Chemistry SM-1232 Week 10 Lesson 2

1. Chemistry SM-1232Week 10 Lesson 2 Dr. Jesse Reich Assistant Professor of Chemistry Massachusetts Maritime Academy Spring 2008

2. Class Today • We will meet for class this Friday • We’ll have next Friday off • I’ll give you a take home quiz after class on Friday • Tests and quizes are not yet graded. I’ll have them as quickly as I can. • Chemical Equilibrium, Dynamic Equilibrium, Equilibrium Constant • Disturbing Equilibrium: concentration change, volume change, temperature change • Ksp • Wiki project to be discussed on Friday!

3. Chemical Equilibrium • Chemical and physical changes can happen slowly or quickly. • Chemists want to control the speed of a reaction and the products of a reaction that form. • Chemists study a topic called equilibrium, which I think of as how reactions would balance on a tight rope.

4. Reaction Rates • Reactions that turn large amounts of reactants to products in short amounts of time are considered fast reactions. • Reactions that turn small amounts of reactants to products in long periods of time are considered slow reactions.

5. How Reactions Occur • Collision theory: • Chemical reactions occur by molecules of different type coming into physical contact with one another. • Not all collisions make products. Some need to have enough energy to react or else they just bounce off one another. • The amount of energy required to make a reaction occur is called “Activation Energy.”

6. Making reactions occur • Molecules react when they physically touch and when there is enough energy in them to react. • To control the speed of a reaction chemists have two main controls (like dials). • 1. Temperature: kinetic energy is relative to temperature so if we change temperature we can change if a molecule has enough energy to overcome activation energy • 2. Concentration. If we load up a reaction there will be lots of collisions. If the concentration is very small it becomes very unlikely two molecules will come in contact.

7. Gas Particle Simulator • Remember this simulator from last semester? • http://www.phy.ntnu.edu.tw/ntnujava/index.php?topic=632.0 • If we crank up the temp and “n” you should see a lot more collisions occuring.

8. Reactions going in both directions • If reactions only go in one direction you can see how high temp and concentration reactions would consume the starting materials very quickly. • Consider the reaction of H2 + I2 2HI • When they mix the products will form, but • 2HI  H2 + I2 • The back reaction occurs, so as we make HI and they collide we’ll reform H2 and I2

9. Equilibrium • If we let the bottle of hydrogen gas, iodine gas, and hydroiodic acid gas sit at high temperature eventually it will settle into a routine. The concentrations will shift until amount of HI gas created is exactly equal to the amount of HI gas reacted. • When the forward reaction and the backward reaction are equal we call it dynamic equilibrium. • We know dynamic equilibrium has been reached because concentrations no longer change.

10. Equilibrium: Don’t get confused • Just because the concentrations stop changing overall doesn’t mean the reaction stopped occurring! • Just because we reached equilibrium doesn’t mean concentrations will be equal.

11. Measuring Equilibrium • We can assign a number to describe the state of equilibrium in a reaction. The number itself is called the equilibrium constant. • The equilibrium constant is defined as the ratio at equilibrium of the products raised to their stoichiometric coefficients divided by the concentrations of the reactants raised to their stoichiometric coefficients.

12. Equilibrium constant • wA + xB yC + zD • w, x, y, and z are coefficients in a balanced reaction. • [A], and [B] are concentrations of the reactants. [C] and [D] are concentrations of the products.

13. Pure solids and pure liquids in Keq equations • DON’T INCLUDE THEM • 2 CO (g)  CO2(g) + C (s) • The Keq is just

14. Pure solids and liquids in Keq continued • Similarly, don’t include pure liquids • CO2 (g) + H2O(l)  H+ (aq) + HCO3-(aq)

15. Equilibrium Control • Chemists can control equilibrium by changing the “dials” available to us. • But it’s like sailing a boat. If you want to turn right you put the rudder in the opposite direction. • Le Chatelier’s Principle! When a chemical system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance

16. Le Chatelier’s Principle: concentration changes • N2O4 2 NO2 • If we add in more NO2 we’ll force the reaction further to the left. • If we add in more N2O4 we’ll force the reaction further to right

17. Le Chatelier’s Principle: concentration changes • 2 BrNO  2 NO + Br2 • What happens if we add in BrNO into the reaction chamber? • What happens if we add in Br2? NO?

18. Le Chatelier’s Principle: volume changes • Pressure and volume are inversely related • If there is more pressure there is less volume • If there is more volume there is less pressure • Consider N2 + 3H2 2NH3 • All gases take up the same amount of space, but count how many moles of gas are on the left side. Count how many are on the right side.

19. Le Chatelier’s Principle: Volume Changes • N2 + 3H2 2NH3 • If we increase the pressure the system wants to relieve the stress. • In order to relieve the stress from pressure we can shrink volume. • In which direction would the reaction shrink volume? So, if we increase pressure, which side of the reaction will be favored?

20. Le Chatelier’s Principle: Volume Changes • N2 + 3H2 2NH3 • If we increase the volume the system wants to relieve the stress. • In order to relieve the stress from a volume increase we can grow volume of gas. • In which direction would the reaction move if the goal was to increase gas volume?

21. Le Chatelier’s Principle: Changing Temperature • Exothermic • Out heat • A + B  C + D + HEAT • Endothermic • In Heat • A + B + HEAT  C + D

22. Le Chatelier’s Principle: Changing Temperature • A + B  C + D + HEAT • Which side will be favored if I add more heat in? • Need help think about condensing water: H2O(g)  H2O (l) + heat • A + B + HEAT  C + D • Which side will be favored if I take heat out? • Think about boiling water. • H2O(l) + heat  H2O (g)

23. E Chatelier’s Principle: Changing Temperature • N2 + 3H2 2 NH3 + Heat • What happens if you remove heat? • N2O4 + heat  2 NO2 • What happens if you remove heat?

24. Solubility • Solubility is an equilibrium expression too. • I told you there is a tug of war between the forces that would break up a solid and the forces that would hold it together • We have a number we can ascribe to talk about how much a solid dissolves

25. Ksp • Consider CaF2(s)  Ca2+(aq) + F-(aq) • Write the Keq • Now take away the solid component • We call that the Ksp

26. Ksp • CaF2(s)  Ca2+ (aq) + 2F-(aq) • You should have found Ksp= [Ca][F-]2 • Same rules apply here for interpreting the results. • Greater than one favors products • Less than one favors the reactants • In other words if it’s less than 1 it wants to stay a solid!

27. Ksp • We typically don’t look for Ksp of things that are very soluble, so you’ll almost always see number that are much smaller than one. • Write the Ksp for BaSO4, Mn(OH)2, and Ag2CrO4

28. To do • Read through the end of the chapter. • YOU MUST COPY EXAMPLE 15.9!!! Page 560! • This is very important! • There is also a very relevant story on maritime chemistry that I suggest you read up OUTSIDE OF CLASS on here: http://tinyurl.com/2tcnbl