2. Chemical Kinetics and Collision Theory Aim KE1 How do chemical reactions actually happen?

3. Collision Theory - A chemistry theory that explains how chemical reactions occur through effective collision of particles –Effective collision - one in which the colliding particles approach each other at the proper angle and with the proper amount of energy The greater the rate of effective collisions, the greater the reaction rate (speed of reaction)

4. An Analogy - Collision theory is similar to bowling: –You throw a ball very weakly –do you knock all the pins down? Why? –You throw the ball at the 10 pin (back right corner) –do you knock down all the pins? Why?

5. Reaction mechanisms Chemical reactions where only two atoms/molecules collide to make a product are rare Most reactions go through a series of steps called the reaction mechanism Example of steps in a reaction: the formation of ammonia (a process called the Haber Process) 3H 2 + N 2  2NH 3 + 92.2 kJ There are three basic steps in the reaction: –Step 1:three H 2 are broken up into six H atoms –Step 2:one N 2 is broken up into two N atoms –Step 3:each N atom collides with three H atoms to make NH 3 (three separate collisions)

6. Rates of Reaction Mechanisms Each step requires collisions Some happen quickly, some happen more slowly In the Haber Process, 3H 2 + N 2  2NH 3 + 92.2 kJ Step 1:the break up of H 2 is FAST Step 2:the break up of N 2 is FAST Step 3:the formation of NH 3 from the H and N is SLOW The slowest step of the reaction mechanism is called the rate determining step It determines how fast the overall reaction will take to happen

7. Transition State Theory As a chemical reaction progresses »Atoms rearrange themselves forming intermediate products (not quite reactants, but not quite products) »Example – In the reaction of hydrogen with oxygen to form water at left These intermediate products are called transition state complexes or activated complexes

8. Transition State Theory Activated complexes –exist for only brief periods of time while the atoms rearrange themselves –have high energy due to their formation by high energy collisions –they are unstable but need to form in order to make the final product(s) –the energy needed to form the activated complex is called the activation energy

9. Activation energy –gives the chemical system enough energy to reach the activated complex –The reaction can than continue to completion, with new products formed –Activation energy varies with the nature of substances in the reaction The type of reaction (endothermic vs exothermic reactions) If reactions didn’t need activation energy, they would all simply happen But this assumes we don’t have an activated complex

10. Transition State Theory –The energy changes in a chemical reaction can be shown in a graph called a potential energy diagram –the high energy product is called an activated complex or a transition state complex –the energy needed to form the activated complex is the activation energy –The reaction pathway represents time or the progress of the reaction –The potential energy shows the amount of energy the chemicals have at different points in the reaction

11. Reading a Potential Energy Diagram The PE Diagram shows the amount of energy that changes in a chemical reaction PE reactants = energy the reactants possess already PE products = the energy the products possess after the reaction occurs E a = the activation energy needed to reach the activated complex  H = the heat of reaction – the heat change between the products and the reactants

12. Rates of Reaction Aim KE 2 What is the difference between exothermic and endothermic reactions? Mr. Foley, may I be excused? My brain is full.

13. Role of Energy in Reactions Enthalpy is the difference between the potential energy (PE) of the products and the reactants  H = H products – H reactants heat of reaction or enthalpy (  H) –The heat energy (enthalpy) released or absorbed during a chemical reaction –In an exothermic reaction heat is released  H is negative ( -  H ) –In an endothermic reaction heat is absorbed  H is positive ( +  H )

14. 53.0 kJ + H 2 + I 2  2 HI H 2 + I 2 2 HI Potential Energy Reaction Pathway Activation Energy  H = + 53 kJ In an endothermic reaction –heat is added to reactants in order to make the products –for the heat of reaction for an endothermic reaction is +  H (heat of reaction) –energy was added to the reactant side of the reaction

15. Example 2 H 2 + O 2  2 H 2 O + 484 kJ  H 2 + O 2 2 H 2 O Potential Energy Reaction Pathway  H = - 484 kJ Activation Energy In an exothermic reaction –heat is removed from the reactants in order to make the products –for the heat of reaction for an exothermic reaction is -  H –energy was removed and comes out as a product

