1. Unit 11: Energy in Chemical Reactions 1

2. The Universe Is made up of the system and the surroundings Energy can be transferred between the system and the surroundings 2

3. Open, closed and isolated system 3

4. Chemical Reaction A chemical reaction is a chemical change. A new substance or substances is/are formed with different properties. Most reactions happen in aqueous or gas state because it is easy for the particles to touch. 4

5. In order to have a chemical reaction Involves breaking of existing bonds (requires energy) and formation of new bonds (releases energy) 5

6. Evidence of a chemical reaction 6

7. In a physical change: No new substances are formed No bonds are broken or formed 7

8. Intensive and Extensive properties Physical properties such as: Intensive (do not vary with amount) Color Odor Luster Malleability Ductility Conductivity Hardness Boiling Point/Melting Point Density (ratio of 2 extensive properties) Temperature Extensive (do vary with amount) Mass Weight Volume Length Watch this (3:51) Watch this 8

9. Stop and Think Watch this (1:35 BN) Name some intensive properties of matter that are different after a chemical change. Na + Cl  NaCl 9

10. Chemical Change or Physical Change? 10

11. parts of a chemical equation Reactants Products state of matter A chemical equation represents a chemical reaction. 11

12. Write an equation to represent this chemical change Nitrogen Chlorine 12

13. Collision Theory “Atoms, ions, or molecules must collide in order to react” 1.Correct reactant particles must collide. 2.Reactant particles must have the correct orientation when they collide. 3.Reacting substances must collide with sufficient energy to break bonds. 13

14. collision orientation 14

15. No Rxn No Rxn Collision Rebound Products Correct Orientation Sufficient Energy Incorrect Orientation Sufficient Energy Correct Orientation Insufficient Energy 15

16. Every collision does not result in a chemical reaction. Watch This (0:54) Watch This 16

17. Collision Summary Questions 1.What is an “effective collision”? 2.Do all collisions cause a reaction? 3.What can you do in the lab to make more effective collisions occur? 17

18. A potential energy diagram plots the change in potential energy that occurs during a chemical reaction. Watch This (1:16) Watch This 18

19. Potential Energy Diagrams Show the PE changes that occur as reactants become products during a chemical change http://www.kentchemistry.com/links/Kinetics/PEDiagrams.htm 19

20. PE Diagram Terms Activation Energy (E a ): minimum amount of energy input needed for a chemical reaction to occur. 20

21. Activation Energy http://staff.prairiesouth.ca/~chemistry/chem30/2_kinetics/kinetics3_2.htm E a is measured from the starting point of the reaction to the highest point on the curve Where would the activation energy be shown in each of these? Why? 21

22. PE Diagram Terms Activated Complex: a temporary, unstable arrangement of atoms after an effective collision; where old bonds are broken and new bonds are formed http://www.chem.ox.ac.uk/vrchemistry/chapter15/html/page29.htm 22

23. PE Diagram Terms Enthalpy Change (ΔH): Difference between the PE of the products and reactants in a chemical reaction. Energy can be absorbed or released. 23

24. Enthalpy Change (ΔH) http://staff.prairiesouth.ca/~chemistry/chem30/2_kinetics/kinetics3_2.htm Where would change in enthalpy be shown? Would ΔH be positive or negative? +∆H-∆H 24

25. Endothermic Reaction 25

26. Exothermic Reaction 26

27. What do exothermic and endothermic reactions look like? Watch this (4:35 Bozeman) Watch this 27

28. Enthalpy (H) A measure of heat content of a system. The heat lost or gained in a reaction. An extensive property Measured in units of Joules/mole (J/mol) or kJ/mol  H = change in heat  H = ∆H f ° (products) – ∆H f ° (reactants) This equation can be found on the bottom of your STAAR chart. “ ° “ refers to STP conditions 28

29. How to determine  H Method 1: By using the ∆H f ° values for each reactant and product Method 2: By analyzing a potential energy diagram Method 3: By looking at a chemical equation that includes the energy term Method 4: By making observations of a chemical reaction Method 5: Calorimetry Method 6: Hess’ Law Watch This (8:03 Bozeman) L stop at 4:46 Watch This 29

30. Method 1: Enthalpy calculations Use the table of ∆H f ° values and the equation  H = ∆H f ° (products) – ∆H f ° (reactants) to determine the enthalpy of the reaction. -  H indicates that a reaction is exothermic +  H indicates that a reaction is endothermic 30

31. Using enthalpies of formation, calculate the standard change in enthalpy for the thermite reaction: Fe 2 O 3 (s)+2Al(s)  Al 2 O 3 (s)+2Fe(s) ΔH f Fe 2 O 3 = (-826 kJ/mol) ΔH f Al 2 O 3 = (-1676 kJ/mol) Hints: i. all elements have a ΔH f = 0. ii. if the substance is multiplied by a coefficient, multiply ΔHº f by the same coefficient. 31

32. Practice problem Is the reaction described by the following equation an endothermic or exothermic reaction? What is the  H value? CH 4(g) + O 2(g)  CO 2(g) + 2H 2 O(l) Substance  H f  (kJ/mol) CH 4 -74.8 CO 2 (g)-393.5 H 2 O(l)-285.8 32

33. Method 2: Analyzing PE diagram Energy of reactants is bigger than energy of products Excess energy leaves to the surroundings. Exothermic (0:56) Exothermic 33

