1. REACTION RATES AND REVERSIBLE REACTIONS
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2. How Fast Does the Reaction Go?
Reaction Rate
How Fast Does the Reaction Go?
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3. For an industrial process, it is important to obtain the product as
Reaction rate is a measure of how much of the reactants are consumed or how much of the products are formed per unit time.
The consideration of the rate or speed of the reaction is very important in chemistry particularly in industrial processes.
For an industrial process, it is important to obtain the product as
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4. Fast as possible for commercial and economic reasons.
Chemists, by studying factors that affect the rate of reactions, can find out the best conditions for any particular reaction.
Some of the common properties considered are;-
Rate of production of a gas.
Changing mass of a reaction.
Change in intensity of colour.
Formation or disappearance of a precipitate
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5. Factors affecting the rate of a reaction
1. CONCENTRATION
Effect of concentration on rate of reaction.
Flask
Volume of acid(cm³)
Volume of water (cm³)
Concentration (Mol dm⁻³)
Time taken (t sec) A 40 39 B 30 10 42 C 20 86 D
607
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6. conclusion
The higher the concentration of the acid, the faster the reaction.
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7. When the number of particles reduces, the collision rate decreases.
When the number of particles increases, the rate of collision also increases.
When the number of particles reduces, the collision rate decreases.
How does the rate of reaction change as the reaction proceeds?
Consider reaction between magnesium and HCl
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8. Mg(s) + 2HCl(aq) MgCl₂(aq) + H₂(g)
The graph below shows shape of the graph obtained by plotting the volume of the gas produced against time.
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9. Average rate = Volume of hydrogen gas produced Time taken
The rate of production of hydrogen gas is highest at the end of the experiment, then it slows down and eventually stops when one or all of the reactants is used up.
The rate of reaction at any given time of the reaction can be calculated as
Average rate = Volume of hydrogen gas produced
Time taken
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10. e.g.
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11. Average rate of reaction between A and B is given by;
= (V₂ - V₁) cm³ (t₂ - t₁) sec The rate of a reaction at a particular time(t) is determined by drawing a tangent to the curve at a certain point and obtaining its gradient.
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13. Since the rate of reaction is inversely proportional to time taken
The rate of reaction at point c on the curve when time is in seconds is given by;-
rate ∆V = (V₄ - V₃) cm³
∆t (t₄ t₃) sec
Since the rate of reaction is inversely proportional to time taken
rate α ⅟t
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14. A graph of ⅟t(s⁻¹) against concentration of the acid is a straight line.
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15. 20 cm³ of hydrogen peroxide solution was mixed with 20cm³ of sulphuric acid, 30cm³ of potassium iodide(excess), 5 cm³ of sodium thiosulphate and 1 cm³ of starch solution.
The hydrogen peroxide reacts with potassium iodide at a reasonable rate, forming iodine.
As soon as iodine is formed, it reacts rapidly with sodium thiosulphate untill all the sodium thiosulphate is used up.
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16. The time taken for the blue colour to appear was recorded.
The iodine formed after the sodium thiosulphate is used up gives a blue colour with the starch.
The time taken for the blue colour to appear was recorded.
The experiment was repeated using hydrogen peroxide solutions of different concentration and the results are shown below.
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17. Time (sec) for blue colour to appear ⅟t( sec⁻¹)
Hydrogen peroxide
(mol ⁻¹)
Time (sec) for blue colour to appear
⅟t( sec⁻¹)
0.30
20.0
0.050
0.25
24.5
0.041
0.20
30.3
0.033
0.15
40.0
0.0250
0.10
60.6
0.0165
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18. What does ⅟time represent?
Plot a graph of ⅟time against concentration of hydrogen peroxide.
Explain why this graph shows the rate of the reaction is proportional to the concentration of hydrogen peroxide solution.
Use the graph to determine the time that would have been taken for the blue colour to appear if the concentration of hydrogen peroxide had been;
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19. 0.05M
0.18M
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20. 2. TEMPERATURE
Generally a rise in temperature tends to increase the rate of a reaction.
When a mixture of substances is heated, the particles move faster.
This has two effects:
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21. As temperature rises, the particles of the reacting substances move faster hence they collide with each other more often. The rate of reaction therefore increases.
