I.1 REACTION KINETICS

Transcription

1 I.1 REACTION KINETICS KEY QUESTION: Why do reactions occur and how do you control them? REACTION KINETICS is the study of the REACTION RATES Express REACTION RATE as Example 1: The rate of a reaction is g of Mg per second. Calculate the number of moles of Mg used up in 6.0 minutes. Example 2: An experiment is done to determine the rate of the following reaction: Al (s) + 6Cl (aq) 3 2(g) + 2AlCl 3(aq) It is found that the rate of production of 2(g) is g/s. Calculate the mass of aluminum reacted in 3.0 minutes. 1

2 Reaction rates are not usually linear relationships. So, often an average rate is calculated over a time interval instead. I.2 METODS OF MEASURING REACTION RATES Properties can be monitored to determine reaction rate. Different properties work for different reactions. Decide the method by looking at 1. Write equation in ionic form. 2. Look at subscripts and use your common sense. Consider: CaCO 3(s) + 2Cl (aq) 2 O (l) + CO 2(g) + CaCl 2(aq) ionic form: Observe: Cu (s) + 4NO 3(aq) Cu(NO 3 ) 2(aq) O (l) + 2NO 2(g) + heat 2

3 Method ow can we use this to measure rate? 1. Colour Change 2. Concentration Change 3. Temperature Change 4. Pressure Change For endothermic reactions For exothermic reactions must pay attention to moles of gas on reactant side vs. product side 5. Volume Change 6. Mass Change of solid of gas escaping ebden: Do Set 1 #1, 2, 4 9 (p. 2-5) I.3 FACTORS AFFECTING REACTION RATE As concentration increases, reaction rate As temperature increases, reaction rate As pressure increases, reaction rate 3

4 If a catalyst is added, reaction rate Nature of reactants As surface area increases, reaction rate I.4 NATURE OF REACTANTS The NATURE OF REACTANTS refers to the of a substance. PASE AQUEOUS IONS > GASES AND LIQUIDS > SOLIDS OMOGENEOUS reactions are than ETEROGENEOUS reactions Rate depends on BOND TYPE BONDS Breaking reactant bonds and reforming product bonds. A reaction with bonds to break and/or form will have a FASTER reaction rate. Examples: Electron Transfer Precipitation Acid-Base (proton transfer) 4

5 Example 3: Which of the following reactions will be slowest at 25 C? Explain. A. Cu (s) + S (s) CuS (s) B. + (aq) + O - (aq) 2 O (l) C. Pb 2+ (aq) + 2Cl - (aq) PbCl 2(s) D. 2NaOCl (aq) 2NaCl (aq) + O 2(g) CATALYSTS AND INIBITORS A is a substance which can the rate of the reaction but it is. An is a substance which a reaction rate by combining with a catalyst or one of the reactants in such a way as to prevent the reaction from occurring. Examples: poisons, common antibiotics, sunscreens, antidepressants ebden: Do Set 2 #10 12, 13 b, d, e, 15, 17, 19 (p.7 11) I.5 COLLISION TEORY Reaction rates can be explained using the COLLISION TEORY or KINETIC MOLECULAR TEORY which uses a particle view of reactions. 2 + I 2 2I Particles are In at different speeds with each other Reaction involves the Reaction rate depends on the Use the Collision Theory to complete the Reason column in the factors affecting reaction rates Table. Example 1: Use collision theory to explain the effect of CONCENTRATION. 5

6 Example 2: Use collision theory to explain the effect of TEMPERATURE. I.6 ENERGIES AND ENTALPY CANGES BOND ENERGY is the amount of energy required to break a bond between two atoms. Example: Cl 2(g) + 243kJ 2Cl (g) POTENTIAL ENERGY (PE) is the energy due to an object s in space, as well as the sum of all the attractive and repulsive forces existing among the particles which make up the object. It is related to bond energy, and the number and type of atoms in the molecules. KINETIC ENERGY (KE) is the energy which a system possesses because of within the system. WAT IS ENTALPY? ENTALPY is the = heat of reaction or energy or reaction symbol = symbol for change in enthalpy = ow does enthalpy apply to reactions? Reaction can Reaction can Exothermic Reactions Endothermic Reactions Thermal chemical equation: Thermal chemical equation: 6

