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Equilibrium. Non reversible reactions Some chemical and physical reactions occur until one or all the reactants are used up Example 1 Evaporation of water.

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Presentation on theme: "Equilibrium. Non reversible reactions Some chemical and physical reactions occur until one or all the reactants are used up Example 1 Evaporation of water."— Presentation transcript:

1 Equilibrium

2 Non reversible reactions Some chemical and physical reactions occur until one or all the reactants are used up Example 1 Evaporation of water from a beaker (a physical process) Reaction is complete when all the H 2 O molecules are in the gas phase gaseous molecules from evaporation water in liquid phase Evaporation continues Fig 1: Physical process going to completion

3 Non reversible chemical reactions Hydrogen gas Magnesium Example 2 Magnesium in excess dilute acid solution Eventually all the magnesium reacts and gas production stops In examples 1 and 2 the forward reaction continues until the reactant(s) are used up this is written as: Reactants Products Fig 2: Chemical process going to completion

4 Reversible Reactions If we warm our beaker of water with glad wrap over the top we get a reversible reaction because the reaction is now in a closed system Evaporation is occurring but so is the opposing process of condensation Example 3 Gaseous water molecules from evaporation Water in liquid phase The forward reaction is from left to right H 2 O (l) H 2 O (g) The reverse reaction is H 2 O (g) H 2 O (l) The overall equilibrium reaction is written as H 2 O (l) H 2 O (g) This reaction has reached Equilibrium when evaporation and condensation occur at the same rate

5 Equilibrium When reversible reactions reach a point where the rate of forward reaction equals the rate of the backward reaction the system is said to have reached equilibrium The equilibrium is said to be dynamic which means the forward and backward reactions do not stop

6 Important Points so Far 1.Not all reactions are equilibrium reactions (ie some reactions are not reversible) 2.If an equilibrium reaction is to occur it must be in a closed system 3.When a reaction has reached dynamic equilibrium this means that the forward reaction is occurring at the same rate as the reverse reaction and this is written as: reactants products

7 Chemical Equilibrium Many chemical reactions behave as a dynamic equilibrium An Example is the Iron(lll) / thiocyanate equilibrium Fe 3+ (aq) + SCN - (aq) yellow colourless ion In the reaction A.The red colour forms quickly then remains constant B.The system is closed as the ions Fe 3+, SCN - and FeSCN 2+ are all contained in the same aqueous solution FeSCN 2+ (aq) bright red complex

8 Iron(lll) / thiocyanate equilibrium continued The reaction hasnt gone to completion because it can be shown that their are unreacted Fe 3+, SCN - ions in the red solution, this shows the reverse reaction is also occurring FeSCN 2+ (aq) Fe 3+ (aq) + SCN - (aq) The equilibrium for the reaction is written: Fe 3+ (aq) + SCN - (aq) FeSCN 2+ (aq) At equilibrium the rate at which the complex ions (FeSCN 2+ ) form is equal to the rate at which the complex ions (FeSCN 2+ ) break up

9 Part A Fe 3+ (Iron III) ions (yellow) & SCN - (thiocyanate) ions (colourless) These two ions react to form a reddish complex ion: Fe 3+ + SCN - FeSCN 2+ Make some of this and place on 4 petrie dishes

10 Control Add more Fe 3+ Add more SCN - Add F -

11 Why does a reaction stop? Until now, you expect it to be because one or both of the reactants have run out In this reaction, neither the Fe 3+, nor the SCN - have run out, yet the reaction appears to stop Why?

12 Two processes occur one to make FeSCN 2+ the other to decompose FeSCN 2+ The reaction stops when these two reactions balance out Fe 3+ + SCN - forming complex FeSCN 2+ Fe 3+ + SCN - decomposing FeSCN 2+ This is written as: Fe 3+ + SCN - FeSCN 2+

13 Question - Which ones correct? Equilibrium occurs only when: A.The reactants are used up B.The concentration of reactants is equal to the concentration of products C.All chemicals stop reacting D.The products react together to form the reactants at the same rate as reactants form products Ans - D

14 Starter Questions What two conditions do you need for an equilibrium reaction to occur? The reaction must be a reversible reaction The reaction must be in a closed system What is meant by the term equilibrium? The forward reaction is occurring at the same rate as the reverse reaction (rate forward = rate reverse) What factors can affect an equilibrium reaction? Concentration of reactants or products Temperature Surface area Pressure

15 Equilibrium constant Scientists wanted to put a number on the proportions of products and reactants when equilibrium is reached They did heaps of experiments and eventually they came upon a method for finding the equilibrium constant (K) for each particular reaction

16 Equilibrium Constant For a general reversible reaction: aA + bB cC + dD a moles of b moles of c moles of d moles of substance A substance B substance C substance D The ratio of products to reactants is expressed as the equilibrium constant symbol K and is written as remember as

17 Writing Equilibrium Equations Writing the following Equilibrium expression for: CO (g) + H 2 O (g) CO 2 (g) + H 2 (g) Where: [B] is the concentration of B in molL -1 at equilibrium [A] is the concentration of A in molL -1 at equilibrium

18 Equilibrium Constants Equilibrium constants (K) are specific for a given reaction and temperature. An equilibrium constant is useful for determining how far a reaction has proceeded. If the equilibrium constant (K) is large eg > 10 5 then at equilibrium the reaction has proceeded almost to completion.

