2 Non reversible reactions Some chemical and physical reactions occur until one or all the reactants are used upExample Evaporation of water from a beaker (a physical process)Reaction is complete when all the H2O molecules are in the gas phasegaseous molecules from evaporationwater in liquid phaseEvaporation continuesFig 1: Physical process going to completion
3 Non reversible chemical reactions Example 2 Magnesium in excess dilute acid solutionEventually all the magnesium reacts and gas production stopsHydrogen gasMagnesiumFig 2: Chemical process going to completionIn examples 1 and 2 the forward reaction continues until the reactant(s) are used up this is written as:Reactants Products
4 Reversible ReactionsIf 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 systemExample 3Evaporation is occurring but so is the opposing process of condensationGaseous water molecules from evaporationWater in liquid phaseThe forward reaction is from left to rightH2O (l) H2O (g)This reaction has reached Equilibrium when evaporation and condensation occur at the same rateThe reverse reaction isH2O (g) H2O (l)The overall equilibrium reaction is written asH2O (l) H2O (g)
5 EquilibriumWhen 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 equilibriumThe equilibrium is said to be dynamic which means the forward and backward reactions do not stop
6 Important Points so Far Not all reactions are equilibrium reactions (ie some reactions are not reversible)If an equilibrium reaction is to occur it must be in a closed systemWhen 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 EquilibriumMany chemical reactions behave as a dynamic equilibriumAn Example is the Iron(lll) / thiocyanate equilibriumFe3+ (aq) SCN-(aq)yellow colourless ionIn the reactionThe red colour forms quickly then remains constantThe system is closed as the ions Fe3+, SCN- and FeSCN2+ are all contained in the same aqueous solutionFeSCN 2+(aq)bright red complex
8 Iron(lll) / thiocyanate equilibrium continued The reaction hasn’t gone to completion because it can be shown that their are unreacted Fe3+, SCN- ions in the red solution, this shows the reverse reaction is also occurringFeSCN 2+(aq) Fe3+ (aq) SCN-(aq)The equilibrium for the reaction is written:Fe3+ (aq) SCN-(aq) FeSCN 2+(aq)At equilibrium the rate at which the complex ions (FeSCN2+) form is equal to the rate at which the complex ions (FeSCN2+) break up
9 Part AFe3+ (Iron III) ions (yellow) & SCN- (thiocyanate) ions (colourless)These two ions react to form a reddish complex ion:Fe3+ + SCN- FeSCN2+Make some of this and place on 4 petrie dishes
11 Why does a reaction stop? Until now, you expect it to be because one or both of the reactants have run outIn this reaction, neither the Fe3+, nor the SCN- have run out, yet the reaction appears to stopWhy?
12 Two processes occur one to make FeSCN2+ the other to decompose FeSCN2+ The reaction stops when these two reactions balance outFe3+ + SCN- forming complex FeSCN2+Fe3+ + SCN- decomposing FeSCN2+This is written as:Fe3+ + SCN- FeSCN2+
13 Question - Which one’s correct? Equilibrium occurs only when:The reactants are used upThe concentration of reactants is equal to the concentration of productsAll chemicals stop reactingThe products react together to form the reactants at the same rate as reactants form productsAns - D
14 Starter QuestionsWhat two conditions do you need for an equilibrium reaction to occur?The reaction must be a reversible reactionThe reaction must be in a closed systemWhat 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 productsTemperatureSurface areaPressure
15 Equilibrium constantScientists wanted to put a number on the proportions of products and reactants when equilibrium is reachedThey did heaps of experiments and eventually they came upon a method for finding the equilibrium constant (K) for each particular reaction
16 Equilibrium ConstantFor a general reversible reaction:aA bB cC dD‘a’ moles of ‘b’ moles of ‘c’ moles of ‘d’ moles ofsubstance A substance B substance C substance DThe ratio of products to reactants is expressed as the equilibrium constant symbol K and is written asremember as
17 Writing Equilibrium Equations Where:[B] is the concentration of B in molL-1 at equilibrium[A] is the concentration of A in molL-1 at equilibriumWriting the following Equilibrium expression for:CO (g) + H2O (g) CO2 (g) + H2 (g)
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 largeeg > 105 then at equilibrium the reaction has proceeded almost to completion.
