Chem 12 Unit 1 Kinematics Student 2017 notes PDF

S. 1.1 - Introduction to Reaction Kinetics. Reaction Kinetics ​ - study of rates of rx. and the factors which affect the rates. (note: “rx” = reaction(s)) ​ Expressing Rates ​ rate = Δ amount (a reactant or product) ​ ​ ​ ​ ​...

S. 1.1 - Introduction to Reaction Kinetics. Reaction Kinetics ​ - study of rates of rx. and the factors which affect the rates. (note: “rx” = reaction(s)) ​ Expressing Rates ​ rate = Δ amount (a reactant or product) ​ ​ ​ ​ ​ Δ time ​ ​ Note: A time unit is always in the denominator of a rate equation. ​ eg. Zn(s) + 2HCl(aq) H2(g) + ZnCl2(aq) ​ ​ eg. If 20 g of H2SO4 are used up after 16 min, what is the average reaction rate? Note: Some rxs, when written in ionic form show that some ions don’t change concentration. ​ eg. Mg(s) + 2HCl(aq) H2(g) + MgCl2(aq)​ ​ Note: To write an equation in IONIC FORM, dissociate all the aqueous (aq) compounds: + - 2+ - ​ ionic form :​ Mg(s) + 2H (aq) + 2Cl (aq) H2(g) + Mg (aq) + 2Cl (aq) (use ion chart)​ ​ ​ ​ ​ ​ ​ ​ ​ 1 Calculations Involving Reaction Rates ​ When doing calculations involving rate, amount (grams, moles, Litres etc.) use the general equation: When converting units, make sure units can be crossed out. You also must use molar mass to go grams moles. eg.) You would use 22.4 L for conversions moles L (STP) for gases. 1 mol eg.) 2  NOTE: This conversion is only used for gases at STP! Try this problem: The rate of a reaction is 0.034 g of Mg per second. Calculate the number of moles of Mg used up in 6.0 minutes. Comparing rates using balanced equations -​ Use coefficient ratios - only proportional to mol /s (not to g/s) ​ ​ eg. ethane ​ ​ ​ 2C2H6 + 7O2 4CO2 + 6H2O ​ ​ consumed produced ​ ​ eg. If ethane is consumed at a rate of 0.066 mol /s, calculate the rate of consumption of O2 in mol /s ​ ​ ​ If ethane is consumed at a rate of 0.066 mol /s calculate rate of production of CO2 ​ ​ ​ 3 ​ ​ ​ ​ ​ ​ ​ ​ ​ -​ When other units used – you must use moles to (go over the “mole” bridge). (you may go from L L of one gas to another at STP) ​ ​ eg. Given: 2Al + 3Br2 2AlBr3 If 67.5 g of Al are consumed per second - calculate the rate of consumption of Br2 in g/s. You may have to use a few conversions and the “rate equation” to arrive at an answer. As you did in Chem. 11, make a “plan” first and make sure your units all cancel the correct way! eg. An experiment is done to determine the rate of the following reaction: ​ 2Al(s) + 6 HCl (aq) 3 H2(g) + 2 AlCl3 (aq) It is found that the rate of production of H2(g) is 0.060 g/s. Calculate the mass of Aluminum reacted in 3.0 minutes. 4 Homework: -​ Pg. 2, Q. 1 – 5. S. 1.2 – Methods of Measuring Reaction Rates. Monitoring Reaction Rates. -​ Properties which can be monitored (measured at specific time intervals) in order to determine rx rate. ​ Note: Must consider -subscripts (s) (l) (g) (aq) ​ ​ - coefficients of gases ​ ​ - heat (endo or exo?) ​ 1)​ Colour changes -​ Only in reactions where coloured reactant is consumed or new coloured product formed. eg. Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g) + heat ​ copper clear ​ blue​ clear​ brown -​ In this case could measure - intensity of _____________________ solution ​ ​ ​ ​ - intensity of ______________________ gas ​ eg. Cu(NO3)2(aq) + Zn(s) Cu(s) + Zn(NO3)2(aq) ​ blue​ grey​ reddish colourless -​ As this reaction proceeds the blue colour fades ​ 2+ - 2+ - ​ In ionic form: Cu (aq) + 2NO3 (aq) + Zn(s) Cu(s) + Zn (aq) + 2NO3 (aq) 2+ 2+ ​ Net ionic:​ Cu (aq) + Zn(s) Cu(s) + Zn (aq) ​ ​ [________ is blue!] ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ -​ Colour intensity can be measured quantitatively using a spectrophotometer ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ Rate of rx equation for colour change. ​ ​ rate = 5  2)​ Temp changes -​ In exothermic reaction temperature of surroundings will increase. -​ In endothermic reaction temperature of surroundings will decrease. -​ Measured in insulated container (calorimeter). Rate of rx equation for temp change.