General, Organic, and Biochemistry, 8e

1. General, Organic, and Biochemistry, 8e Bettelheim,Brown, Campbell, and Farrell

2. Chapter 8 Reaction Rates and Chemical Equilibrium

3. Chemical Kinetics • Some chemical reactions takes place rapidly, others are very slowly. • Ex: AgNO3 + NaCl AgCl + NaNO3 fast • Ex : C6H12O6 + 6 O2 6 CO2 + 6 H2O v. slow • Chemical kinetics: the study of the rates of chemical reactions. • Every reaction has its own rate • Ex: Consider the reaction that takes place when chloromethane and sodium iodide are dissolved in acetone; the net ionic equation for this reaction is:

4. Chemical Kinetics • The rate of reaction is the increase in concentration of iodomethane divided by the time interval. • For example, the concentration of CH3I might increase from 0 to 0.12 mol/L over a 30 minute time period • The reaction rate over this period is: • Ex 8.1:for the above example, suppose the concentration of iodide was 0.24 mol/L at the start of the reaction. At the end of 20 min, the concentration dropped to 0.16 mol/L. what is the reaction rate? • The rate of reaction is the decrease in concentration of iodide divided by the time interval = 0.16 – 0.24/20 = - 0.004mol/L.min

5. Chemical Kinetics • The Rate of the reaction: • It is of great important to know the rates of chemical reaction in both laboratories and inside our bodies. The reaction that goes more slowly than we need may be useless, whereas a reaction that goes too fast may be dangerous.

6. Molecular Collisions • In order for two species, A and B (they may be molecules or ions), to react, they must firstly collide, and secondly the collision must be effective. • Not all collisions result in a reaction. • A collision that results in a reaction is called an effective collision. • There are two main reasons why some collisions are effective and others are not; 1-activation energy and 2-the relative orientations of the colliding particles. • Activation energy: the minimum energy required for a reaction to take place.

7. Molecular Collisions • In most chemical reactions, one or more covalent bonds must be broken and energy is required for this to happen. • This energy comes from the collision between A and B. • If the collision energy is large, there is sufficient energy to break the necessary bonds, and reaction takes place. • If the collision energy is too small, no reaction occurs.

8. Molecular Collisions • 2- Orientation at the time of collision • Even if two molecules colloid with an energy greater than activation energy, reaction may not take place if molecule are not oriented properly at the time of the reaction. For example, to be an effective collision between H2O and HCl molecules, the oxygen of H2O must collide with the H of HCl so that the new O-H bond can form and the H-Cl bond can break.

9. Energy Diagrams • The reaction of H2 and N2 to form ammonia is exothermic: • In this reaction, six covalent bonds are broken and six now ones formed: • Breaking a bond requires energy, and forming a bond releases energy. • In this reaction, the energy released in making the six new bonds is greater than the energy required to break the six original bonds; the reaction is exothermic.

10. Energy Diagrams • Energy diagram for an exothermic reaction.

11. Energy Diagrams • Energy diagram for an endothermic reaction.

12. Energy Diagrams • Transition state: it is unstable state, where an activated complex is formed (short life complex). • The transition state for the reaction between H2O and HCl probably looks like this, in which the new O-H bond is partially formed and the H-Cl bond is partially broken.

13. Factors Affecting Rate • The rates of chemical reactions are affected by the following factors: • 1-Nature of reactants • Reaction between ions in aqueous solution are very fast (activation energies are very low). • Reaction between covalent compounds, whether in water or another solvent, are slower (their activation energies are higher). • 2-Concentration • In most cases, reaction rate increases when the concentration of either or both reactants increases. • For many reactions, there is a direct relationship between concentration and reaction rate; when concentration doubles the rate doubles. Rate α [reactants] • 2 H2O2(l) 2 H2O(l) + O2(g), rate = K [H2O2]

14. Factors Affecting Rate • 3-Temperature • Nearly in all reactions, the rate increases as temperature increases, because of the following: • When temperature increases, molecules move faster (they have more kinetic energy), which means that they collide more frequently; • more frequent collisions mean more effective collision occures and higher reaction rates. • 4- Catalyst: • Temperature can increase the reaction rate, but not recommended all the time because it may increase the rate of unwanted reactions. • a catalyst is substance that increases the rate of a chemical reaction without itself being used up.

15. Factors Affecting Rate • Catalysts lower the activation energy and increase the reaction rate. • Each catalyst has its own way of providing an alternative path way. • There are two kinds of catalyst: 1- heterogeneous catalyst (Different phaseof the reactant, Pd, Ni, Pt in reac. between gases) 2- homogenous catalyst (same phase of the reactant, Example enzymes)

16. Factors Affecting Rate • Many catalysts provide a surface on which reactants can meet. • The reaction of ethylene with hydrogen is an exothermic reaction. • If these two reagents are mixed, there is no visible reaction even over long periods of time. • When they are mixed and shaken with a finely divided transition metal catalyst, such as Pd, Pt, or Ni, the reaction takes place readily at room temperature.

