To understand the collision model of chemical reactions To understand activation energy

1. Objectives 17.1 Reaction Rates and Equilibrium • To understand the collision model of chemical reactions • To understand activation energy • To understand how a catalyst speeds up a chemical reaction • To explore reactions with reactants or products in different phases • To learn how equilibrium is established • To learn about the characteristics of chemical equilibrium

2. Big Idea 4 Rates of chemical reactions are determined by the details of the molecular collisions

3. A. How Chemical Reactions Occur • Collision model – molecules must collide in order for a reaction to occur • Rate depends on concentrations of reactants and temperature.

4. A. How Chemical Reactions Occur • Many reactions proceed via a series of elementary reactions • Elementary reactions are mediated by collisions between molecules. Only collisions having sufficient energy and relative orientation of reactants lead to products

5. B. Conditions That Affect Reaction Rates • Concentration – increases rate because more molecules lead to more collisions • Temperature – increases rate • Why?

6. B. Conditions That Affect Reaction Rates • Reaction rates that depend on temperature and other environmental factors are determined by measuring changes in concentrations of reactants or products over time

7. B. Conditions That Affect Reaction Rates • Activation energy – minimum energy required for a reaction to occur

8. B. Conditions That Affect Reaction Rates • Catalyst – a substance that speeds up a reaction without being consumed • Enzyme – catalyst in a biological system

9. Catalyst Reaction rates may be increased by the presence of a catalyst • Catalysts function by lowering the activation energy of an elementary step in a reaction mechanism, and by providing a new and faster reaction mechanism • Important classes in catalyst include acid-base catalysis, surface catalysis, and enzyme catalysis

10. C. Heterogeneous Reactions • Homogeneous reaction – all reactants and products are in one phase • Gas • Solution • Heterogeneous reaction – reactants in two phases

11. C. Heterogeneous Reactions

12. D. The Equilibrium Condition • Equilibrium – the exact balancing of two processes, one of which is the opposite of the other

13. D. The Equilibrium Condition • Chemical equilibrium – a dynamic state where the concentrations of all reactants and products remain constant

14. Although time continues to pass, the numbers of reactant and product molecules are the same as in (c). No further changes are seen as time continues to pass. The system has reached equilibrium. The reaction continues as time passes and more reactants are changed to products. The reaction begins to occur, and some products (H2 and CO2) are formed. Equal numbers of moles of H2O and CO are mixed in a closed container. E. Chemical Equilibrium: A Dynamic Condition

15. E. Chemical Equilibrium: A Dynamic Condition • Why does equilibrium occur?

16. Objectives 17.2 Characteristics of Equilibrium • To understand the law of chemical equilibrium • To learn to calculate values for the equilibrium constant • To understand how the presence of solids or liquids affects the equilibrium expression   

17. Big Idea 6 Any bond or intermolecular attraction that can be formed can be broken. These two processes are in a dynamic competition, sensitive to initial conditions and external perturbations.

18. A. The Equilibrium Constant: An Introduction • Law of chemical equilibrium • For a reaction of the type • aA + bB↔cC + dD • Equilibrium expression • Each set of equilibrium concentrations is called an equilibrium position.

19. A. The Equilibrium Constant: An Introduction • Chemical equilibrium is a dynamic, reversible state in which rates of opposing processes are equal • Systems at equilibrium are responsive to external perturbations, with the response leading to a change in the composition of the system

20. B. Heterogeneous Equilibria • Heterogeneous equilibria – an equilibrium system where the products and reactants are not all in the same state

21. Objectives 17.3 Applications of Equilibria • To learn to predict the changes that occur when a system at equilibrium is disturbed • To learn to calculate equilibrium concentrations • To learn to calculate the solubility product of a salt • To learn to calculate solubility from the solubility product

22. A. Le Chatelier’s Principle • Le Chatelier’s Principle – when a change is imposed on a system at equilibrium the position of the equilibrium shifts in a direction that tends to reduce the effect of that change

23. A. Le Chatelier’s Principle • Effect of a Change in Concentration

24. A. Le Chatelier’s Principle • Effect of a Change in Concentration • When a reactant or product is added the system shifts away from that added component. • If a reactant or product is removed, the system shifts toward the removed component.

25. The system is initially at equilibrium. The piston is pushed in, decreasing the volume and increasing the pressure. The system shifts in the direction that consumes CO2 molecules, lowering the pressure again. A. Le Chatelier’s Principle • Effect of a Change in Volume

26. A. Le Chatelier’s Principle • Effect of a Change in Volume • Decreasing the volume • The system shifts in the direction that gives the fewest number of gas molecules.

27. A. Le Chatelier’s Principle • Effect of a Change in Volume • Increasing the volume • The system shifts in the direction that increases its pressure.

28. A. Le Chatelier’s Principle • Effect of a Change in Temperature • The value of K changes with temperature. We can use this to predict the direction of this change. • Exothermic reaction – produces heat (heat is a product) • Adding energy shifts the equilibrium to the left (away from the heat term). • Endothermic reaction – absorbs energy (heat is a reactant) • Adding energy shifts the equilibrium to the right (away from the heat term).

29. B. Applications Involving the Equilibrium Constant • The Meaning of K • K > 1  the equilibrium position is far to the right • K < 1  the equilibrium position is far to the left

30. B. Applications Involving the Equilibrium Constant • The value of K for a system can be calculated from a known set of equilibrium concentrations. • Unknown equilibrium concentrations can be calculated if the value of K and the remaining equilibrium concentrations are known.

31. C. Solubility Equilibria • The equilibrium conditions also applies to a saturated solution containing excess solid, MX(s). • Ksp = [M+][X] = solubility product constant • The value of the Ksp can be calculated from the measured solubility of MX(s).

32. C. Solubility Equilibria • Chemical equilibrium plays an important role in acid-base chemistry and in solubility • The equilibrium constant is related to temperature and the difference in Gibbs free energy between reactants and products