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Chapter 7

Copyright  The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 7. Energy, Rate, and Equilibrium . Denniston Topping Caret 5 th Edition. 7.1 Thermodynamics. Thermodynamics - the study of energy, work, and heat Applied to chemical change

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Chapter 7

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  1. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 7 Energy, Rate, and Equilibrium Denniston Topping Caret 5th Edition

  2. 7.1 Thermodynamics • Thermodynamics - the study of energy, work, and heat • Applied to chemical change • Calculate the quantity of heat obtained from combustions of one gallon of fuel oil • Applied to physical change • Determine the energy released by boiling water • The laws of thermodynamics help us to understand why some chemical reactions occur and others do not

  3. The Chemical Reaction and Energy Basic Concepts – from Kinetic Molecular Theory • Molecules and atoms in a reaction mixture are in constant, random motion • Molecules and atoms frequently collide with each other • Only some collisions, those with sufficient energy, will break bonds in molecules • When reactant bonds are broken, new bonds may be formed and products result 7.1 Thermodynamics

  4. Change in Energy and Surroundings • Absolute value for energy stored in a chemical system cannot be measured • Can measure the change in energy during these chemical changes • System - contains the process under study • Surroundings- the rest of the universe 7.1 Thermodynamics

  5. Changes in the System • Energy can be lost from the system to the surroundings • Energy may be gained by the system at the expense of the surroundings • This energy change is usually in the form of heat • This change can be measured 7.1 Thermodynamics

  6. Law of Conservation of Energy • The first law of thermodynamics - energy of the universe is constant • This law is also called the Law of Conservation of Energy • Where does the reaction energy come from that is released and where does the energy go when it is absorbed? 7.1 Thermodynamics

  7. Changes in Chemical Energy A-B + C-D  A-D + C-B • Consider the reaction converting AB and CD to AD and CB • Each chemical bond is stored chemical energy • If a reaction will occur: • Bonds must break • Breaking bonds requires energy 7.1 Thermodynamics

  8. Exothermic Reactions If the energy required to break the bonds is less than the energy released when the bonds are formed, there is a net release of energy • This is called an exothermic reaction • Energy is a product in this reaction 7.1 Thermodynamics A-B + C-D  A-D + C-B These bonds must be broken in the reaction, requiring energy These bonds are formed, releasing energy

  9. Endothermic Reactions If the energy required to break the bonds is larger than the energy released when the bonds are formed, there will need to be an external supply of energy • This is called an endothermic reaction 7.1 Thermodynamics A-B + C-D  A-D + C-B These bonds must be broken in the reaction, requiring energy These bonds are formed, releasing energy

  10. Exothermic Reaction Combustion CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + 211 kcal 7.1 Thermodynamics Endothermic Reaction Decomposition 22 kcal + 2NH3(g)  N2(g) + 3H2(g)

  11. Enthalpy • Enthalpy - represents heat energy • Change in Enthalpy (DHo) - energy difference between the products and reactants of a chemical reaction • Energy released, exothermic reaction, enthalpy change is negative • In the combustion of CH4, DHo = –211 kcal • Energy absorbed, endothermic, enthalpy change is positive • In the decomposition of NH3, DHo = +22 kcal 7.1 Thermodynamics

  12. Spontaneous and Nonspontaneous Reactions • Spontaneous reaction - occurs without any external energy input • Most, but not all, exothermic reactions are spontaneous • Thermodynamics is used to help predict if a reaction will occur • Another factor is needed, Entropy 7.1 Thermodynamics

  13. Spontaneous and Nonspontaneous Reactions 7.1 Thermodynamics DS o is positive DSo is negative

  14. Spontaneous and Nonspontaneous Reactions Are the following processes exothermic or endothermic? • Fuel oil is burned in a furnace • C6H12O6(s) 2C2H5OH(l) + 2CO2(g) DH = –16 kcal • N2O5(g) + H2O(l) 2HNO3(l) + 18.3 kcal 7.1 Thermodynamics

