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Rates & Extents of Reactions.

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Presentation on theme: "Rates & Extents of Reactions."— Presentation transcript:

1 Rates & Extents of Reactions

2 Overview Quiz A What prevents reaction from going?
Increasing the Rate of reaction Temperature Quiz B Catalysis Collisions Industry Rates & Extents of Reactions Measuring Change Temperature Quiz C Quiz G Pressure Concn. Mass of System Le Chatelier’s Principle The Equilibrium Constant Establishing Equilibrium Quiz F Quiz E Dynamic Eqm. Quiz D

3 Quiz A- Common misconceptions
When solid copper carbonate reacts with excess acid, carbon dioxide is produced. The curves shown were obtained under two different sets of conditions. X Volume CO2 of time Y The change from X to Y could be brought about by : A. increasing the concentration of the acid. B. decreasing the particle size. C. adding a catalyst. D. increasing the temperature. E. decreasing the mass of copper carbonate.

4 Quiz A- Common misconceptions
2. Chemical reactions are in a state of dynamic equilibrium only when: A. the rate of the reverse reaction is equal to the rate of the forward reaction. B. the reaction involves a zero enthalpy change. C. the activation energy of the forward reaction equals the activation energy of the reverse reaction. D. the reaction goes to completion. E. the ratio of the products to reactants is equal to exactly 1.

5 Quiz A- Common misconceptions
X(g) is placed in a flask and the following reaction proceeds to equilibrium X(g) ⇌ Y(g) ΔH = +ve Which of the following statements is correct? A. The forward reaction rate increases as the reaction “gets going”. B. The forward reaction rate always equals the reverse reaction rate. C. The activation energy of the forward reaction of the system considered will always be higher than the activation energy of the reverse reaction no matter the reaction conditions. D. When equilibrium is re-established after a disturbance (e.g. adding more Y) the rate of the forward and reverse reactions will be equal to those at the initial equilibrium. E. The forward reaction is completed before the reverse reaction begins.

6 Quiz A- Common misconceptions
The graph below shows the variation of concentration of a reactant , X, with time as the reaction proceeds. [X] / Time /s 30 60 0.020 0.005 What is the average reaction rate (mol dm-3 s-1) during the first 30 seconds? A B 0.005 C 0.015 D 0.020 E The slope of the tangent to the curve at t = 30s

7 Quiz A- Common misconceptions
Excess marble chips (CaCO3) were added to 50 cm3 of 1 M HCl. The experiment was repeated using the same mass of marbles chips and 50 cm3 of 1M CH3COOH. Which of the following would have been the same for both experiments? A The average rate of reaction. B The rate at which the first 2 cm3 of gas was evolved. C The time taken for the reaction to be complete. D The mass of the marble chips remaining after the reaction had stopped. E None of the above.

8 What prevents reactions from going?
Potential Energy Reaction co-ordinate Ea Energy Profile 2NO  N2 + O2 2NO ΔH N2 + O2

9 Rate = k [A]x [B]y Increasing the Rate of Reaction Orders of reactants
Rate constant Rate = k [A]x [B]y Non-examinable ! But useful for understanding.

10 Increasing the Rate of Reaction
A. Temperature Everyone knows that as we increase the temperature, the reaction rate goes up…… But how do we explain this observation scientifically? Q A The answer lies in the Maxwell Distribution where the idea of a range of molecular speeds is introduced.

11 What makes reactions go faster?
Fraction of molecules Molecular Velocity T2 Maxwell Distribution of velocities T2 > T1 KEmin = ½mv2min vmin T2 T1

12 Increasing the Rate of Reaction
If we increase the temperature we will increase the fraction of the molecules with the energy to overcome the activation energy. f = e –Ea/RT Where f = fraction of molecular collisions with energy greater or equal to Ea If T  then f  If f  then k  If k  then Rate Note : If Ea due to a catalyst then f  and the Rate 

13 Increasing the Rate of Reaction
B. Catalysis  Catalysts are employed to speed–up the attainment of equilibrium.  Catalysts do not change the position of equilibrium .  Catalysts offer a different mechanism for the reaction to occur . For the catalyst to be successful it must offer a mechanism with a lower activation energy, Ea. Catalysts can either be in the same phase as the reactant – homogeneous catalysts -or a different phase –heterogeneous catalysts.

14 Increasing the Rate of Reaction
Potential Energy Reaction co-ordinate Ea Energy Profile N2(g) + O2(g)  2NO(g) Ea 2NO ΔH N2 + O2 O N N O N O Catalyzed Non – catalyzed N N O N O O

15 Increasing the Rate of Reaction
C. Increasing collisions The more collisions there are the more likely reactants are to react. We can influence the amount of collisions by: 1. Increasing the concentration of reactants (pressure in the case of gases). 2. Increasing the surface area of solid reactants (or solid catalysts).

