistry Chapter 16

8. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Exothermic Chemical Reactions System releases energy Energy of the products < energy of the reactants ∆ H is negative example: 2H 2 (g) + O 2 (g)  2H 2 O(g) + 483.6 kJ 2H 2 (g) + O 2 (g)  2H 2 O(g) ∆ H = –483.6 kJ Products are more stable than the reactants because they have lower energy.

9. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy of Reaction, continued In an exothermic reaction, energy is evolved, or given off, during the reaction; ∆ H is negative. Section 1 Thermochemistry Chapter 16

10. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Endothermic Chemical Reactions System gains energy Energy of the products > energy of the reactants ∆H is positive example: 2H 2 O(g) + 483.6 kJ  2H 2 (g) + O 2 (g) ∆ H = +483.6 kJ Reactants are more stable because they have lower energy Bank account example Section 1 Thermochemistry Chapter 16

11. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy of Reaction, continued In an endothermic reaction, energy is absorbed; in this case, ∆ H is designated as positive. Section 1 Thermochemistry Chapter 16

12. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy of Reaction, continued The physical states of reactants and products must always be included in thermochemical equations, because the states of reactants and products influence the overall amount of energy as heat gained or lost. Section 1 Thermochemistry Chapter 16

13. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 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. Enthalpies of formation are given for a standard temperature and pressure so that comparisons between compounds are meaningful. To signify standard states, a 0 sign is added to the enthalpy symbol, and the subscript f indicates a standard enthalpy of formation: Section 1 Thermochemistry Chapter 16

14. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy of Formation, continued Some standard enthalpies of formation are given in the appendix of your book. 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. The thermochemical equation to accompany an enthalpy of formation shows the formation of one mole of the compound from its elements in their standard states. Section 1 Thermochemistry Chapter 16

15. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Stability and Enthalpy of Formation Compounds with a large negative enthalpy of formation are very stable. Section 1 Thermochemistry Chapter 16 example: the of carbon dioxide is –393.5 kJ per mol of gas produced. Elements in their standard states are defined as having = 0. This indicates that carbon dioxide is more stable than the elements from which it was formed.

16. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Stability and Enthalpy of Formation, continued Compounds with positive values of enthalpies of formation are typically unstable. Section 1 Thermochemistry Chapter 16 example: hydrogen iodide, HI, has a of +26.5 kJ/mol. It decomposes at room temperature into violet iodine vapor, I 2, and hydrogen, H 2.

17. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Calculating Enthalpies of Reaction, continued Any enthalpy of reaction may be calculated using enthalpies of formation for all the substances in the reaction of interest Mathematically, the overall equation for enthalpy change will be in the form of the following equation: ∆H = H products – H reactants ∆ H 0 = sum of [( of products)  (mol of products)] – sum of [( of reactants)  (mol of reactants)] Section 1 Thermochemistry Chapter 16

18. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 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) + O 2 (g)  NO 2 (g) Use the enthalpy-of-formation data in the appendix. Calculating Enthalpies of Reaction, continued Section 1 Thermochemistry Chapter 16

19. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Sample Problem B Solution Section 1 Thermochemistry Chapter 16 Calculating Enthalpies of Reaction, continued Given: Unknown:

20. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Enthalpy and Reaction Tendency Most reactions in nature are exothermic  lower energy state Some endothermic reactions do occur spontaneously. Something other than enthalpy change can help determine whether a reaction will occur. Section 2 Driving Force of Reactions Chapter 16

21. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Entropy and Reaction Tendency Melting = endothermic process that occurs naturally An ice cube melts spontaneously at room temperature as energy is transferred from the warm air to the ice. The well-ordered arrangement of water molecules in 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 a negative enthalpy change does so with an increase in entropy. Section 2 Driving Force of Reactions Chapter 16

22. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu 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) Section 2 Driving Force of Reactions Chapter 16 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.

23. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Entropy and Reaction Tendency, continued There is a tendency in nature 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, can be defined in a simple qualitative way as a measure of the degree of randomness of the particles, such as molecules, in a system. Section 2 Driving Force of Reactions Chapter 16

24. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Standard Entropy Changes for Some Reactions Section 2 Driving Force of Reactions Chapter 16

25. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Entropy and Reaction Tendency, continued To understand the concept of entropy, consider the comparison between particles in solids, liquids, and gases. In a solid, the particles are in fixed positions, and we can easily determine the locations of the particles. In a liquid, the particles are very close together, but they can move around. Locating an individual particle is more difficult. The system is more random, and the entropy is higher. Section 2 Driving Force of Reactions Chapter 16

26. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Entropy and Reaction Tendency, continued In a gas, the particles are moving rapidly and are far apart. Locating an individual particle is much more difficult, and the system is much more random. The entropy is even higher. Positive entropy is favored whereas negative enthalpy is favored. Section 2 Driving Force of Reactions Chapter 16

27. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemists have found that chemical reactions occur at widely differing rates. The speed of a chemical reaction depends on the energy pathway that a reaction follows and the changes that take place on the molecular level when substances interact. Chapter 17 Section 1 The Reaction Process

28. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Collision Theory In order for reactions to occur between substances, their particles must collide. The set of assumptions regarding collisions and reactions is known as collision theory. Reactant molecules must collide with a favorable orientation and with enough energy to merge the valence electrons and disrupt the bonds of the molecules to form to the products. Chapter 17 Section 1 The Reaction Process

29. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Particle Collisions Chapter 17 Section 1 The Reaction Process

30. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Collision Theory, continued A chemical reaction produces new bonds which are formed between specific atoms in the colliding molecules. Unless the collision brings the correct atoms close together and in the proper orientation, the molecules will not react. Chapter 17 Section 1 The Reaction Process

31. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Possible Collision Orientations for the Reaction of H 2 and I 2 Chapter 17 Section 1 The Reaction Process

32. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Collision Theory, continued Collision theory provides two reasons why a collision between reactant molecules may fail to produce a new chemical species: the collision is not energetic enough to supply the required energy the colliding molecules are not oriented in a way that enables them to react with each other Chapter 17 Section 1 The Reaction Process

33. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Activation Energy http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/activa2.swf Chapter 17 Section 1 The Reaction Process An initial input of energy is needed to overcome the repulsion forces that occur between reactant molecules when they are brought very close together. Bonds must be broken in order for new bonds to be formed. Bond breaking is an endothermic process (energy must be added) Molecules must possess kinetic energy equal to or greater than activation energy This initial energy input activates the reaction.

34. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Activation Energy, continued Activation energy (E a ) is the minimum energy required to transform the reactants into an activated complex. Chapter 17 Section 1 The Reaction Process

35. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Activation Energies Chapter 17 Section 1 The Reaction Process

36. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Activation Energy Differences in Exothermic and Endothermic Reactions Chapter 17 Section 1 The Reaction Process

37. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The change in concentration of reactants per unit time as a reaction proceeds is called the reaction rate. The area of chemistry that is concerned with reaction rates and reaction mechanisms is called chemical kinetics. Chapter 17 Section 2 Reaction Rate

38. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu For reactions other than simple decompositions to occur, particles must come into contact in a favorable orientation and with enough energy for activation. The rate of a reaction depends on the collision frequency of the reactants and on the collision efficiency (orientation). The nature of the reactants, the surface area, the temperature, the concentration, and the presence of a catalyst are factors that influence the rate of a chemical reaction. Rate-Influencing Factors Chapter 17 Section 2 Reaction Rate

39. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Nature of Reactants The rate of reaction depends on the particular reactants and bonds involved. Surface Area When reactants are in different phases (heterogeneous), the reaction rate depends on the area of contact of the reaction substances. ↑ surface area = ↑ rate of heterogeneous reactions Rate-Influencing Factors, continued Chapter 17 Section 2 Reaction Rate

40. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Temperature ↑ T = ↑ average KE and ↑ number of effective collisions = ↑ reaction rate At higher temperatures, more particles possess enough energy to form the activated complex when collisions occur. A rise in temperature produces an increase in collision energy as well as in collision frequency. Rate-Influencing Factors, continued Chapter 17 Section 2 Reaction Rate

41. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Concentration The reaction rate depends on the concentration of the reacting species. example: A substance that oxidizes in air oxidizes more vigorously in pure oxygen. Rate-Influencing Factors, continued Chapter 17 Section 2 Reaction Rate

42. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Concentration Vs. Possible Collisions Chapter 17 Section 2 Reaction Rate

43. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Concentration, continued ↑ [reactants] = ↑ reaction rate (in general) The actual effect of concentration changes on reaction rate must be determined experimentally. Rate-Influencing Factors, continued Chapter 17 Section 2 Reaction Rate

44. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Presence of Catalysts Sometimes reaction rates can be increased dramatically by the presence of a catalyst. A catalyst is a substance that changes the rate of a chemical reaction without itself being permanently consumed. The action of a catalyst is called catalysis. (catalyst = yeast in our Elephant Toothpaste experiment) Rate-Influencing Factors, continued Chapter 17 Section 2 Reaction Rate

45. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Presence of Catalysts, continued A catalyst provides an alternative energy pathway or reaction mechanism in which the potential-energy barrier between reactants and products is lowered. The catalyst may be effective in forming an alternative activated complex that requires a lower activation energy. Catalysts do not appear among the final products of reactions they accelerate. Rate-Influencing Factors, continued Chapter 17 Section 2 Reaction Rate

46. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The relationship between the rate of a reaction and the concentration of one reactant is determined experimentally. A series of experiments reveals how the concentration of each reactant affects the reaction rate. An equation that relates reaction rate and concentrations of reactants is called the rate law for the reaction. Rate Laws for Reactions- OPTIONAL Chapter 17 Section 2 Reaction Rate It is applicable for a specific reaction at a given temperature.

47. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The general form for the rate law is given by the following equation: R = k[A] n [B] m The rate law is applicable for a specific reaction at a given set of conditions and must be determined from experimental data. The power to which a reactant concentration is raised is called the order in that reactant. The value of n is said to be the order of the reaction with respect to [A], so the reaction is said to be “nth order in A.” Rate Laws for Reactions, continued Using the Rate Law Chapter 17 Section 2 Reaction Rate

48. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The sum of all of the reactant orders is called the order of the reaction, or overall order. Rate Laws for Reactions, continued Using the Rate Law, continued Chapter 17 Section 2 Reaction Rate NO 2 (g) + CO(g) NO(g) + CO 2 (g) R = k[NO 2 ] 2 second order in NO 2 zero order in CO second order overall The orders in the rate law may or may not match the coefficients in the balanced equation.

49. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The specific rate constant (k) is the proportionality constant relating the rate of the reaction to the concentrations of reactants. 1.k is determined experimentally 2.Value of k specific to particular reaction 3.k changes with temperature or in the presence of a catalyst Rate Laws for Reactions, continued Specific Rate Constant Chapter 17 Section 2 Reaction Rate