1. 1 Kinetics and Equilibrium

2. 2 All substances contain chemical potential energy A  B high PE low PE Energy released -- Exothermic low PE high PEEnergy absorbed -- Endothermic Most chemical reactions require some energy to get started -- activation energy

3. 3 Potential Energy Diagrams Progress of Reaction PE A – PE of reactants B – Activation energy C – PE of products D – Heat of reaction; ΔH In this diagram, PE decreasedenergy released; exothermic ΔH is negative A B C D 50 40 30 20 10 kJ

4. 4 Exothermic reactions have a negative ΔH Endothermic reactions have a positive ΔH See reference table I

5. 5 How do we tell if a reaction will occur spontaneously? Consider 2 factors: 1. Heat of Reaction (ΔH) --low energy conditions are more stable -- exothermic reactions (- ΔH ) are favored 2. Change in entropy (ΔS) --entropy is a measure of disorder, or randomness --increasing entropy (+ ΔS ) is favored gasliquidsolid disordered neat, orderly high entropylow entropy

6. 6 Favorable conditions are: Exothermic -ΔH Increasing entropy +ΔS Any reaction with -ΔH, +ΔS is always spontaneous +ΔH, -ΔS is never spontaneous Reactions where only one factor is favorable are spontaneous under some conditions – depends on temperature. Example: 2C(s) + 3H 2 (g) → C 2 H 6 (g) ΔH = -84.0 kJExothermic, favorable ΔS solid to gas, entropy increases, favorable So this reaction is spontaneous

7. 7 Reaction Rates Collision Theory --reactions occur when molecules physically collide -- the collision must be effective – it must have 1. enough energy 2. proper orientation -- any change that increases the number of effective collisions will increase the reaction rate

8. 8 Factors affecting reaction rate 1. Temperature --faster motion of molecules means more frequent and more effective collisions 2. Surface Area --more contact between reactants means more frequent collisions (increase surface area by dividing a solid into many small pieces) 3.Concentration of Reactants – higher concentration = more frequent collisions in solutions – raise concentration by dissolving more reactant, or by reducing water in gases – raise concentration by raising pressure

9. 9 4. Nature of Reactants -- --if bonds must be broken first, reaction is slow --if not, reaction is fast ex., reactions between solutions 5.Catalysts– a catalyst is a substance that increases the rate of a chemical reaction catalysts are not used up during the reaction catalysts do not affect the ΔH or the spontaneity of the reaction catalysts work by lowering the activation energy

10. 10 low energy collision ineffective no catalyst effective catalyst (no catalyst)

11. 11 Equilibrium -- a state of balance between opposing forces -- no overall changes occur at equilibrium -- in a dynamic equilibrium, opposing changes occur at equal rates and cancel each other Kinds of Equilibrium 1. Phase equilibrium: ex.) ice and water at 0ºC ex.) water and water vapor in a closed jar 2. Solution equilibrium 3. Chemical equilibrium

12. 12 Solutions -- any homogeneous mixture Not to be confused with: a. suspensions: small particles of one substance suspended in another b. colloids: a suspension that doesn’t settle Parts of a Solution 1. Solute : the material that gets dissolved makes up less than 50% 2. Solvent : material that dissolves the solute makes up more than 50% Often the solvent is a liquid --but not always solvent solute

13. 13 Solution Concentration Qualitative:lots of solute per unit solvent = concentrated little solute per unit solvent = dilute 1.Molarity (M) moles per liter solute solution M = moles liters Quantitative:

14. 14 Example: 100 g of MgO are dissolved in enough water to make 5.0 L of solution. What is the molarity? Solution: first, find the moles of solutemoles = 100 40 = 2.5 moles next, use molarity formulaM = 2.5 5.0 = 0.5M Example: How many grams of BaCl 2 are needed to make 300.0 mL of a 0.70M solution? Solution: find the moles of solute required 0.7 = moles 0.3 = 0.21 moles now, change moles to grams 0.21 = grams 207 44 g =

15. 15 2. Percent (%) amount of solute. amount of total solution % = ( ) x 100 amounts can be volumes or masses, but units must agree -- except it’s OK to mix g and mL Example: 100 mL alcohol is diluted with enough water to form a final volume of 250 mL. % alcohol = ? Solution: 100 250 x 100 = 40% Example: 100 mL alcohol is mixed with 250 mL water. % alcohol=? Solution: 100 350 x 100 = 29%

16. 16 3. Parts per Million (ppm) ppm = amount of solute. amount of total solution ) x 1,000,000 ( Example: a 50.0 gram sample of apple contains 3.2 x 10 -4 g of pesticide. ppm = ? Solution: 3.2 x 10 -4 50.0 1,000,000 ()( ) = 6.4 ppm When solutes dissolve in liquids, they change some physical properties of the liquid, such as: 1. The freezing point is lowered -- salt used to melt ice on roads -- salt used to freeze home-made ice cream -- antifreeze used to prevent car freeze-ups