  15. Entropy • The second law of thermodynamics - the universe spontaneously tends toward increasing disorder or randomness • Entropy (So) – a measure of the randomness of a chemical system • High entropy – highly disordered system, the absence of a regular, repeating pattern • Low entropy – well organized system such as a crystalline structure • No such thing as negative entropy 7.1 Thermodynamics

  16. Entropy of Reactions DSo of a reaction = So(products) - So(reactants) • A positive DSo means an increase in disorder for the reaction • A negative DSo means a decrease in disorder for the reaction 7.1 Thermodynamics

  17. Processes Having Positive Entropy Phase change Melting Vaporization Dissolution 7.1 Thermodynamics All of these processes have a positive DSo

  18. Entropy and Reaction Spontaneity • If exothermic and positive DSo… SPONTANEOUS • If endothermic and negative DSo… NONSPONTANEOUS • For any other situations, it depends on the relative size of DHo and DSo 7.1 Thermodynamics

  19. Greatest Entropy • Which substance has the greatest entropy? • He(g) or Na(s) • H2O(l) or H2O(g) 7.1 Thermodynamics

  20. Free Energy • Free energy (DGo) - represents the combined contribution of the enthalpy and entropy values for a chemical reaction • Free energy predicts spontaneity of chemical reactions • Negative DGo…Always Spontaneous • Positive DGo…Never Spontaneous 7.1 Thermodynamics DGo = DHo - TDSo T in Kelvin

  21. Free Energy and Reaction Spontaneity • Need to know both DH and DS to predict the sign of DG, making a statement on reaction spontaneity • Temperature also may determine direction of spontaneity • DH+, DS- : DG always +, regardless of T • DH-, DS+ : DG always -, regardless of T • DH+, DS+ : DG sign depends on T • DH-, DS- : DG sign depends on T 7.1 Thermodynamics

  22. 7.2 Experimental Determination of Energy Change in Reactions • Calorimetry - the measurement of heat energy changes in a chemical reaction • Calorimeter - device which measures heat changes in calories • The change in temperature is used to measure the loss or gain of heat

  23. Heat Energy in Reactions • Change in temperature of a solution, caused by a chemical reaction, can be used to calculate the gain or loss of heat energy for the reaction • Exothermic reaction – heat released is absorbed • Endothermic reaction – reactants absorb heat from the solution • Specific heat (SH) - the number of calories of heat needed to raise the temperature of 1 g of a substance 1oC 7.2 Determination of Energy Change in Reactions

  24. Heat Energy in Reactions • Specific heat of the solution along with the total number of grams of solution and the temperature change, permits calculation of heat released or absorbed during the reaction • SH for water is 1.0 cal/goC • To determine heat released or absorbed, need: • specific heat • total number of grams of solution • temperature change (increase or decrease) 7.2 Determination of Energy Change in Reactions

  25. Calculation of Heat Energy in Reactions • Q is the product • ms is the mass of solution in the calorimeter • DTs is the change in temperature of the solution from initial to final state • SHs is the specific heat of the solution • Calculate with this equation • Units are: calories = gram x ºC x calories/gram - ºC 7.2 Determination of Energy Change in Reactions

  26. Calculating Energy Involved in Calorimeter Reactions If 0.10 mol HCl is mixed with 0.10 mol KOH in a “coffee cup” calorimeter, the temperature of 1.50 x 102 g of the solution increases from 25.0oC to 29.4oC. If the specific heat of the solution is 1.00 cal/goC, calculate the quantity of energy evolved in the reaction. DTs = 29.4oC - 25.0oC = 4.4oC Q = msx DTs x SHs = 1.50 x 102 g solution x 4.4oC x 1.00 cal/goC = 6.6 x 102 cal 7.2 Determination of Energy Change in Reactions