16 Quiz B- Increasing reaction rates
A catalyst lowers the activation energy of a reaction from 100 kJ/mol to 50 kJ/mol. What increase in f will result? The activation energy for a reaction is known to be 50 kJ/mol of reactant. If the temperature is increased from 300 K to 310 K what increase in the rate of reaction would you expect?

17 Q A Measuring Change A (g)  B(g) Consider the following reaction:
How could we measure the rate of this reaction? A Well let’s consider what might change during this reaction. Im bending your spoon with my mind! ! ! ! ! Mass of system Temperature Pressure Concentration

18 Measuring Change A. Temperature * See Maxwell Distribution
If the reaction is exothermic or endothermic then the temperature will change. Exothermic; ΔH is –ve; T  Endothermic; Δ H is +ve; T  But if the temperature changes then the rate of reaction will change! Why? * T  Rate of reaction  Therefore if we want to measure the rate of reaction we normally keep the temperature of the reactor constant. This is called thermostatting the reactor. Therefore we cannot measure the rate of a reaction using the change in temperature * See Maxwell Distribution

19 Measuring Change B. Pressure We know from Grade 11 that : PV = nRT
Therefore if we keep the volume of the thermostatted reactor constant then: RT V × n P = = const × n P  n Will PA change during the reaction? Q A Of course : PA  as nA  Will PB change during the reaction? Q A Of course : PB  as nB  Will PTOTAL change during the reaction. Q A NO! PTOT   nTOT and nTOT won’t change for this reaction .

20 Measuring Change Therefore we cannot measure the rate of this reaction
using ptotal

21 Measuring Change C. Mass of the system
This one is easy – if the reactor is a closed system then the mass cannot change. Mass must be conserved. If the reaction is taking place in an open system then gases might leave the system and the mass of the system will decrease. In this particular case we have a closed system. Therefore we cannot measure the rate of This reaction using the mass of the system

22 Measuring Change D. Concentration
This is by far the most common means of measuring rates, because concentrations ALWAYS change during a reaction. In this example : [A]  with time while [B]  We need to be able to measure concentrations and the most used technique is SPECTROSCOPY where a reactant or product selectively absorbs or emits electromagnetic radiation, leaving it coloured. The intensity or amount of the radiation absorbed (A) is proportional to the concentration. A  [ ] Normally chemists measure concentrations to determine reaction rates A substance does not need to be coloured to absorb electromagnetic radiation. CO2 for instance is colourless as it does not absorb visible radiation but it does absorb infra-red radiation (the cause of global warming!). Therefore we could use infra-red spectroscopy to monitor [CO2].

23 Quiz C- Measuring change
1. Which of the following reactions could we use PTOTAL for the measurement of the reaction rate? a) 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g) b) O3(g) + NO(g)  NO2(g) + O2(g) c) 2CaO(s) + 5C(s)  2CaC2(l) + CO2(g) 2. Predict the necessary experimental conditions to be able to measure the rate of the following reaction by monitoring the mass of the system. 2CaO(s) + 5C(s) 2CaC2(l) + CO2(g) Sulfur dioxide gas can be oxidized to sulfur trioxide gas by oxygen gas. In a particular experiment a stoichiometric amount of sulfur dioxide and oxygen are allowed to react in a glass vessel. If the initial pressure in the flask is 3 atm , what would be the predicted final pressure if the reaction goes to completion and the temperature remains constant.

24 Quiz C- Measuring change
4. Calculate the rate of CO2 production (moles of CO 2 / dm3 sec) from an industrial process given the data in the two graphs shown below.

25 Quiz C- Measuring change

26 Establishing Equilibrium
Closed System Open System

27 Establishing Equilibrium
5 10 15 20 Molecules /cm3 time Establishing Equilibrium Solution A at t = 0 sec K = [B] [A] = 3 A ⇌ B

28 5 10 15 20 Molecules /cm3 time Dynamic Equilibrium

29 Quiz D- Establishing Equilibrium
Which of the following systems are at equilibrium ?(Assuming they have been left long enough to establish equilibrium.) a) A boiling kettle. b) A sealed thermostatted test-tube containing a drop of water. c) A completely flat battery. d) A water tank under the following conditions: Water H2O 5 L s-1 300 L min-1

30 aA + bB ⇌ cC + dD The Equilibrium Constant
The Law of Mass action, which is independent of kinetic theory, states that for a reaction- aA + bB ⇌ cC + dD the ratio [C]c [D]d [A]a [B]b will be a constant when the system is at equilibrium. This constant is known as the equilibrium constant, Kc. One must remember that the value of Kc is specific to a particular reaction equation(i.e. one in which the stoichiometry is fixed) and that it is specific to a given temperature.