When two molecules collide, they will only react with each other if they have sufficient energy- activation energy.
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22. The reaction rate will therefore depend upon the proportion of molecules which have this or higher energy.
KLB PG 80
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23. Things that Affect Rate
Temperature
Higher temperature, faster particles.
More and harder collisions.
Faster Reactions.
Concentration
More concentrated molecules closer together
Collide more often.
Faster reaction.
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24. Collision theory and activation energy
Collision theory states that for a reaction to occur between reactants A and B to product C, portions of A and B must collide with sufficient energy to react.
Many collisions occur between reactant particles but only those with sufficient energy will result in a reaction.
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25. The sufficient energy is the activation energy.
The minimum amount of energy that the particles or reactants must have in order to react is called the activation energy.
Activation energy is an energy barrier which must be overcome before a reaction can take place.
The collisions which succeed in producing a reaction are called effective collisions.
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26. Non active particles can be activated by absorbing some energy.
The particles having activation energy(Ea) are known as activated particles.
Non active particles can be activated by absorbing some energy.
During a reaction, increase in temperature results in an increase in kinetic energy of particles making particles to move faster and collide more frequently with sufficient energy to cause more effective collisions per unit time.
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27. High concentration of particles of the reactants increases the number of collisions which increases the number of successful collisions and hence increase in rate of a reaction.
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28. Energy
Reactants
Products
Reaction coordinate
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29. Activation Energy - Minimum energy to make the reaction happen.
Reactants
Products
Reaction coordinate
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30. Activated Complex or Transition State
Energy
Reactants
Products
Reaction coordinate
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31. 3. Surface area
Is significant in heterogenous reaction where the reactants are not in the same phase.
Increasing surface area of a solid reactant increases reaction rate.
The reaction
CaCO₃ (s) + 2HCl (aq) CaCl₂ (aq) +H₂O (l)+ CO₂ (g)
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32. Increase in the surface area of particles increases the area of contact of the reacting particles.
The greater the surface area of contact between reacting particles, the higher the rate of a reaction.
The rate of a reaction of powdered marble chips with acid reaches completion earlier than unground marble.
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33. NB: Both curves flatten out at the same volume.
CO₂ produced is the same in both experiment.
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34. 1/2/2019
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35. 4. Catalysts
substances that alters the rate of a chemical reaction but remains chemically unchanged at the end of a reaction.
Catalysts provide an easier route for the reactants to attain the minimum energy required for a reaction to take place.
The new path has a lower activation energy.
Most catalysts work by providing a surface on which particles can react hence solid catalysts are effective in powder form e.g. in contact process.
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36. Other catalysts act by adsorption.
Catalysts can also increase the rate of a reaction by having a short lived intermediate compound or complex(activated complex) which involves the catalyst and which breaks up to release the catalyst and the products.
Other catalysts act by adsorption.
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37. NB:
Energy of the reactants and products remains the same but the activation energy is lowered.
Since catalysts are unchanged by the end of the reaction, they can be re used.
Some common catalysts
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38. The heat change for the reaction (∆H) remains the same
catalyst
Industrial process
Iron
Manufacture of ammonia in haber process
Vanadium(V) oxide; platinum
Manufacture of sulphuric (VI) acid.
Enzymes (yeast)
Production of alcoholic drinks by fermentation of sugar
The heat change for the reaction (∆H) remains the same
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39. Energy
Reactants
Products
Reaction coordinate
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40. Catalyst
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41. Catalysts Hydrogen bonds to surface of metal. Break H-H bonds
Pt surface
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42. Catalysts H H C C H H H H H H
Pt surface
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43. Catalysts C C The double bond breaks and bonds to the catalyst.
Pt surface
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44. Catalysts C C The hydrogen atoms bond with the carbon Pt surface H H H
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45. Catalysts H H C C H H H H H H
Pt surface
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46. 5. Light
Some reactions speed up considerably when exposed to sunlight.
In such reactions, molecules absorb energy in form of light energy rather than heat energy and so they are called photochemical reactions.
The ultra violet part of the sunlight provides energy that speeds up the reaction.
The influence of light on a chemical reaction is called photocatalysis
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47. Examples
Reaction between chlorine and hydrogen does not occur under ordinary temperature and pressure in darkness.