17. 17 2. The boiling point is raised --antifreeze also prevents boiling These changes depend on the concentration of solute particles, not what kind Some substances cause larger-than-expected changes -- ionic compounds (salts)-- acids (H + ) and bases (OH - ) These substances dissociate, or ionize, when dissolved in water (separate into + and – ions) Example: NaCl (s)  1 mole 1 mole 1 mole So, 1M NaCl affects the fp and bp as if it were a 2M solution 2 moles Examples: CaCl 2 (s)  Ca +2 (aq) + 2Cl - (aq) Na 2 CO 3 (s)  2Na + (aq) + CO 3 -2 (aq) Na + (aq) + Cl - (aq)

18. 18 These ions behave independently in solution Example: Na 2 CO 3 + CaCl 2  2NaCl + CaCO 3 2Na + (aq) + CO 3 -2 (aq) + Ca +2 (aq) + 2Cl - (aq)  2Na + (aq) + 2Cl - (aq)+ CaCO 3 (s) This is an Ionic Equation The solid that forms from a solution (CaCO 3 ) is a precipitate The Na + and Cl - do not participate in the reaction -- they are spectator ions The Net Ionic Equation is: Ca +2 (aq) + CO 3 -2 (aq)  CaCO 3 (s)

19. 19 Reactions between ions in solution will occur only if one of the products is insoluble (gas or solid), or water Example: NaCl + KNO 3  ? Solution: the possible products, NaNO 3 and KCl, are both soluble so, no reaction Use reference table F to predict solubilities Pb +2 + 2I - = PbI 2

20. 20 Solution Equilibrium When solute is added to solvent, it dissolves until an equilibrium is reached: solid dissolved The rate of dissolving = the rate of precipitation (the maximum amount of solute is dissolved) The solubility of a substance is its concentration at equilibrium (at saturation) Factors affecting solution equilibrium: A. solid/liquid solutions -- Temperaturealmost all solids are more soluble at high temps when the solution is saturated

21. 21 B. gas/liquid solutions -- Temperature gases are less soluble at high temperatures -- Pressure gases are more soluble at high pressures See reference table G

22. 22 Chemical Equilibrium Many chemical reactions are reversible A  B forward reaction A  B reverse reaction B  A Both reactions use the same PE diagram act. E ΔH act. E ΔH forward: ΔH +; endo reverse: ΔH  ; exo ΔH’s are equal but opposite A B

23. 23 A  B Equilibrium occurs when: 1. the rate of the forward and reverse reactions are equal 2. the concentration of the product and reactant remain constant 3. no visible changes are occurring

24. 24 Equilibrium Shifts LeChatelier’s principle: When a system at equilibrium is subject to a change, the equilibrium will shift to undo the change Three factors can affect chemical equilibrium: 1. Concentration of product or reactant Example: A + B  C + D is at eq.; then more D is added the equil. will shift to reduce the amount of D more A and B will form, and C will be used up to combine with the excess D. the equilibrium shifts left Example: AgCl(s)  Ag + (aq) + Cl - (aq) at eq.; then NaCl is added adding NaCl would add more Cl - ; therefore, the equil. shifts left, and AgCl(s) will precipitate Result:

25. 25 2. Pressure: affects equil. involving gases If pressure increases,the equil. shifts to the side with fewer moles of gas Example: N 2 (g) + 3H 2 (g)  2NH 3 (g) at eq., then pressure 1 mole + 3 moles  2 moles Result: to cause the pressure to go back down, the eq. shifts right Example: CaCO 3 (s) + HCl(aq)  CaCl 2 (aq) + H 2 O(l) + CO 2 (g) What happens to the amount of CaCl 2 if pressure decreases? Result: 0 moles  1 mole Eq. shifts right to restore pressure, producing more CaCl 2 If gases are equal on both sides, equil. does not shift

26. 26 3. Temperature: consider heat as a product or reactant Example: exothermic reaction NaOH + HCl  NaCl + H 2 O + heat Result: Eq. shifts left to get rid of excess heat Example: endothermic reaction N 2 + O 2 + heat  2NO Result: Eq. shifts right Catalysts do not shift the equilibrium -- they speed up the forward and reverse reactions equally at eq., then temp

27. 27 Types of Reactions heat of reaction(ΔH) [favorable?] entropy (ΔS) [favorable?] 1. decrease ( - ) [yes]increase ( + )[yes] YES 2. increase ( + ) [no] decrease ( - ) [no] NO 3. decrease ( - ) [yes] decrease ( - ) [no] 4. increase ( + ) [no]increase ( + )[yes] The result in cases 3 and 4 depends on temperature Example : H 2 O (solid)  H 2 O (liquid) -- solid to liquid, energy increases, energy absorbed ΔH positive, unfavorable -- solid to liquid, entropy increases ΔS positive, favorable So, ice melts spontaneously only at high temperatures Only at low temp Only at high temp