16. Potential energy diagram #1 Match each letter with the appropriate label using the word bank below: 1. PE of Reactants = ____ 3. Activation Energy = _____ 2. PE of Products = ____ 4. Heat of reaction = _____ 5. Mark an X where the activated complex is located. 6. What type of reaction is this, endothermic or exothermic? ________________ 7. How do you know? _________________ _________________ 8. Is the heat of reaction positive or negative? __________________

17. Potential energy diagram #2 Match each letter with the appropriate label using the word bank below: 1.PE of Reactants = ____ 3. Activation Energy = _____ 2.PE of Products = ____ 4. Heat of reaction = _____ 5. Mark an X where the activated complex is located. 6. What type of reaction is this, endothermic or exothermic? ________________ 7. How do you know? _________________ _________________ 8. Is the heat of reaction positive or negative? __________________

18. Table I – Heats of Rxn Various reactions are presented in this chart All are at standard pressure and room temperature (298 K) The  H represents the energy absorbed (+) or released in the reaction (-) This energy represents the total energy of all the moles formed in each reaction

19. Examples from Table I Combustion of methane CH 4(g) + 2O 2(q)  CO 2(g) + 2H 2 O (g) + 890.4 kJ  H = - 890.4 kJ; an exothermic reaction Formation of water vapor 2H 2(g) + O 2(q)  2H 2 O (g) + 483.4 kJ  H = - 483.4 kJ; an exothermic reaction Synthesis of nitrogen dioxide 66.4 kJ + N 2(g) + 2O 2(g)  2 NO (g)  H = + 66.4 kJ; an endothermic reaction Dissolving (ionization) of a salt (cold pack salt) H 2 O 14.78 kJ + NH 4 Cl (s) ---> NH 4 + (aq) + Cl - (aq)  H = +14.78 kJ; an endothermic reaction

20. Using Table I to fill in the chart below: ReactionExothermic or Endothermic?  H value C 3 H 8(g) +5O 2(g)  3CO 2(g) +4H 2 O (l) Exothermic-890.4 kJ N 2(g) + 3H 2(g)  2NH 3(g) Exothermic-91.8 kJ H 2 O (g) LiBr (s)  Li + (aq) + Br - (aq) Exothermic- 48.83 kJ The formation of liquid water from hydrogen and oxygen gas Exothermic-571.6 kJ The dissolving of sodium chloride in water to form sodium ions and chlorine ions Endothermic+3.81

22. Rates of Reaction Aim KE 3 What speeds up chemical reactions?

23. Factors affecting Rates of Chemical Reactions Just like when we spoke about the rate of dissolving and the effects on solubility Various factors affect the speed or rate of a chemical reaction –Concentration – the number of particles in a given sample –Pressure – affects only gases –Nature of Particles – how large they are and how well they react –Surface Area – the amount of molecular surface exposed between the reactants –Temperature – the amount of kinetic energy

24. Increasing Chemical Reaction Rates 1 - Increase the concentration of the reactants –Increasing concentration could mean adding more reactants (increasing molarity) –or increasing the density by making the volume of the reaction vessel smaller –More collisions = increasing reaction rate

25. Increasing Rates of Chemical Reactions 2 - Increase the pressure of gaseous reactants –Squeezing together gases increases concentration –Increasing pressure forces reactants closer together - but only for gases

26. Increasing Rates of Chemical Reactions 3 - Nature of Reactants –Chemical reactions occur by breaking and rearranging existing bonds. –The less electrons that need to be rearranged, the faster the reaction is. –As a result, reactions between ionic substances in aqueous solution, such as double replacement reactions, are rapid at room temperature –Ex: KI (aq) + Pb(NO 3 ) 2(aq)  KNO 3 (aq) + PbI 2(solid) –On the other hand, reactions in which covalent bonds are broken, such as the decomposition of hydrogen peroxide, occur slowly at room temperature.

