Download
chapter 16 controlling the yield of reactions n.
Skip this Video
Loading SlideShow in 5 Seconds..
Chapter 16 – Controlling the Yield of Reactions PowerPoint Presentation
Download Presentation
Chapter 16 – Controlling the Yield of Reactions

Chapter 16 – Controlling the Yield of Reactions

116 Views Download Presentation
Download Presentation

Chapter 16 – Controlling the Yield of Reactions

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Chapter 16 – Controlling the Yield of Reactions

  2. Chemical Equilibrium • The stage when the quantities of reactants and products remains unchanged is called the chemical equilibrium. • Reaction yield is the extent of conversion of reactants into products, or how far the reaction will go.

  3. Reversible Reactions • Some physical and chemical changes can be reversed. • For example ice to liquid water and back again. • A double arrow ( ) is used to indicate a reversible reaction.

  4. Equilibrium Explained • We can use the idea that processes can be reversed to understand why some reactions reach an equilibrium. • Chemists have shown that in these reactions the forward and reverse reactions occur simultaneously.

  5. Equilibrium Explained Example • For example, nitrogen gas and hydrogen gas were added to a sealed container at a constant temperature. The nitrogen and hydrogen start to react immediately, forming ammonia. The following sequence of events then occurs: • As the forward reaction proceeds, the concentration of nitrogen and hydrogen decrease, so the rate of ammonia production decreases. • At the same time the ammonia is being formed, some ammonia molecules react to re-form nitrogen and hydrogen. The rate of this reverse reaction increases as the concentration of ammonia increases. • Eventually the forward and back reactions proceed at the same rate. When this situation is reached, the ammonia is formed at exactly the same rate it is breaking down. • The concentrations of the hydrogen, nitrogen and ammonia will remain constant. • At the equilibrium position no further change will take place in the rate of either the forward or back reaction. The reaction has reached a point of balance, an equilibrium.

  6. Equilibrium Explained cont…

  7. Equilibrium Explained cont… • Equilibrium is a dynamic state, since the forward and back reactions have not ceased. • They occur simultaneously at the same rate. • During dynamic equilibrium: • The amounts and concentrations of chemical substances remain constant. • The total gas pressure is constant (if gases are involved) • The temperature us constant • The reaction is incomplete (all of the substances are present in the equilibrium mixture).

  8. How far do equilibrium reactions go? • Different reactions proceed to different extents. • For example a strong acid, like HCl, almost completely ionises. • Whereas a weak acid, like ethanoic acid, does not completely ionise. • This is shown by their electrical conductivity. • As such their point of equilibrium is different.

  9. The Equilibrium Law • K is known as the equilibrium constant. • K can be found by finding the concentration fraction, also known as the reaction quotient. • While the concentration fraction can be calculated for any mixture of reactants and products, it is only at equilibrium that it has a constant value. K = [products] [reactants]

  10. Equilibrium Law Example… • N2(g) + 3H2(g) 2NH3(g) K = [NH3]2 [N2][H2]3 K= [0.074]2 [0.25][0.75]3 K = 0.005 0.25x0.42 K = 0.005 0.11 K = 0.048

  11. Equilibrium Law cont… • From studies of equilibria, chemists have found that: • Different chemical reactions have different values of K • For a particular reaction, K is constant for all equilibrium mixtures at a fixed temperature.

  12. What does an equilibrium constant tell us? • The equilibrium constant gives us an indication of the extent of the reaction. • For values of K that are between 10-4 and104, there will be significant amounts of reactants and products present at equilibrium. • For values of K that are large (>104), the equilibrium mixture consists mostly of products with small amounts of reactants. • For values of K that are small (<10-4) the equilibrium mixture will consist mostly of reactants with small amounts of products.

  13. Effect of Temperature on Equilibria • The value of K, for a particular reaction, depends only on temperature. • It is not affected by addition of reactants or products, changes in pressure or the use of catalysts. • As temperature increases: • For exothermic reaction, the amount of products decreases so the value of K decreases. • For endothermic reactions, the amount of products increases and so the value of K increases.

  14. Effect of Temperature on Equilibria

  15. Calculations using equilibrium constants Calculate the value of the equilibrium constant for the reaction represented by the equation: H2(g) + I2(g) 2HI(g) At 460°C, if a 2.00L vessel contains an equilibrium mixture of 0.0860mol of H2, 0.124 mole of I2 and 0.716mol of HI. • [H2] = n(H2)/v(H2) = 0.0860/2.00 = 0.0430M • [I2] = n(I2)/v(I2) = 0.124/2.00 = 0.0620M • [HI] = n(HI)/v(HI) = 0.716/2.00 = 0.358M • K = [HI]2 / [H2][I2] = 0.3582/(0.0430x0.0620) = 48.1

  16. Example 2 At a particular temperature 0.0500mol of SO2, 0.0100mol of O2 and 0.1500mol of SO3 were mixed in a 2.00L vessel and allowed to reach equilibrium according to the equation: 2SO2(g) + O2(g) 2SO3(g) Analysis showed that 0.1400mol of SO3 were present in the gas mixture at equilibrium . Calculate the value of the equilibrium constant at this temperature. Because the amount of SO3 has decreased from 0.1500 to 0.1400, a net back reaction must have occurred. The amount of SO3 that has reacted is 0.1500-0.1400 = 0.0100mol. From the equation, 2 mol of SO3 decomposes to 2 mol of SO2 and 1 mol of O2. Therefore, 0.0100mol of SO3 will form 0.100mol of SO2 and 0.00500 mol of O2.