19 Equilibrium Constants (continued) If the equilibrium constant (K) is smalleg < 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 Be very careful! Writing Equilibrium expressions: Task: Write the equilibrium expression forCH4 (g) H2O (g) CO (g) H2 (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 expressionEg: write the equilibrium expression forC (s) H2O (g) CO (g) H2 (g)Be very careful!
22 Be very careful! Writing more Equilibrium expressions: Eg: write the equilibrium expression forS(s) O2 (g) SO2 (g)Be very careful!
24 write the equilibrium expression for: 2N2(g) H2(g) NH3(g)
25 Moving Equilibrium Systems Le Chatelier’s Principle: When change is applied to a system at equilibrium, the system responds so that the effects of the change are minimised
26 Example: H2(g) + Cl2(g) 2HCl(g) In the above reaction what would be the effect on the amount of HCl formed if Cl2 gas was added?The amount of HCl formed would increase, because the system will try to minimise the change and remove Cl2 by shifting the equilibrium to the RHS producing more HCl.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 H2 by shifting the equilibrium to the LHS.
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 Forward reaction is endothermic Backward reaction is exothermic N2O4(l) 2NO2(g) ΔH = +54kJ(colourless) (brown)Or…N2O4(l) + energy 2NO2(g)Forward reaction is endothermicBackward reaction is exothermicCooling favours exothermicHeating 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 This would decrease the amount of NH3 produced. ExampleN2(g) + 3H2(g) NH3(g) H= -92 kJH= +92 kJThe 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 NH3 produced.
31 Consider another form of the N2O4/NO2 system N2O4(g) 2NO2(g)(colourless) (brown)Fewer molecules More moleculesFewer moleculesMore molecules
32 If the pressure is decreased, the system opposes this by increasing the number of molecules in the systemForms more NO2(g)colour will go from light brown to a darker brown as new equilibrium establishedFewer moleculesMore molecules
33 Pressure N2(g) + 3H2(g) 2NH3(g) Now what would happen to the amount of NH3 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 NH3 would decrease.
34 Reactant concentration Increase – reaction moves forwards (to the right)Decrease- reaction moves backwards (to the left)Product concentrationIncrease – reaction moves backwardsDecrease – reaction moves forwardsCatalyst No change in equilibrium position, but equilibrium reached more quickly.
35 ExerciseFor the equilibriumN2O4(g) NO2(g) rH= -92 kJ mol -1what would happen to the number of NO2 molecules if the following changes are made to the system at equilibriuma) an increase in pressureb) the concentration of N2O4 is increasedc) the pressure of the system is decreasedd) the temperature is decreasede) a catalyst is added
36 Starter Equilibrium Changes PCl3 (g) + Cl2(g) PCl5 (g) H= +110 kJ mol -1what would happen to the amount of PCl5 if the following changes are made to the system at equilibriuma) an increase in pressureb) the concentration of Cl2 is increasedc) the volume of the container is increasedd) the temperature is decreased
38 Strategies to lower costs include Buy raw materials as cheaply as possibleUse energy efficientlyRecycle 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 processCatalysts assist in the rapid establishment of an equilibrium
40 Fritz Haber developed the “Haber Process” used today to make ammonia A severe shortage of nitrogen-based fertilisers at the start of the 20th century started hunt for a method of using nitrogen in the airNitrogen = 71% of airFritz Haber developed the “Haber Process” used today to make ammonia
41 Developed “Ecstasy”Did fundamental research into the Bunsen flameDeveloped a process for extracting gold from seawaterNoble prize 1918 for Haber process
42 Haber died in exile after standing up to Hitler’s anti-Semitic policies which he felt were morally wrong and were decimating German research institutes
43 The Equilibrium N2(g) + 3H2(g) 2NH3(g) -ΔH To move the equilibrium to the right means a pressure drop and the production of heatHigh temperatures favour the backward reaction but make the reaction rate fasterHigh pressures favour the forward reaction and increase the yield of ammonia
44 % Yield of Ammonia Using The Haber Process Pressure (atm)T(0C)10501003001000200517482909815395271934004254780500161126576000.5251413