19 Equilibrium Constants (continued) If the equilibrium constant (K) is small eg < then at equilibrium only a small amount of reactant is used up and only a small amount of product has formed then the reaction has effectively not happened

20 Writing Equilibrium expressions: Task: Write the equilibrium expression for CH 4 (g) + H 2 O (g) CO (g) + 3H 2 (g) Be very careful!

21 For equilibrium systems involving solid(s) and pure liquids The solid(s) and liquids are not included in the equilibrium expression Eg: write the equilibrium expression for C (s) + H 2 O (g) CO (g) + H 2 (g) Be very careful!

22 Writing more Equilibrium expressions: Eg: write the equilibrium expression for S (s) + O 2 (g) SO 2 (g) Be very careful!

23 N 2 (g) + 3H 2 (g) 2NH 3 (g) Starter : write the equilibrium expression for:

24 2N 2 (g) + 6H 2 (g) 4NH 3 (g) write the equilibrium expression for:

25 Moving Equilibrium Systems Le Chateliers Principle: When change is applied to a system at equilibrium, the system responds so that the effects of the change are minimised

26 Example: H 2 (g) + Cl 2 (g) 2HCl(g) What would be the effect on the amount of HCl formed if hydrogen gas was removed ? The amount of HCl formed would decrease because the system would try to produce more H 2 by shifting the equilibrium to the LHS. In the above reaction what would be the effect on the amount of HCl formed if Cl 2 gas was added? The amount of HCl formed would increase, because the system will try to minimise the change and remove Cl 2 by shifting the equilibrium to the RHS producing more HCl.

27 Turn to page 147 in your lab book – Equilibrium Systems Lets see how changing concentration and temperature affect the position of the equilibrium

28 N 2 O 4(l) 2NO 2(g) ΔH = +54kJ (colourless) (brown) Or… N 2 O 4(l) + energy2NO 2(g) Forward reaction is endothermic Backward reaction is exothermic Cooling favours exothermic Heating favours endothermic

29 So heating this system will move the equilibrium to the right and the system will appear dark brown Cool it and move the equilibrium system to the left and the system will appear pale brown

30 The endothermic (+ H) direction is the backward reaction, so an increase in temperature would move the equilibrium in the backward direction to take more energy in and therefore minimise the increase in temperature. This would decrease the amount of NH 3 produced. Example N 2 (g) + 3H 2 (g) 2NH 3 (g) H= -92 kJ H= +92 kJ

31 Consider another form of the N 2 O 4 /NO 2 system N 2 O 4(g) 2NO 2(g) (colourless) (brown) Fewer molecules More molecules Fewer molecules More molecules

32 If the pressure is decreased, the system opposes this by increasing the number of molecules in the system –Forms more NO 2(g) –colour will go from light brown to a darker brown as new equilibrium established Fewer molecules More molecules

33 Pressure N 2 (g) + 3H 2 (g) 2NH 3 (g) Now what would happen to the amount of NH 3 produced if the pressure was decreased? For the same reaction at equilibrium, an increase in the volume of the container, (or decrease in pressure), will shift the equilibrium to the side with more gas particles. This means the concentration of NH 3 would decrease.

34 Reactant concentration –Increase – reaction moves forwards (to the right) –Decrease- reaction moves backwards (to the left) Product concentration –Increase – reaction moves backwards –Decrease – reaction moves forwards Catalyst No change in equilibrium position, but equilibrium reached more quickly.

35 Exercise 1.For the equilibrium N 2 O 4 (g) 2NO 2 (g) r H= -92 kJ mol -1 what would happen to the number of NO 2 molecules if the following changes are made to the system at equilibrium a) an increase in pressure b) the concentration of N 2 O 4 is increased c) the pressure of the system is decreased d) the temperature is decreased e) a catalyst is added

36 PCl 3 (g) + Cl 2 (g) PCl 5 (g) H= +110 kJ mol -1 what would happen to the amount of PCl 5 if the following changes are made to the system at equilibrium a) an increase in pressure b) the concentration of Cl 2 is increased c) the volume of the container is increased d) the temperature is decreased Starter Equilibrium Changes

37

38 Strategies to lower costs include Buy raw materials as cheaply as possible Use energy efficiently Recycle unused reactants and by- products, if possible

39 Heat exchangers may be used to remove energy from an exothermic process and deliver it to an endothermic process Catalysts assist in the rapid establishment of an equilibrium

40 A severe shortage of nitrogen- based fertilisers at the start of the 20 th century started hunt for a method of using nitrogen in the air –Nitrogen = 71% of air Fritz Haber developed the Haber Process used today to make ammonia

41 Developed Ecstasy Did fundamental research into the Bunsen flame Developed a process for extracting gold from seawater Noble prize 1918 for Haber process

42 Haber died in exile after standing up to Hitlers anti- Semitic policies which he felt were morally wrong and were decimating German research institutes

43 The Equilibrium N 2(g) + 3H 2(g) 2NH 3(g) -ΔH To move the equilibrium to the right means a pressure drop and the production of heat –High temperatures favour the backward reaction but make the reaction rate faster –High pressures favour the forward reaction and increase the yield of ammonia

44 % Yield of Ammonia Using The Haber Process Pressure (atm) T( 0 C)


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