  17. Example 2 cont… So at equilibrium: n(SO2) = 0.0100 + 0.0500 = 0.0600mol n(O2) = 0.00500 + 0.0100 = 0.0150mol n(SO3) = 0.1400mol [SO2] = n(SO2)/V(SO2) = 0.0600/2.00 = 0/0300M [O2] = 0.0150/2.00 = 0.00750M [SO3] = 0.1400/2.00 = 0.0700M K = [SO3]2 / [SO2]2[O2] = 0.07002 / 0.03002 x 0.00750 = 726

  18. Changing the Equilibrium Position of a Reaction • The equilibrium position of a reaction may be changed by: • Adding or removing a reactant or product • Changing the pressure by changing the volume (equilibria involving gases) • Dilution (for equilibria in solution) • Changing the temperature

  19. Adding Extra Reactant or Product • Suppose a vessel contains an equilibrium mixture represented by the equation: N2(g) + 3H2(g) 2NH3(g) • If extra nitrogen gas were added to the vessel without changing the volume or temperature, the mixture would momentarily not be in equilibrium. • The following events occur as the composition of the mixture adjusts to return to the equilibrium: • The increased concentration of nitrogen gas causes the rate of the forward reaction to increase and more ammonia is formed. • As the concentration of the ammonia increases, the rate of the back reaction to reform N2 and H2 increases. • Ultimately the rates of the forward and back reaction become equal again and a new equilibrium is established.

  20. Adding Extra Reactant or Product

  21. Adding Extra Reactant or Product • When equilibrium is re-established the concentration of all substances have changed. • The overall effect of adding nitrogen is to increase the concentration of ammonia at equilibrium – a net forward reaction. • Addition of more NH3, increases the rate of the back reaction.

  22. Adding Extra Reactant or Product • To summarise: • Addition of a reactant leads to the formation of more products (a net forward reaction). • Addition of a product leads to the formation of more reactants (a net back reaction). Le Chatelier’s Principle: if an equilibrium system is subjected to a change, the system will adjust itself to partially oppose the effect of the change.

  23. Changing the Pressure – by changing the volume • The pressure of gases in an equilibrium mixture can be changed by increasing or decreasing the volume of the container while keeping the temperature constant.

  24. Example 1… • During sulfuric acid manufacture, one step involves sulfur dioxide reacting with oxygen to form sulfur trioxide gas: 2SO2(g) + O2(g) 2SO3(g) 3 gas particles 2 gas particles • In this equilbrium, the forward reaction involves a reduction in the number of particles of gas, causing a reduction in pressure. An equilibrium will respond to an increase in pressure by adjusting to reduce the pressure. A net forward reaction will occur, increasing the amount of sulfur trioxide present at equilibrium.

  25. Example 2… • Colourless dinitrogen tetroxide gas and brown nitrogen gas exist in equilibrium: N2O4 2NO2(g) 1 gas particle 2 gas particles • When an equilibrium mixture of the gases is compressed it is observed that, after an initial darkening because of the higher concentration of the brown gas, the colour of the gas mixture fades. Some nitrogen dioxide has converted to dinitrogen trioxide. The system adjusts to the increased pressure by undergoing a net back reaction: the equilbrium moves in the direction that produces fewer particles and reduces the pressure. Note the concentration of NO2 in the new equilbrium will be higher than in the initial equilbrium, but not as high as it would be if there had not been a net back reaction.

  26. Changing the Pressure by Changing the Volume cont… • In general the effect of a change in pressure, by changing the container volume, depends on the relative number of gas particles on both sides of the equation.

  27. Changing the Pressure by Adding an Inert Gas • The total pressure of an equilibrium mixture of gases may also be changed by adding a non-reacting gas such as helium, neon or argon. • Despite the increase in total pressure that occurs when this gas is added, the concentration of the individual chemicals involved in the equilibrium are not affected by the presence of the extra gas. • The system therefore stays in equilibrium and there is no net forward or back reaction.

  28. Dilution • The effect of diluting a solution by adding water is similar to changing the volume in gaseous equilibria. • Where possible, a net reaction occurs in the direction that produces the greater number of particles. • For example, dilution of a solution containing the equilibrium system: Fe3+(aq) + SCN-(aq) Fe(SCN)2+(aq) 2 particles in solN 1 particle in solN • Results in an increases in the amount of Fe3+ and SCN- ions. • In terms of Le Chatelier’s principle, a net back reaction increases the total concentration of particles in solution, offsetting the effect of dilution. • Despite the net back reaction, the concentration of Fe3+ and SCN- at the new equilibrium will be lower than their concentration prior to the dilution as the equilibrium shift partially offsets the change.

  29. Changing the Temperature • Heating increases the energy of the substances in the mixture. • Applying Le Chatelier’s Principle to an exothermic reaction, the reaction opposes an increase in energy by removing energy – that is a net backwards reaction occurs. • For an endothermic reaction, net forwards reactions occur as the temperature rises. • An increase in temperature in an equilibrium mixture reaction results in: • A net backward reaction (less products) for exothermic reactions • A net forward reaction (more products) for endothermic reactions.

  30. The Influence of Catalysts • Catalysts increase the rate of reactions. • It has been shown that they increase the rate of forward and back reactions equally. • As a consequence, the presence of catalyst does not change the position of equilibrium. • A catalyst may greatly increase the rate at which equilibrium is reached.

  31. Do all reaction reach equilibrium? • Many reactions can be regarded as continuing until they are complete. • These include: • Reactions that produce products such as gases that escape from the reaction mixture as they are formed. Continual loss of products drives these reactions forwards. • Reactions that form equilibria in which only minute quantities of reactants are present.

  32. Le Chatelier’s Principle Summarised