​ ​ rate = 3)​ Pressure changes​ (constant volume or sealed container) ​ -​ If more moles of gas (coefficient) in products pressure will go up ​ ​ ​ Zn(s) + 2HCl(aq) H2(g) + ZnCl2(aq) ​ ​ O m.o.g.​ ​ 1 m.o.g. -​ If more MOG in reactants - pressure will decrease. Rate of rx equation for pressure change.​ ​ ​ rate = ​ (constant volume --- closed container) ​ ​ ​ -​ If equal MOG, pressure will not change: ​ ​ NO2(g) + CO(g) CO2(g) + NO(g) ​ ​ 2 m.o.g.​ ​ 2 m.o.g. ​ 4)​ Volume change​ (constant pressure eg. balloon or manometer) ​ ​ eg. If more gas is produced, volume of balloon will increase Rate of rx equation for volume change. ​ ​ rate = ​ (constant pressure) ​ ​ ​ 5)​ Mass changes -​ If only one solid is used up -​ Could remove periodically and weigh it: 6 ​ ​ ​ Mg(s) + 2HCl(aq) H2(g) + MgCl2(aq) ​ (periodically remove Mg and weigh what is left) -​ If one gas is produced and escapes, measure mass of what’s left in container (mass of container and contents) eg. ​ heat + CaCO3(s) CaO(s) + CO2(g) ​ ​ Rate of rx equation for mass change. rate = Note: it’s not practical to measure masses of (aq) substances separately since they are mostly water. ​ eg.​ Ca(s) + 2HNO3(aq) H2(g) + Ca(NO3)2(aq) ​ ​ Δ mass of HNO3(aq)​ ​ ​ Δ time 6)​ Changes in molar concentration of specific ions ​ eg.​ Mg(s) + 2HBr(aq) H2(g) + MgBr2(aq) ​ + - 2+ - ionic form: Mg(s) + 2H (aq) + 2Br (aq) H2(g) + Mg (aq) + 2Br (aq) + -​ Could monitor [ H ] - it will decrease ​ eg. rate = ​ - Note: Does the [Br ] change? No. Explain. Spectator ion. ​ Rate of rx equation for change in specific ions. rate = 7 ​ The concentration of a specific ion can be measured: -​ Using spectrophotometer. -​ Periodic samples taken and titrated to measure conc. + 7) Changes in Acidity​ ​ [H ]​ -​ Special case of #6 ​ Rate = ​ ​ pH is a measure of acidity ​ ​ pH​ 0​ ​ 7​ ​ 14 ​ ​ ​ ​ more acidic​ more basic ​ ​ (less basic)​ (less acidic) ​ + -​ If H is a reactant (or any acid HCl, HNO3 etc.) + [H ] will decrease so pH will INCREASE! ​ ​ ​ ​ ​ ​ (less acidic) Rate of rx equation for change in acidity. rate = ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ Homework: -​ Q. 7 – 9, pg. 5. 8 S. 1.3 – Factors Affecting Reaction Rates. Factors affecting reaction rates -​ 2 kinds of reactions: Homogeneous reactions -​ All reactants are in the same phase ​ ​ ​ ​ (don't consider products) ​ ​ eg.​ 3H2(g) + N2(g) 2NH3(g) ​ ​ (both gases) + - ​ ​ Ag (aq) + Cl (aq) AgCl(s) ​ ( both (aq) ) Heterogeneous Reactions -​ More than one phase in reactants. ​ eg.​Zn(s) + 2HCl(aq) H2(g) + ZnCl2(aq) ​ ​ (2 diff. phases) ​ eg.​C(s) + O2(g) CO2(g) ​ (2 diff. phases) Factors that affect both homogeneous & heterogeneous. reactions 9  1.​ Temperature -​ As temperature increases, rate increases 2.​ Concentration of reactants -​ As cons. of one or more reactants increases, rate increases. -​ Also partial pressure of a gas (partial pressure of a gas is the pressure exerted by ​ ​ that gas in a mixture of gases - it’s proportional to concentration) ​ 3.​ Pressure -​ Affects reactions with gases in reactants. ​ eg.​C(s) + O2(g) ---> CO2(g) -​ As pressure increases, rate increases. Note: a decrease in the volume of reaction container increases the pressure (therefore rate) 4.​ Nature of reactants -​ Rate depends on how strong & how many bonds in reactants need to be broken. -​ In general covalent bonds are strong and slow to break. ​ ​ ​ C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)​ (slow at room temp) ​ ​ ​ ​ ​ ​ 2- - + ​ eg. 5C2O4 + 2MnO4 + 16H 10CO2 + 8H2O ​ Many bonds have to be broken and many new bonds have to form. So this reaction is slow at room temperature. ​ eg. H2(g) + Cl2(g) 2HCl(g) ( H2 and Cl2 are diatomic) ​ ​ H - H​ +​ Cl - Cl 10 ​ ​ ​ ​ ​ ​ slow at room temp. ​ Consider Phase ​ ​ A(s) + B(s) AB ​ both solids​ ​ ​ slow at room temp. ​ Fast reactions at room temperature: -​ Simple electron transfer (no bonds broken). 