1. Collision Theory  Collision theory is a theory proposed independently by Max Trautz in 1916 and William Lewis in 1918, that qualitatively explains how chemical reactions occur and why reaction rates differ for different reactions.  Molecules/atoms first have to collide, and then they may react.  Affected by orientation, temperature, pressure, concentration, surface area and catalysts.

2. Rate of reactions  The rate of a reaction is the speed at which a chemical reaction happens.  If a reaction has a low rate, that means the molecules combine at a slower speed than a reaction with a high rate.  Rate is usually not constant throughout the reaction

3. Activation Energy  In chemistry, activation energy is a term introduced in 1889 by the Swedish scientist Svante Arrhenius that means the minimum energy that must be input to a chemical system with potential reactants to cause a chemical reaction.

4. Endothermic and Exothermic

5.  Endothermic: reactants absorb heat during the reaction  Exothermic: products include a release of heat during the reaction

6. Factors affecting reaction rates  An increase in temperature usually results in the atoms or molecules moving more rapidly and therefore are more likely to collide and increases the initial energy which will more likely meet or exceed the activation energy

7. Factors affecting reaction rates  An increase in concentration & pressure means more atoms or molecules of a reactant in a given volume of space which increases the probability they will collide and possibly react.  As pressure increases, the volume the gas takes up decreases and thus the concentration increases and the rate will also increase.

8. Factors affecting reaction rates  An increase in Surface area may increase the reaction rate if the reactants are solids or liquids  The reaction can only take places at the surface and thus the amount of surface area exposed will change the rate.

9. Factors affecting reaction rates  A catalyst increases in the rate of a chemical reaction due to the participation of an additional substance called a catalyst.  With a catalyst, reactions occur faster and with less energy.  Because catalysts are not consumed, they are recycled. Often only tiny amounts are required.

10. Dynamic Equilibrium  Many chemical reactions can easily run in both forward and reverse directions.  When there is a balance between the forward reaction and the reverse reaction, the system is in dynamic equilibrium

11. Reversible Reaction  Many chemical reactions can easily run in both forward and reverse directions and are called reversible reactions.  The forward reaction direction and the reverse reaction direction will run all the time.  Other chemical reactions can only run in one direction, going only from the reactants on the left side of the arrow to the products on the right side of the arrow.

12. Le Chatelier’s Principle  Le Chatelier’s Principle : If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.

13. Le Chatelier’s Principle  Suppose you have an equilibrium established between four substances A, B, C and D.  What would happen if you changed the conditions by increasing the concentration of A?  According to Le Chatelier, the position of equilibrium will move so that the concentration of A increases again. That means that more C and D will react to replace the A that has been removed. The position of equilibrium moves to the left.

14. Factors affecting reactions….  Concentration  Pressure  Temperature

15. HEAT – Changing Temperature  If temperature is increased, that means that the position of equilibrium will move so that the temperature is reduced again.  To determine the effect of heat, you need to know if the reaction is exothermic (given out) or endothermic (absorbed).

16. HEAT – Changing Temperature  Assume that our forward reaction is exothermic (heat is evolved):  The value ∆H ( ∆Q in physics) is always given as if the reaction was one-way in the forward direction.  The negative means heat was released (also called evolved )  A positive means heat was absorbed.

17. HEAT – Changing Temperature  Suppose the system is in equilibrium at 300°C, and you increase the temperature to 500°C.  How can it cool itself down again?  To cool down, it needs to absorb the extra heat that you have just put in.  In the case we are looking at, the back reaction is ENDOTHERMIC (absorbs heat). The position of equilibrium therefore moves to the left. The new equilibrium mixture contains more A and B, and less C and D.

18. Example:  2SO 2 (g) + O 2 (g) ↔ 2SO 3 (g) + heat  Is  H (  Q) positive or negative in this reaction?  Adding heat will cause it to ABSORB heat (←)  Removing heat will cause it to release heat (→)

19. Effect of Pressure  This only applies to reactions involving gases  Pressure is caused by gas molecules hitting the sides of their container. The more molecules you have in the container, the higher the pressure will be. The system can reduce the pressure by reacting in such a way as to produce fewer molecules.