17. Reversible Reactions • Reversible reaction: a reaction that can be made to go in either direction: • If we mix CO and H2O in the gas phase at high temperature, CO2 and H2 are formed: • We can also make the reaction take place the other way by mixing CO2 and H2: • The reaction is reversible, the rate of forward reaction gradually decrease and the reverse rate gradually increase until the two rates become equal

18. Reversible Reactions • Equilibrium: a dynamic state in which the rate of the forward reaction is equal to the rate of the reverse reaction. • At equilibrium there is no change in concentration of either reactants or products. • Reaction, however, is still taking place; reactants are still being converted to products and products to reactants, but the rates of the two reactions are equal. • Equilibrium constant, K: the product of the concentration of products of a chemical equilibrium divided by the concentration of reactants, each raised to the power equal to its coefficient in the balanced chemical equation

19. Equilibrium Constants • for the general reaction: • the equilibrium constant expression is: • Problem 8.3: write the equilibrium constant expression for this reversible reaction: • solution: the equilibrium constant expression is:

20. Equilibrium Constants • Problem 8.4: write the K expression for: • O2 + 4 ClO2 2 Cl2O5 • Problem 8.5: when H2 and I2 react at 427°C, the following equilibrium is reached: • The equilibrium concentrations are [I2] = 0.42 mol/L, [H2] = 0.025 mol/L, and [HI] = 0.76 mol/L. Using these values, calculate the value of K. • Solution: This K has no units because molarities cancel.

21. Equilibrium and Rates • There is no relationship between a reaction rate and the value of K. • It is possible to have a large K and a slow rate as in reaction between glucose and oxygen • It is also possible to have a small K and a fast rate as in reaction of silver nitrate and sodium chloride. • It is also possible to have any combination of K and rate in between these two extremes. • Reaction with very large K value proceeds almost to completion (to the right) • N2 + 3 H2 2 NH3 K = 1018

22. LeChatelier’s Principle • LeChatelier’s Principle: when a stress is applied to a chemical system at equilibrium, the position of the equilibrium shifts in the direction to relieve the applied stress. • We look at five types of stress that can be applied to a chemical equilibrium: • addition of a reaction component • removal of a reaction component • change in temperature • Change in the pressure • Adding catalyst

23. LeChatelier’s Principle • Addition of a reaction component • suppose this reaction reaches equilibrium: • Suppose we now disturb the equilibrium by adding some acetic acid or ethanol, The rate of the forward reaction increases ( reaction shift to right) and the concentrations of ethyl acetate and water increase. As this happens, the rate of the reverse reaction also increases. In time, the two rates will again become equal and a new equilibrium will be established.

24. LeChatelier’s Principle • The system has relieved the stress by increasing the components on the other side of the equilibrium. • We say that the system has shifted to minimize the stress (reaction shift to right). • If we add H2O or ethyl acetate, the reaction shift to the left to minimize the stress by increasing the components of the other side of the equilibrium. • Ex 8.7: N2O4(g) 2 NO2(g) (colorless) (brown) When more N2O4 is added to the equilibrium mixture, the brown color become darker, why? More NO2 is formed because the addition of reactant shift the equilibrium to the right

25. LeChatelier’s Principle • Removal of a reaction component • Removal of a component shifts the position of equilibrium to the side that produces more of the component that has been removed. • Suppose we remove ethyl acetate from this equilibrium: • If ethyl acetate is removed, the position of equilibrium shifts to the right to produce more ethyl acetate and restore equilibrium. • The effect of removing a component is the opposite of adding one.

26. LeChatelier’s Principle • Problem 8.8: when acid rain (H2SO4(aq))attacks marble (calcium carbonate), the following equilibrium can be written: How does the fact that CO2 is a gas influence the equilibrium? • Solution: CO2 gas diffuses from the reaction site, and is removed from the equilibrium mixture; the equilibrium shifts to the right and the marble continues to erode.

27. LeChatelier’s Principle • Change in temperature • The effect of a change in temperature on an equilibrium depends on whether the forward reaction is exothermic or endothermic. • Consider this exothermic reaction: • Adding heat (increasing the temperature) as adding product, pushes the equilibrium to the left. • Removing heat (decreasing the temperature) pushes the equilibrium to the right. • Consider the endothermic reaction: 39 kcal + 2N2(g) + O2(g) 2 N2O(g)

28. LeChatelier’s Principle • Summary of the effects of change of temperature on a system in equilibrium • Ex 8.9: 2 NO2(g) N2O4(g) + 13700 cal • the conversion of nitrogen dioxide to di-nitrogen tetra oxide is an exothermic reaction. The brown color is darker at 50ºC than it is at 0ºC. Explain? To go from 0ºC to 50ºC we must add heat so reaction shift to left.

29. LeChatelier’s Principle • Change in pressure: • Change in the pressure influences the equilibrium only if one ore more components of the reaction is gas. Ex: N2O4(g) 2 NO2 • Increase the pressure will shift the reaction to less no of mole (to left) • Decrease the pressure will shift the reaction to high no of moles (to right) • Ex 8.10: N2(g) + 3 H2(g) 2NH3(g) • What kind of pressure change would increase the yield of ammonia? Increase pressure

30. LeChatelier’s principle • Effect of catalyst • Catalyst increase the reaction rate ( forward and reverse reactions) to reach equilibrium faster. Catalyst has no effect on the position of equilibrium.

31. Chapter 8 Reaction Ratesand Chemical Equilibrium End Chapter 8