1. Chemical Kinetics Rates of chemical reactions Mechanisms of chemical reactions

2. Chemical reaction Particles involved can be ions, atoms and/or molecules

3. Reaction Rate Measure of the # of moles of reactant, used up in a volume of reaction mixture, per unit of time How much use in seconds Example Fast –NaCl + AgNO3  white ppt. Almost instantly Slow –Digestion of school cafeteria lunch

4. Avg. rate = change in # moles of product  time =  (moles Product)  t

5. Rates in terms of concentration A + B  C + D Avg. rate =  [A]  t use [ ] to express conc. in M molarity = [A] final time – [A] initial time = M/s units t f - t i

7. Plot data on graph Graph curve used to determine instantaneous rate at any point Instantaneous rate is the slope of the straight line tangent that touches the curve at the point of interest

8. Rates and Stoichiometry AgNO 3 + NaCl  AgCl + NaNO 3 Each reactant and product is 1:1 ratio Therefore rate =  [AgNO3] =  [AgCl]  t  t Not all balanced equations are 1:1 2HI  H 2 + I 2 In this case the disappearance of HI is twice the rate of the appearance of both H 2 and I 2

9. Rate = -1  [H 2 ] =  [I 2 ] 2  t  t  t General equation is aA + bB  cC + dD Rate= 1  [A] = 1  [B] = 1  [C] = 1  [D] a  t b  t c  t d  t

10. Reaction Mechanism Series of steps by which reacting particles rearrange themselves to form the products of a chemical reaction Hard to do Intermediate products often short lived

11. Effective Collision Based on collision theory Particles must collide for chem. reaction to occur Ex. –A + B  C As A approaches B they begin to interact e- shift positions Old bonds are broken New bonds are formed Collision must be EFFECTIVE

12. The reactants must approach each other at Proper angle ( CORRECT GEOMETRY) and with Sufficient energy Activation energy Minimum amount of energy required for a reaction to occur Electron rearrangement needs energy Laser disc #9 Cpt.20

13. Reaction rates Why necessary to understand? Body uses chem. Reactions to stay healthy Medicines Industry Curtail rust Speed up production of products

14. 5 Main Factors affecting rate 1. Nature of reactants In reactions bonds are rearranged Bond types affect rate due to electronegativities Slight e- arrangement, fast reaction at room temp. Ex. Ionic double displacement Covalent bonds Large # of rearrangements Slow reactions at room temp. Ex. CH 4 +2O 2  2H 2 O + CO 2

15. 2. Concentration An increase conc. of one or more reactants will usually but not always increase the reaction rate if the mixture is homogenous

16. 3. Temperature Increase temp. increase rate of most reactions An increase in temp. of 10 0 C will approximately double reaction rate Due to collision theory –More KE to change to PE –More and faster collisions –More chance for an Effective collision –Laser disc #9 Cpt. 30

17. 4. Catalysts Substance which speeds up a reaction without being permanently altered –Recoverable –Ex. Enzymes, catalytic converter Inhibitors –Chemicals that slow down reactions

18. 5. Phase Reactants in same or more than one phase Homogenous –Same phase –Fast reaction rate Heterogeneous –Reactants in different phases Slower reaction rate –Due to reactants only being able to react at the interface site (where meet) Pressure in gases –Higher pressure forces closer-faster –Low pressure further apart-slower

19. Reaction mechanisms Series of steps Reactants rearrange to form products Most reactions follow a simple series of steps Ex. Chem.A + Chem.B  Chem.C – Step 1 Reactants  Intermediate Prod – Step 2 Intermediate  Product Net equation A + B  C A + B  I 1 I 1  C A + B  C

20. Step 1 A + B  I 1 Step 2I 1 + A  I 2 Step 3 I 2  C Net Eq 2A + B  C

21. A collides with B to form intermediate product I 1 I 1 collides with another A particle to form I 2 I 2 rearranges into the final product C When final equation is written it is in the net form Done this way due to difficulty in actually knowing what the intermediate is Intermediate products are highly unstable and short lived

22. Increase conc. usually increases rate Not always true Remember there are steps Steps reactions occur at different rates Slowest step is known as RATE DETERMINING STEP To speed up reaction must increase conc. Of reactants in rate determining step Demo

23. Example 2A + B + C  D STEP 1 A + B  I 1 STEP 2 I 1 + A  I 2 rate determining step STEP 3 I 2 + C  I 3 (slowest) STEP 4 I 3  D Net Eq. 2A + B + C  D

24. A + B have collided with an Effective collision The compound initially formed that has e- undergoing rearrangement is called the Activated Complex or Transition State Complex –It is an intermediate –Unstable and short lived –2 processes it can undergo Break down back into A + B Form product(s) Shown using a Potential Energy Diagram using Activation energy

28. Thermodynamics Study of changes in energy in chemical reactions and the influence of temperature on those changes

29. Every system or form of matter has stored energy Energy in bonds P.E. due to phase, pressure and volume K.E. from random motion of particles Total of all these forms of energy referred to as heat content Internal energy is symbolized as  U

30. Heat of formation (  H, heat of reaction) is known as ENTHALPY Change in energy when reactants form products Enthalpy changes by amount of energy gained or lost Energy given off:  H is neg.value –Exothermic Energy absorbed:  H is pos. –Endothermic Show in equation Expressed in kJ/mol

31. Heat of formation depends on temp. and press. at which reaction occurs and phase Standard Heat of Formation  H f o is at 25 0 C and 1 Atm.(101.325kPa)pressure Expressed as 298K and 101.3kPa plus phase Ex. Water liquid  H f 0 = 286kJ/mol H 2(g) + 1/2O 2(g)  H 2 O (l) + 286kJ Since energy is released(part of product)  H is neg. value ΔH= -286kJ

32. Hess’s Law of Constant Heat Summation –When a reaction can be expressed as the algebraic sum of two or more other reactions, then the heat of the reaction is the algebraic sum of the heats of these other reactions –Ex. CuO (s) + H 2(g)  Cu (s) + H 2 O (g)  H=? CuO (s)  1/2O 2 + Cu (s)  H = 155kJ H 2(g) + 1/2O 2(g)  H 2 O (g)   H = -242kJ

33. Entropy Measure of disorder, randomness or lack of organization of a system Large entropy- great state of disorder –Solid is organized pattern with low entropy and as melts the particles become random and have higher disorder of larger entropy; increases as change to gas Represented by  S  S = S f – S i  S value pos. spontaneous change can occur

34. Gibbs Free Energy Equation Combines enthalpy and entropy changes to determine spontaneity Neg.value spontaneous Pos. value non spontaneous  G =  H - T  S