34. Method 2: Analyzing PE diagram The energy of the products is bigger than the energy of the reactants Energy is needed from the surroundings Endothermic (5:45) Endothermic 34

35. Method 3: chemical equation EXOTHERMIC 4Fe + 3O 2  2Fe 2 O 3 + 1625 kJ or 4Fe + 3O 2  2Fe 2 O 3 + energy Heat is written on the product side because it exits the system. 35

36. Method 3: chemical equation ENDOTHERMIC 27 kJ + NH 4 NO 3  NH 4 + + NO 3 - or energy + NH 4 NO 3  NH 4 + + NO 3 - Heat is written on the reactant side because it enters the system. 36

37. Method 4: observations of chemical reaction Feels hot -----EXOTHERMIC Feels cold -----ENDOTHERMIC 37

38. Method 5: Calorimetry (10:44) Calorimetry Calorimetry - the act of measuring the heat of chemical reactions or physical changes, or the science of making such measurements Calorimeter – the instrument used in calorimetry Q = mΔCT Q is the same as ΔH 38

39. Method 6: Hess Law Calculate the  H for the reaction: 2C(s) + H 2 (g)  C 2 H 2 (g) C 2 H 2 (g) + 5/2O 2 (g)  2CO 2 (g) + H 2 O(l)  H = -1299.6 kJ C(s) + O 2 (g)  CO 2 (g)  H = -393.5 kJ H 2 (g) + 1/2O 2 (g)  H 2 O(l)  H = -285.8 kJ 39

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42. Reaction Rate Rate at which reactant is converted to product depends on: – Concentration (more particles … more collisions… more reactions) – Temperature (particles moving faster.. more collisions.. more reactions) – Presence of a catalyst (lowers the activation energy, speeds up reaction rate) or inhibitor (raises the activation energy, slows down the reaction rate) – Nature of reactants (think back to “reactivity trends” for metals and nonmetals…K is more reactive than Li so K will react faster than Li) – Watch This (6:24 Bozeman) Watch This – Watch This (1:20) Watch This 42

43. Catalyst They increase the frequency of collisions or change the orientation of the molecules so more collisions are effective 43

44. Inhibitor Pathway with inhibitor Inhibitors raise the activation energy. Food preservatives are inhibitors. 44

45. Equilibrium When a reaction starts, the reactants are used up and products are made. Reactants Products After awhile, the products re-form to make reactants. Reactants Products Processes that proceed in both the forward and reverse direction are said to be reversible. Reactants Products Watch ThisWatch This (1:47) 45

46. Arrow Conventions Chemists commonly use two kinds of arrows in reactions to indicate the degree of completion of the reactions. A single arrow indicates all the reactant molecules are converted to product molecules at the end. A double arrow indicates the reaction stops when only some of the reactant molecules have been converted into products. –  in these notes 46

47. Reaction Dynamics When a reaction starts, the reactants are consumed and products are made. – The reactant concentrations decrease and the product concentrations increase. – As reactant concentration decreases, the forward reaction rate decreases. Eventually, the products can react to re-form some of the reactants, assuming the products are not allowed to escape. – As product concentration increases, the reverse reaction rate increases. Processes that proceed in both the forward and reverse direction are said to be reversible. reactants  products 47

48. Dynamic Equilibrium All reactions are reversible as long as they are in a closed container. Closed system= no interaction with the surroundings. As the forward reaction slows and the reverse reaction accelerates, eventually they reach the same rate. Dynamic equilibrium is the condition wherein the rates of the forward and reverse reactions are equal. Once the reaction reaches equilibrium, the concentrations of all the chemicals remain constant because the chemicals are being consumed and made at the same rate. Watch This (4:18) Watch This 48

49. H 2 (g) + I 2 (g)  2 HI(g) At time 0, there are only reactants in the mixture, so only the forward reaction can take place. [H 2 ] = 8, [I 2 ] = 8, [HI] = 0 49

50. H 2 (g) + I 2 (g)  2 HI(g) [H 2 ] = 6, [I 2 ] = 6, [HI] = 4 At time 16, there are both reactants and products in the mixture, so both the forward reaction and reverse reaction can take place. 50

51. H 2 (g) + I 2 (g)  2 HI(g) At time 32, there are now more products than reactants in the mixture, the forward reaction has slowed down as the reactants run out, and the reverse reaction accelerated. [H 2 ] = 4, [I 2 ] = 4, [HI] = 8 51

52. H 2 (g) + I 2 (g)  2 HI(g) At time 48, the amounts of products and reactants in the mixture haven’t changed; the forward and reverse reactions are proceeding at the same rate. It has reached equilibrium. 52

53. H 2 (g) + I 2 (g)  2 HI(g) As the concentration of product increases and the concentrations of reactants decrease, the rate of the forward reaction slows down, and the rate of the reverse reaction speeds up. 53

54. H 2 (g) + I 2 (g)  2 HI(g) At dynamic equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction. The concentrations of reactants and products no longer change. 54

55. Equilibrium  Equal The rates of the forward and reverse reactions are equal at equilibrium. But that does not mean the concentrations of reactants and products are equal. Some reactions reach equilibrium only after almost all the reactant molecules are consumed; we say the position of equilibrium favors the products. Other reactions reach equilibrium when only a small percentage of the reactant molecules are consumed; we say the position of equilibrium favors the reactants. 55