  27. Calculating Energy Involved in Calorimeter Reactions Is the reaction endothermic or exothermic? • 0.66 kcal of heat energy was released to the surroundings—the solution • The reaction is exothermic What would be the energy evolved for each mole of HCl reacted? • 0.10 mol HCl used in the original reaction • [6.6 x 102 cal / 0.10 mol HCl] x 10 = 6.6 kcal 7.2 Determination of Energy Change in Reactions

  28. Bomb Calorimeter and Measurement of Calories in Foods Nutritional Calorie(large “C” Calorie) = • 1 kilocalorie (1kcal) • 1000 calories • The fuel value of food • Bomb calorimeter is used to measure nutritional Calories 7.2 Determination of Energy Change in Reactions

  29. Calculating the Fuel Value of Foods 1 g of glucose was burned in a bomb calorimeter. 1.25 x 103 g H2O was warmed from 24.5oC to 31.5oC. Calculate the fuel value of the glucose (in Kcal/g). • DTs = 31.5oC - 24.5oC = 6.1oC Surroundings of calorimeter is water with specific heat capacity = 1.00 cal/g H2OoC Fuel Value = = 1.25 x 103 g H2O x 6.1oC x 1.00 cal/g H2OoC = 7.6 x 103 cal 7.6 x 103 cal x 1 Calorie / 103 cal = 7.6 nutritional Calories 7.2 Determination of Energy Change in Reactions

  30. 7.3 Kinetics • Thermodynamics determines if a reaction will occur spontaneously, but tells us nothing about the amount of time the reaction will take • Kinetics - the study of the rate (or speed) of chemical reactions • Also supplies an indication of the mechanism – step-by-step description of how reactants become products

  31. Kinetic Information • Kinetic information represents changes over time, seen here: • Disappearance of reactant, A • Appearance of product, B 7.3 Kinetics

  32. Alternative Presentation of Kinetic Data Rather than the graph shown before, this figure demonstrates the change from purple reactant to green product over time from the molecular perspective 7.3 Kinetics

  33. Kinetic Data Assessed by Color Change • Change in color over time can be used to monitor the progress of a chemical reaction • The rate of color change can aid in calculating the rate of the chemical reaction 7.3 Kinetics

  34. The Chemical Reaction CH4(g) + 2O2(g)  CO2(g) +2H2O(g) + 211 kcal • C-H and O=O bonds must be broken • C=O and O-H bonds must be formed • Energy is required to break the bonds • This energy comes from the collision of the molecules • If sufficient energy available, bonds break and atoms recombine in a lower energy arrangement • Effective collision is one that produces product molecules • Leads to a chemical reaction 7.3 Kinetics

  35. Activation Energy and the Activated Complex • Activation energy - the minimum amount of energy required to initiate a chemical reaction • Picture a chemical reaction in terms of changes in potential energy occurring during the reaction • Activated complex - an extremely unstable, short-lived intermediate complex • Formation of this activated complex requires energy (Ea) to overcome the energy barrier to start the reaction 7.3 Kinetics

  36. Activation Energy and the Activated Complex • Reaction proceeds from reactants to products via the activated complex • Activated complex - can’t be isolated from the reaction mixture • Activation energy (Ea) is the difference between the energy of the reactants and that of the activated complex 7.3 Kinetics • To be an Exothermic reaction requires a net release of energy (DHo)

  37. Activation Energy in the Endothermic Reaction • This figure diagrams an endothermic reaction • Reaction takes place slowly due to the large activation energy required • The energy of the products is greater than that of the reactants 7.3 Kinetics

  38. Factors That Affect Reaction Rate • Structure of the reacting species • Concentration of reactants • Temperature of reactants • Physical state of reactants • Presence of a catalyst 7.3 Kinetics