When exposed to direct light, the reaction is explosive.
Cl₂(g) + H₂(g) HCl(aq)
U.V light
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48. ii. Reaction between bromine and methane CH₄ (g) + Br₂ (g) CH₃Br (g) + HBr (g) iii. Photosynthesis
U.V light
(g)
(l)
(g)
(aq)
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49. NB Photographic plates are normally coated with silver bromide.
Processing of films is done in a dark room to avoid decomposition of the silver bromide.
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50. Effect of pressure of a gas on the rate of reactions
Pressure is significant in reactions involving gaseous reactants.
Increase in pressure reduces the volume means the concentration of particles increases, the collision rate therefore increases consequently increasing the rate of reaction.
When pressure decreases, the volume increases hence reducing concentration of particles and in effect reducing the rate of a reaction.
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51. 1/2/2019
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52. Reversible physical changes
Reversible reactions are those reactions which can be made to go either direction by changing conditions such as temperature.
Reversible physical changes
These involve change of state only e.g. from solid to liquid which a solid melts or from liquid to gas when a liquid evaporates.
Ice liquid steam
Heat
cool
Heat
cool
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53. I₂ (s) I₂ (g) (sublimation) where no liquid is formed.
Heat
cool
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54. Reversible chemical reactions (changes)
Expt klb pg 92
Heating hydrated copper (II) sulphate.
Reversible chemical reactions are those in which products, under certain conditions can be combined to reform to reactants.
When hydrated copper (II) sulphate crystals are heated, a white powder is formed.
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55. CuSO₄.5H₂O (s) CuSO₄ (s) + 5H₂O (l)
blue crystals white powder colourless
liquid
When the colourless liquid is poured back into the tube containing the white solid( anhydrous copper(II)sulphate turns blue and crystalline structure of hydrated copper (II) sulphate is restored.
CuSO₄.5H₂O (s) CuSO₄ (s) + 5H₂O (l)
Heat
cool
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56. Where is the sign for the reaction which goes both ways
The left to right reaction is referred to as forward reaction.
The right to left reaction is the reverse reaction.
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57. Question
Hydrated cobalt(II)chloride exist as pink crystals and anhydrous cobalt(II) chloride is a blue powder. Describe a laboratory experiment that can be used to show that the action of heat on hydrated cobalt(II)chloride is a reversible reaction.(KCSE 2010)
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58. Answer
Heat the hydrated salt, the pink substance changes to blue, collect vapour and cool, add liquid to solid, it turns pink.
Heat the hydrated salt in a sealed container, the pink substances changes to blue; allow the blue substance to cool.It changes to pink substance;
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59. Characteristics of reversible reactions
At the beginning of a reaction.
The rate of forward reaction is faster
The concentration of reactants is greater.
The concentration of products is lowest.
At zero time, the rate of reverse reaction is zero, because no products are formed.
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60. The rate of forward reaction decreases with time .
The rate of reverse reaction and the concentration of the products increases with time.
A time is reached when forward and reverse rates are equal.
Other examples of reversible reactions.
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61. COCl₂.6H₂O (s) COCl₂ (s) + 6H₂O (g) pink blue 3Fe (s) + 4H₂O (g) Fe₃O₄ (s) + 4H₂ (g) grey brown CaCO₃ (s) CaO (s) + CO₂ (g) 2SO₂ (g) + O₂ (g) 2SO₃ (g) N₂ (g) + 3H₂ (g) 2NH₃ (g)
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62. 2HI (g) H₂ (g) + I₂ (g) NH₄Cl (s) NH₃ (g) + HCl (g) N₂ (g) + O₂ (g) 2NO (g) 2HgO (s) 2Hg (l) + O₂ (g)
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63. EQUILIBRIUM
In a reversible chemical reaction the two opposing reactions(forward and reverse reactions) eventually reach a point where no changes in the reactant mixture seem to be occuring.
Rate of forward reaction = rate of reverse reaction.
At this stage, both forward and reverse reactions are taking place at the same rate but in opposite directions.
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64. The reaction is said to have attained “ a state of balance”
Microscopic (invisible) processes continue but in a balance that yields macroscopic(visible) changes.