27. Increasing Rates of Chemical Reactions 4 - Increasing surface area –More particles available to react with each other

28. Increasing Rates of Chemical Reactions 5 - Increasing temperature –Increases average kinetic energy of particles –Increases energy and number of collisions 75 o C R 50 o C A 25 o C T E Time

29. Increasing Rates of Chemical Reactions 6 – Catalysts Catalysts speed up reactions without being permanently altered Enzymes in biological systems are organic (carbon based) catalysts They change the pathway the reaction mechanism follows so less activation energy is required to get to the activated complex or transition state

30. Equilibrium Aim KE 4d – What is equilibrium in chemistry?

31. Equilibrium –A balance between two opposing forces or objects –Tends to be static (doesn’t change once reached) –Ex – gravity vs your leg bones or a building’s supports In chemistry –Dynamic equilibrium occurs –So that the balancing forces are always adjusting themselves –Chemical reactions are reversible –They can occur in either direction, so one opposing reaction is exothermic, the other is endothermic

32. Under conditions of dynamic equilibrium –the rates of the forward and reverse reactions are equal –A balance point occurs as the two reactions compete –But the amounts of materials on each side of the reaction may NOT be equal –and the point of equilibrium can change with changing conditions

33. Types of Equilibrium - Phase Equilibrium –We know phase changes are reversible –Ice in a container at 0 o C - S  L –Rate of melting equals rate of freezing –Water in a container at 25 o C - L  G –Rate of evaporation equals rate of condensation –There is a balance or equilibrium point at each of these temperatures –What happens if the temp changes?

34. Types of Equilibrium - Phase Equilibrium –What happens if the temperature changes? –If we write the phase equilibrium as an equation: H 2 O (l)  H 2 O (g) At a temperature of 25 o C, there might be 100 mL of liquid water and 5 mL of water vapor As the temperature increases, more liquid will change to gas and the equilibrium point shifts to the right At first, evaporation will occur faster than condensation and more water vapor will form; eventually, the rates will equal out as more vapor cools to from liquid A new balance point occurs where there 95 mL of liquid water and 10 mL of water vapor

35. Types of Equilibrium – Solution Equilibrium –When a solid dissolves in a saturated solution –The rate of dissolving = the rate of crystallization Examples: –Mr. Foley’s Duncan Donuts Light and Sweet coffee –C 12 H 22 O 11 (s)  C 6 H 12 O 6 (aq) where sugar is dissolving = –C 6 H 12 O 6 (aq)  C 12 H 22 O 11 (s) where sugar is precipitating and settling on the bottom = –The equilibrium equation is C 12 H 22 O 11(s)  C 6 H 12 O 6(aq)

36. Types of Equilibrium – Solution Equilibrium –Remember – a saturated solution holds the maximum amount of solute at that temperature –What happens if we lower the temp of Mr. Foley’s coffee? –The rate of dissolving < the rate of crystallization (at first) –Sugar precipitates faster than the sugar dissolves –A new saturation point is reached as the excess sugar at the new temp falls out of solution –But a new equilibrium occurs with more sugar at the bottom

37. Solution Equilibrium – solids in liquids So lowering the temperature of the coffee will shift the equilibrium or balance point to the left as more sugar precipitates out and settles on the bottom: C 12 H 22 O 11(s)  C 12 H 22 O 11 (aq) Salts in solution will act the same way NaCl (s)  Na + (aq) + Cl + (aq) –An increase in temperature will shift the equilibrium point to the right as more salt dissolves at higher temperatures –A decrease in temperature will shift the equilibrium point to the left as more salt precipitates out of solution

38. Solution Equilibrium – Gases in solution Gases dissolved in Liquids –In a closed system (closed container) –Equilibrium may exist between dissolved gas and the gas dissolved in a solution CO 2(g)  CO 2(aq) –Note: gas equilibrium is affected by both pressure AND temperature! –Increase the pressure – more CO 2 is dissolved –Increase the temperature – less CO 2 is dissolved –In each case the equilibrium is shifted either to left or right in the equation –The rates will be equal, but the amounts may differ