  39. Structure of Reacting Species • Oppositely charged species react more rapidly • Dissociated ions in solution whose bonds are already broken have a very low activation energy • Ions with the same charge do not react • Bond strength plays a role • Covalent molecules bonds must be broken with the activation energy before new bonds can be formed • Magnitude of the activation energy is related to bond strength • Size and shape influence the rate • Large molecules may obstruct the reactive part of the molecule • Only molecular collisions with correct orientation lead to product formation 7.3 Kinetics

  40. The Concentration of Reactants • Rate is related to the concentration of one or more of the reacting substances • Rate will generally increase as concentration increases • Higher concentration means more reactant molecules per unit volume • More reactant molecules means more collisions per unit time 7.3 Kinetics

  41. The Temperature of Reactants • Rate increases as the temperature increases • Increased temperature relates directly to increased average kinetic energy • Greater kinetic energy increases the speed of particles • Faster particles increases likelihood of collision • Higher kinetic energy means a higher percentage of these collisions will result in product formation 7.3 Kinetics

  42. The Physical State of Reactants Reactions occur when reactants can collide frequently with sufficient energy to react • Solid state: atoms, ions, compounds are close together but restricted in motion • Gaseous state: particles are free to move but often are far apart causing collisions to be relatively infrequent • Liquid state: particles are free to move and are in close proximity • Reactions to be fastest in the liquid state and slowest in the solid state • Liquid > Gas > Solid 7.3 Kinetics

  43. The Presence of a Catalyst • Catalyst- a substance that increases the reaction rate • Undergoes no net change • Does not alter the final product of the reaction • Interacts with the reactants to create an alternative pathway for product production 7.3 Kinetics

  44. Use of a Solid Phase Catalyst N2+3H2 2NH3 7.3 Kinetics Haber Process is a synthesis of ammonia facilitated by a solid phase catalyst • Diatomic gases bind to the surface • Bonds are weakened • Dissociation of diatomic gases and reformation of NH3 • Newly formed NH3 leaves the solid surface with the catalyst unchanged

  45. Mathematical Representation of Reaction Rate • Consider a decomposition reaction with the following balanced chemical equation: • When heated N2O5 decomposes to 2 products – NO2 and O2 • When holding all factors constant, except concentration, rate of reaction is proportional to the concentration 7.3 Kinetics

  46. Mathematical Representation of Reaction Rate • Reaction rate is proportional to reactant concentration – • Concentration of N2O5 is denoted as [N2O5] • Replace proportionality symbol with (=) and proportionality constant k • k is called the rate constant 7.3 Kinetics

  47. Rate Equation • For a reaction Aproductswe writethe equation: rate = k[A]n • This is called the rate equation (or rate law) • The exponent n is the order of the reaction • If n=1, first order • If n=2, second order • n must be determined experimentally • This exponent is not the same as the coefficient of the reactant in the balanced equation 7.3 Kinetics

  48. Rate Equation • For the equation A + B  products the rate equation is: • rate = k[A]n[B]m • What would be the general form of the rate equation for the reaction: CH4+2O2CO2+2H2O • Rate = k[CH4]n[O2]m • Knowing the rate equation and the order helps industrial chemists determine the optimum conditions for preparing a product 8.3 Kinetics

  49. Writing Rate Equations • Write the form of the rate equation for the oxidation of ethanol (C2H5OH) • The reaction has been experimentally determined to be first order in ethanol and third order in oxygen • Rate expression involves only the reactants • Concentrations: [C2H5OH][O2] • Raise each to exponent corresponding to its order rate = k [C2H5OH][O2]3 • Remember that 1 as an exponent is understood and NOT written • Knowing the rate equation and the order helps industrial chemists determine the optimum conditions for preparing a product 8.3 Kinetics

  50. 7.4 Equilibrium Rate and Reversibility of Reactions • Equilibrium reactions - chemical reactions that do not go to completion • Completion – all reactants have been converted to products • Equilibrium reactions are also called incompletereactions • Seen with both physical and chemical processes • After no further observable change, measurable quantities of reactants and products remain

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