The reaction is said to have attained “ a state of balance”
Forward and reverse reactions are therefore in a state of equilibrium.
A reaction in equilibrium is represented by the sign
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65. When you first put reactants together the forward reaction starts.
Since there are no products there is no reverse reaction.
As the forward reaction proceeds the reactants are used up so the forward reaction slows.
The products build up, and the reverse reaction speeds up.
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66. This is dynamic equilibrium.
Eventually you reach a point where the reverse reaction is going as fast as the forward reaction.
This is dynamic equilibrium.
The rate of the forward reaction is equal to the rate of the reverse reaction.
The concentration of products and reactants stays the same, but the reactions are still going on
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67. Examples of dynamic equilibrium. Acid base equilibrium
Chromate dichromate equilibrium
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68. Factors affecting equilibrium
The balance or position of equilibrium may be altered by most of the factors that affect the rate of reaction.
They include:
Concentration of the components
Temperature
Pressure
Use of catalysts
Volume of the vessel
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69. When a system at equilibrium is suddenly disturbed, it will respond in such a way that the equilibrium is eventually re-established.
This is in accordance with the le chartelier’s principle which states that;
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70. Le Châtelier’s Principle
States that
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71. If a system at equilibrium is subjected to change, the processes occur which tend to counteract the change;
Or if the stress is applied to a system at equilibrium, the system re-adjusts, if possible to reduce the effect of stress.
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72. Three types of stress are applied. Changing concentration
Changing temperature
Changing pressure
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73. Effect of change of concentation on equilibrium
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74. Changing Concentration
If you add reactants (or increase their concentration).
The forward reaction will speed up.
More product will form.
Equilibrium “Shifts to the right”
Reactants ® products
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75. consider
Br₂ (g) +H₂O (l) OBr⁻ (aq)+ Br⁻ (aq) + 2H⁺ (aq)
When NaOH is added to bromine water(above reaction), it provides hydroxide ions which react with hydrogen ions in solution to form water.
OH⁻ (aq) H⁺ (aq) H₂O (l)
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76. Consequently hydrogen ions are removed from the mixture.
The reaction proceeds to the right( forward reaction is favoured).
Colour changes from yellow –orange to colourless.
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77. Changing Concentration
If you add products (or increase their concentration).
The reverse reaction will speed up.
More reactant will form.
Equilibrium “Shifts to the left”
Reactants ¬ products
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78. Consider
When hydrochloric acid is added to bromine water, hydrogen ions are introduced which react with colorless bromide and hypobromite ions to form yellow orange aqeuous bromine.
The reaction proceeds more to the left hence the increase in intensity of the colour( reverse reaction is favoured)
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79. chromate(VII) dichromate(VI) equilibrium
2K₂CrO₄ (aq) + H₂SO₄ (aq) K₂Cr₂O₇ (aq) + K₂SO₄ (aq) + H₂O (l)
yellow orange
Potassium chromate(VII) potassium dichromate(VI)
2CrO₄²ˉ (aq)+ 2H⁺(aq) Cr₂O₇²ˉ (aq) + H₂O (l)
yellow orange
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80. Cr₂O₇²ˉ (aq)+ OHˉ (a) 2CrO₄²ˉ (aq)+H₂O (l) Orange yellow
Addition of acid causes more chromate (CrO₄²ˉ) to change to dichromate(Cr₂O₇²ˉ) ions making the solutions to become orange.
Addition of hydroxide ions reverses the reaction leading to the formation of yellow chromate solution.
Cr₂O₇²ˉ (aq)+ OHˉ (a) CrO₄²ˉ (aq)+H₂O (l)
Orange yellow
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81. Changing Concentration
If you remove products (or decrease their concentration).
The reverse reaction will slow down.
More product will form.
Equilibrium reverse“Shifts to the right”
Reactants ® products
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82. Changing Concentration
If you remove reactants (or decrease their concentration).
The forward reaction will slow down.
More reactant will form.
Equilibrium “Shifts to the left”.
Reactants ¬ products
Used to control how much yield you get from a chemical reaction.
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83. Effect of temperature on equilibrium
if the temperature is increased, energy is added. To release the stress of extra energy, the equilibrium will shift to produce substances which are formed by absorbing energy.