31 aA(g) + bB(g) ⇌ cC(g) + dD(g)
The Equilibrium Constant For a gaseous reaction the equilibrium constant can be expressed as a ratio of the partial pressures of the products and reactants. For example the equilibrium constant for the reaction : aA(g) + bB(g) ⇌ cC(g) + dD(g) could be expressed as (PC)c (PD)d (PA)a (PB)b KP = Note that if one uses this expression the pressures must be quoted in atmospheres. Also note that the numerical value of KP might not be the same as the numerical value of KC.

32 Quiz E- Equilibrium Constants
The reaction between nitrogen gas and oxygen gas to form nitrogen dioxide gas is shown below. N2(g) + 2O2(g) ⇌ 2NO2(g) a) Write down an expression for the equilibrium constant, Kc, for this reaction. b) Kc is equal to about 2.6 × for this reaction at 25 oC. In a 1 L flask at 25 oC there are 1.0 × molecules of N2, 3.0 × 1015 molecules of O2 and 1.0 × 1012 molecules of NO2. i) Is this system at equilibrium? ii) If this system is not at equilibrium, in what direction will the reaction proceed? What is the numerical value of the equilibrium constants for each of the following reactions? i) ½N2(g) + O2(g) ⇌ NO2(g) ii) 2NO2(g) ⇌ N2(g) + 2O2(g)

33 Quiz E- Equilibrium Constants
At a given constant temperature, a 1 L flask initially containing mol of SO2 and 0.150 mol of O2, is allowed to come to equilibrium. 80% of the SO2 is found to have reacted to form SO3. Calculate the equilibrium constant for the reaction- O2(g) + 2SO2(g) ⇌ 2SO3(g) 3. For the endothermic reaction : 2SO3(g) ⇌ O2(g) + 2SO2(g) state the effect on the equilibrium constant that the following disturbances will have. a) Increasing the concentration of SO3(g). b) Decreasing the concentration of SO2(g). c) Doubling the size of the reaction flask. d) Decreasing the temperature.

34 Le Chatelier’s Principle
If a change is imposed upon a system at equilibrium the position of the equilibrium will shift in a direction that tends to reduce that change. Note : By changing the position of equilibrium we are not changing the equilibrium constant if the temperature remains constant. Equilibrium constants can only change if one changes the temperature of the reaction vessel.

35 Quiz F- Influencing the Equilibrium
Consider one of the gaseous equilibria involved in the industrial preparation of nitric acid by the Ostwald process. 2NO(g) + O2(g) ⇌ 2NO2(g) ΔH = –ve What qualitative effect would the following disturbances have on the position of the equilibrium? a) An increase in PNO. b) An increase in temperature. c) A decrease in reactor volume. d) An increase in pressure via the addition of an inert gas e.g. Ar.

36 Quiz F- Influencing the Equilibrium
2. Consider the following equilibria below : BaCO3(s) + (aq) ⇌ Ba2+(aq) + CO32- (aq) CO32- (aq) + H2O(l) ⇌ CO2(g) + 2OH-(aq) What qualitative effect would the following disturbances have on the position of the equilibrium? a) Making the particle size of the BaCO3 smaller. b) Decreasing the pH of the aqueous solution.

37 Industries want to make money!
Industry Industries want to make money! They want to do things as quickly and as efficiently as possible. However, sometimes, doing something quickly might not mean doing the same thing efficiently. Thermodynamic (i.e. enthalpy ) considerations might sometimes clash with kinetic (i.e. rate) considerations.

38 Quiz G- Equilibrium & rates in industry
Predict the conditions of temperature and pressure required to increase productivity in the following industrial processes. a) N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -ve b) 4NH3(g) + 5O2(g) ⇌ 4NO (g) + 6H2O(g) ΔH = -ve c) 2CaO(s) + 5C(s) ⇌ 2CaC2(l) + CO2(g) ΔH = +ve

39 Quiz G- Equilibrium & rates in industry
2. An industrial process to convert X into Y has the following stoichiometry 2X(g) ⇌ Y(g) ΔH = -ve The reaction is catalysed by a solid heterogenous catalyst. Which of the following set of experimental conditions would an industrial chemist choose to optimize the reaction? X Y + X T = 300K P = 10 atm A) X Y + X T = 600K P = 1 atm B) X Y + X T = 300K P = 10 atm C) X Y + X T = 300K P = 1 atm D) Catalyst

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