7 I.7 KINETIC ENERGY DISTRIBUTIONS 1. Think of kinetic energy as 2. In a room full of reactants, they all have 3. of kinetic energies. KE DISTRIBUTIONS Increase temperature Average KE increases KE DISTRIBUTIONS AT 2 TEMPERATURES Area under high temperature curve almost the area under the low temperature curve. 7

8 This means: In General: KE DISTRIBUTIONS FOR FAST REACTIONS ebden: Do Set 3 #23, 24, 26, 28, 29, 31 (p.13-20) POTENTIAL ENERGY DIAGRAMS So KE PE KE + PE = TOTAL ENERGY (constant) Work with a partner, graph the change in the level of the KE and PE as a reaction between two species proceeds below. (PE + KE = Total Energy = Constant) 8

9 1) At the start of the reaction, most of the energy is KE. Use one colour for KE. As the molecules approach each other, the PE increases and the KE (as velocity) decreases. Graph decreasing KE and increasing PE. After the reaction, the molecules separate and pick up speed as the PE decreases. 2) Do the PE graph in a different colour. Remember when KE goes up, PE goes down. Energy Progress of reaction Compare the 2 collisions I.8 ACTIVATION ENERGY What is ACTIVATION ENERGY E a? What is an ACTIVATED COMPLEX? O O O reactants O O O O O activated complex products 9

10 E a depends on NATURE OF REACTANTS NOT on Δtemperature or Δconcentration # of Molecules Curve at Lower Temp. COLLISION GEOMETRY Activation Energy Kinetic Energy COLLISION GEOMTRY A 2 + B 2 2AB Collision has Therefore, it needs higher energy to have an effective collision. Collision has. Therefore, it needs less energy to have an effective collision. 10

11 Requirements for EFFECTIVE COLLISION: REVERSIBLE REACTIONS: REVERSIBLE REACTIONS DRAWING AND LABELLLING PE DIAGRAMS 1. Draw and label axes. 2. Draw in energy levels for reactants and products (pay attention to endothermic/exothermic!) 3. Connect the energy levels with a bump. 4. E a(f) Draw an arrow pointing from starting reactant of forward reaction to peak 5. E a(r) Draw an arrow pointing from starting reactant (product) of reverse reaction to peak E a(f) = Example 4: Find the E a and for the forward and reverse reactions. Is the forward reaction exothermic or endothermic? E a(f) = kj E a(r) = kj (forward) = (reverse) = kj kj The forward reaction is 11

12 ebden: Do Set 4 #34 38, 40, 41, 44, (45) (p ) I.9 REACTION MECANISM TEORY REACTION MECANISM is a sequence of individual steps that make up the overall reaction. 5C 2 O MnO This reaction involves 23 particles colliding at once to react. Most reactions, take a series of steps to reach the product. Each step can proceed at is own rate. fast Step O O 2 Step 2 3 O 2 + slow + I 2 O + OI Overall O 2 + I - 2 O + OI Example 5: Step 1: Br + O 2 OOBr (slow) Step2: Br + OOBr 2OBr (fast) Step 3: 2OBr + 2Br 2 2 O + 2Br 2 (very fast) Overall: a) What is the overall reaction? (cancel identical species on opposite sides and add up what s left) b) What is the formula for the activated complex of Step 1? (combine all the reactant elements together, add up the charges) c) List the reaction intermediates. (reaction intermediates produced in one step and used up in a later step, so it appears on the right then on the lower left) 12

13 d) Complete the table. Reaction intermediate [ ] is high or low? Reason e) What is the difference between a reaction intermediate and an activated complex in terms of their energy level and stability? I.10 PE DIAGRAMS OF A REACTION MECANISM ebden: Do Set 5# 46, 47, 51, 54, 55 (p.28 30) I.11 CATALYSTS A CATALYST is a substance which provides an overall reaction with an which has a. EFFECT OF CATALYSTS ON E A 13

14 Example 6: a) Draw the PE Diagram for ClO + I IO + Cl (pay attention to # steps and change in ). b) The intermediates are, the catalyst is. int: You must complete the catalyzed reaction mechanism first and identify all species. Catalyzed steps: Step 1 ClO + ClO + O Δ = Step 2 I + ClO IO + slow Δ = + Step 3 + O IO + 2 O fastest Δ = Overall ClO + I IO - + Cl Δ = PE Reaction Proceeds USES OF CATALYSTS Read p and give some examples of the uses of catalysts. ebden: Do Set 6 #56, 58, 59, 61, 62 (p.34) 14