If the temperature of the system in equilibrium is increased, the reaction will move in away that heat is absorbed i.e. the endothermic reaction is favoured.
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84. ( Dinitrogen tetraoxide) Nitrogen(IV)oxide
If the temperature of the system in equilibrium is lowered, the reaction will move in a way to produce more heat i.e. the exothermic reaction is favoured.
E.g.
N₂O₄ (g) NO₂ (g)
pale yellow dark brown
( Dinitrogen tetraoxide) Nitrogen(IV)oxide
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85. The equilibrium shifts from left to right.
The favoured reaction in the above case is endothermic hence is favoured by increase in temperature i.e. heating tends to more dark-brown colour.
The equilibrium shifts from left to right.
The reverse reaction is exothermic thus it will be favoured by decrease in temperature i.e. cooling the mixture shifts the equilibrium to the left and more pale yellow colour is observed.
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86. Summary of effect of temperature on equilibrium.
Consider the reaction
A + B C + D………….. ∆H ± Ve
Summary of effect of temperature on equilibrium.
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87. Effect of change in temperature on equilibrium
Reaction
Effect of change in temperature on equilibrium
Increase
Decrease
Exothermic
∆H -Ve
Shifts to the left and more of A and B would be formed.
Shifts to the right and more C and D would be formed.
Endothermic
∆H +Ve
Shifts to the right and more of C and D would be formed.
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88. Effect of pressure on equilibrium.
Change in pressure will only affect equilibrium involving gases.
If the pressure of a system in equilibrium is increased, the system will favour the direction that has fewer molecules.
A decrease in pressure favours the direction of the reaction that forms more molecules.
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89. 2NO₂ (g) N₂O₄ (g) dark brown pale yellow
when the volume of the gases is reduced by increasing the pressure, a pale yellow colour is observed.
This shows increase in pressure favours forward reaction which forms fewer molecules.
When the volume is increased by lowering the pressure, a dark brown colour is observed.
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90. The reverse reaction occurs which forms more molecules by lowering the pressure.
If the number of molecules of reactants equals number of molecules of the product, a change in pressure would have no effect on the equilibrium.
H₂ (g) I₂ (g) HI (g)
1 mol mol mol
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91. In the above reaction, change in pressure will not affect the equilibrium position but will help in quick attainment of the equilibrium.
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92. Effect of pressure change on equilibrium mixture
Reaction
Effect of pressure change on equilibrium mixture
N₂O₄(g) NO₂(g)
Increase
Decrease
More N₂O₄ formed ( reverse rxn favoured)
More NO₂ formed (forward rxn favoured)
N₂(g) + 3H₂(g) NH₃(g)
2SO₂(g)+ O₂(g) SO₃(g)
4NH₃(g) + 5O₂(g) NO(g)+ 6H₂O (g)
H₂(g) + Cl₂(g) 2HCl(g)
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93. some industrial applications of chemical equilibrium.
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94. Interests of a manufacturer
Maximum quantity of products
Minimum cost
Shortest time possible.
Optimum conditions are conditions required to obtain the greatest yield of the products in the shortest time and minimum cost.
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95. 1. Haber process N₂ (g) + 3H₂ (g) 2NH₂ ∆H = - 92kj/mol
Temperature 450˚C
Pressure 200 atm
Catalyst – iron (platimum is expensive)
Reducing the concentration of ammonia as soon as it is formed.
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96. 2. Contact process 2SO₂ (g) + O₂ (g) 2SO₃ (g) temperature 450˚C
Pressure – atmospheric pressure which gives 96% yield
Catalyst – vanadiun(V)oxide.
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97. Effect of catalyst on equilibrium
A catalyst has no effect on the position of equilibrium but is useful if it is required to reach the equilibrium quickly.
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98. Note Formation of gases and precipitates.
Ordinarily a reversible reaction does not go to completion as it takes place both in the forward and reverse directions.
However, if one of the products formed during the reaction is removed, reverse reaction will not take place but the forward reaction will go to completion.
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99. A reaction in which a gas evolved and allowed to escape or in which a precipitate is formed, will proceed in one direction only even though a reversible reaction might be expected.
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100. End
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