39. Equilibrium Aim KE 5e – What happens in chemical equilibrium reactions?

40. Chemical Equilibrium In any chemical reaction: –Many chemical reactions go to completion –Examples Pb(NO 3 ) 3(aq) + KI (aq)  KNO 3(aq) + PbI (solid) The reverse reaction cannot happen because PbI 2 precipitates out of the solution and is no longer available to react CaCO 3(s)  CaO (s) + CO 2(g) The reverse reaction cannot happen because the CO 2 gas is released and cannot react with the CaO to become CaCO 3 again

41. Chemical Equilibrium In many other chemical reactions the reactions are reversible –Ex: the Haber Process (production of ammonia) In a reaction chamber at equilibrium, the H 2 and N 2 react to create NH 3 3H 2(g) + N 2(g)  2NH 3(g) + 91.8 kJ this is an exothermic rxn; heat is lost in the rxn NH 3 breaks down into H 2 and N 2 at the same rate 2NH 3(g) + 91.8 kJ  3H 2(g) + N 2(g) this is an endothermic rxn; heat is gained in the rxn

42. Chemical Equilibrium A balance or equilibrium occurs between the two reaction rates –The forward rate = the reverse rate 3H 2(g) + N 2(g)  2NH 3(g) + 91.8 kJ –The amounts of H 2, N 2 and NH 3 are constant at a given temperature and pressure –If stress is applied to this system, the equilibrium of the system will change as well LeChatelier’s Principle: –any system in equilibrium can be disturbed by adding a stress to it –Stress will cause the system to reach a new equilibrium point

43. Le Chatelier’s and Chemical Systems In the Haber process of ammonia synthesis N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ At equilibrium the concentrations of nitrogen, hydrogen, and ammonia are constant But a stress added to this system will cause a chemical shift –Either the stress will favor the forward rxn or the reverse rxn –This will change the concentrations of the substances involved –And eventually make the rates of the forward and reverse reactions equal again

44. Le Chatelier’s Principle and Chemical Systems 1. Effect of Concentration on Equilibrium –In the Haber process N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ –The concentration (molarity, M) is expressed in moles per liter –[N 2 ] = concentration of nitrogen, –[H 2 ] = concentration of hydrogen –[NH 3 ] = concentration of ammonia –When the concentration of one of the above increases, –the system will shift to reduce that amount and return the system to a new equilibrium

45. Le Chatelier’s and Chemical Systems –In the Haber process N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ –Increasing the [N 2 ] makes more N 2 molecules available to react with H 2 molecules More NH 3 is then produced, increasing [NH 3 ] but also decreasing the [H 2 ] A new equilibrium point is then reached –Increasing the [NH 3 ] makes more [N 2 ] and [H 2 ] as the increased number of NH 3 molecules allows more to break down

46. Le Chatelier’s Principle and Chemical Systems 2 - Effect of a Change in Pressure on Equilibrium –In the Haber process N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ –increasing the pressure brings all the molecules closer together and the system tries to relieve the stress –A shift will occur in the direction that will produce the fewest number of molecules that take up less space

47. Le Chatelier’s Principle and Chemical Systems N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ –Four molecules are on the left, and only two are on the right –As the pressure increases and NH 3 molecules are formed, there are fewer molecules in the system –Less force against the container side due to less molecules –Less pressure as the system shifts to the right In the reaction below there is no effect on the system by changing the volume… why? H 2(g) + Cl 2(g)  2HCl (g) Two on the left, two on the right, already equal

48. Le Chatelier’s and Chemical Systems 3. Effect of Temperature –In the Haber Process N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ –The forward reaction is exothermic –The reverse reaction is endothermic –Which needs more heat to happen – the endothermic or exothermic reaction? An increase in temperature always favors the endothermic reaction more –In the Haber Process, therefore, the equilibrium shifts to the left as this adds required heat to the right side of the reaction

49. Le Chatelier’s and Chemical Systems 4. Effect of Catalysts on equilibrium –In the Haber Process N 2(g) + 3H 2(g)  2NH 3(g) + 91.8 kJ –A catalyst lowers the activation energy of both forward and reverse reactions –Therefore, there is an equal affect on both reactions and their rates –Catalysts don’t affect equilibrium