20. Example  N 2 (g) + 3H 2 (g) ↔ 2NH 3 (g)  There are 4 molecules on the left and 2 on the right. The forward reaction will reduce molecules, the reverse will increase  Adding pressure will cause the system to REDUCE molecules (→)  Removing pressure will cause it to increase molecules (←)

21. Effect of Concentration  Earlier we said that if we increase the concentration of a species, the reaction will try to reverse that increase.  PCl 5 (g) ↔ PCl 3 (g) + Cl 2 (g)  Adding PCl 5 will cause the system to produce less PCl 5 (→)  Removing PCl 5 will cause the system to produce more PCl 5 (←)

22. System in Equilibrium  The following system is at equilibrium. How would each of the following changes shift the position of the equilibrium?  CO (g) + 2H 2 (g) ↔ CH 3 OH (g) ∆H = -91 kJ  What happens if…..  Raise the temperature.  Increase the size of the container.  Increase pressure.  Add solid Fe mesh to absorb CO(g).  Add more H 2 (g).  Add more Carbon Monoxide

23. Equilibrium Constants Calculations ~ K eq  The Equilibrium Constant, K c or K eq, relates to a chemical reaction at equilibrium.  It can be calculated if the equilibrium concentration of each reactant and product in a reaction at equilibrium is known.

24. Equilibrium Constants Calculations ~ K eq  Below is a “generic” balanced chemical equation: aA + bB  cC + dD  Lower case letter represent the coefficients.  Upper case letters represent the molecules/atoms  The equation for equilibrium constant  Brackets mean concentration in MOLARITY

25. More Equilibrium Constant  Pure solids and liquids should be “excluded” from the equilibrium expression by placing a “1” If K >1 products are favored at equilibrium If K < 1 reactants are favored at equilibrium  A Homogenous equilibrium has everything present in the same phase. The usual examples include reactions where everything is a gas, or everything is present in the same solution.  A Heterogeneous equilibrium has things present in more than one phase. The usual examples include reactions involving solids and gases, or solids and liquids.

26. Examples  Gaseous Dinitrogen Tetroxide is in equilibrium with Gaseous Nitrogen Dioxide. Write and balance this equation, then write the equilibrium expression.

27. Example (cont’d)  Calculate K eq for the reaction in question 2 above when [SO 3 ] = 0.016M, [SO 2 ] = 0.0056M, and [O 2 ] = 0.0021M.  Does this reaction favor reactants or products?

28. Equilibrium-Pressure  You can also write an equation using the pressures of the reactants and products.  The pressure constant and the equilibrium constant are related. R = 0.0821 atm*liter / mol*K

29. Thermodynamics  Movement of heat energy  Specific Heat (or Specific Heat Capacity = C ) the amount of heat necessary to move 1.00 gram of a substance 1.00 °C  Heat Capacity is the amount of heat necessary to move the temperature 1.00 °C  DIFFERENT SOURCES DEFINE THESE DIFFERENTLY – I AM SORRY….SCIENCE IS WEIRD

30. Thermodynamics  Q = Heat Energy in Joules or calories  Specific Heat Capacity of Water  1 cal/g °C  4.184 J/g °C

31. Thermodynamics  1. A hot iron bar is thrust into 200-mL of water at 25 °C. If the water rises to 42 °C, what is the amount of heat energy gained by this endothermic reaction?

32. Thermodynamics  2. In an exothermic reaction, 50-kilocalories are emitted from a hot brick having a heat capacity of 1.32 cal/g °C. What was the mass of the brick if it began at 300 °C and fell to 45 °C?

34. Energy for Phase Changes  Heat of Fusion (H f ) or Molar Heat of Fusion is the energy absorbed to melt/or released to freeze per unit mass  Water H f = 6.01 kJ/mol  Water H f = 3.34 x 10 2 J/g  Heat of Vaporization (H v ) or Molar Heat of Vaporization is the energy absorbed to boil/or released to condense per unit mass  Water H v = 40.7kJ/mol  Water H v = 2.26 x 10 3 J/g

35. Energy for Phase Changes  If you start with 5.0g of ice at 0°C, how much energy must be absorbed to melt it all?  2. How many grams of ice can be melted with the addition of 1.5kJ of energy?  3. How many moles of water can be evaporated with the addition of 20kJ of energy?

36. Energy for Phase Changes

38. Phase Change Graph  Show on the graph the location of the boiling point temperature and the melting point temperature.  What phase of matter is missing? Name and describe it.  What do you note about temperature during the phase change?