2+ 4+ 4+ 2+ ​ eg. Sn + Te Sn + Te (2 electrons have been transferred from Sn to Te ) fast at room temp. -​ Precipitation reactions: 2+ 2- ​ ​ eg. Fe (aq) + S (aq) FeS(s) fast at room temp. ​ ​ ​ both reactants (aq) - no bonds to break. -​ Acid base (proton transfers). -​ Intermediate in rate. + 2- - ​ eg. NH4 + SO3 NH3 + HSO3 ​ 5. Catalysts - ​ A substance which can be added to increase the rate of a rx. without being consumed itself. (reactants are consumed) ​ Demo with H2O2 + MnO2​ ​ 2H2O2(l) 2H2O(l) + O2(g)​ ​ uncatalyzed - slow ​ 2H2O2(l) 2H2O(l) + O2(g)​ catalyzed - fast ​ Inhibitors 11  -​ A substance which can be added to reduce the rate of a reaction. -​ Can combine with a catalyst or a reactant & prevent it from reacting. ​ eg. poisons (cyanide)​ - organophosphates (diazinon) ​ ​ antibiotics ​ ​ antidepressants (serotonin uptake inhibitors) ​ ​ sunscreens Factor which affects only heterogeneous reactions (more than one phase) 6.​ Surface area -​ When 2 different phases react, reaction can only take place on surface Increase surface area by cutting solid into smaller pieces (liquids in smaller droplets) ​ ​ In general​ ​ -​ Reactants with solids are slow (except powdered). -​ Gaseous reactants are faster (but watch for diatomic bonds!). -​ Reactants in ionic solution. are fastest if no bonds to break. ​ ​ ​ ​ ​ n + - eg. ppt Ag (aq) + Cl (aq) AgCl (s) ​ Some points ​ ​ 1.) Temperature affects rate of all reactions ​ ​ 2.) Pressure (or volume) affect reactions with gaseous reactants ​ ​ 3.) Concentration only affects (aq) or (g) reactants 12 ​ ​ 4.) Surface area - affects only heterogeneous reactions. ​ ​ ​ ​ ​ ​ ​ ​ Everyday situations which require control of reaction rate ​ Body chemistry ​ ​ eg. - metabolism ​ ​ - fever can destroy bacteria ​ ​ - neurotransmitters - awareness, sleep etc. ​ ​ hormones - messengers (adrenaline, sex hormones) ​ ​ catalysts - enzymes (digestive etc) ​ ​ - aging​ Fuels ​ - concentration of O2 important ​ ​ ​ - to increase combustion rate - increase [ O2 ] ​ ​ ​ ​ ​ ​ ​ - increase surface area ​ ​ ​ ​ ​ ​ ​ - increase temperature ​ ​ ​ ​ ​ ​ ​ - catalyst (wood stoves etc) ​ ​ ​ - to decrease combustion rate ​ ​ ​ ​ ​ - water on fire -smothers it (decreases O2) ​ ​ ​ ​ ​ ​ ​ - cools it ​ ​ ​ ​ ​ - fire retardant​ - forest fires ​ ​ ​ ​ ​ ​ ​ - children's clothing ​ ​ ​ ​ ​ - airplane fuels​- when spilled Industrial Processes ​ - produce product quickly ​ ​ eg. - fiberglass ​ - uses catalyst (hardener) ​ ​ ​ ​ ​ hardens fast but not too fast ​ ​ ​ - glue ​ - epoxy uses catalyst ​ ​ ​ ​ ​ - contact cement fast ​ ​ ​ - concrete​ - ceramics – paint ​ ​ ​ - oxy- acetylene welding (must be very hot) ​ ​ ​ - oil refining ​ ​ ​ - sewage treatment​ - use microbes to speed up breakdown ​ - slow down reactions. ​ ​ eg. nitroglycerine - keep cool - if too warm explodes ​ Rusting​ -(oxidation) of cars etc. ​ ​ ​ - paint, sealers, etc. prevents O2 from contact with surface 13 ​ ​ ​ - keep cool & dry ​ Cooking​ - improves taste ​ ​ ​ - kills some bacteria ​ ​ ​ - if too hot causes burning and productions of carcinogens ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ (benzopyrenes)​ Food preservation ​ ​ ​ - lower temperature ​ ​ ​ - anti-oxidants (eg. ascorbic acid) ​ ​ ​ - keep from O2 (sealing) ​ ​ ​ - preservatives (nitrates, nitrites) Think of more! Homework: -​ Q. 10 – 17 (P. 7 – 10). 1.4 – Measuring Reaction Rates. Measuring Reaction Rates. -​ Different methods for different reactions. -​ Must look at subscripts & use common sense. eg. CaCO3(s) + 2HCl(aq) H2O(l) + CO2(g) + CaCl2 (aq) + - 2+ - ionic form: CaCO3(s) + 2H (aq) + 2Cl (aq) H2O(l) + CO2(g) + Ca (aq) + 2Cl (aq) + 2+ net ionic form: CaCO3(s) + 2H (aq) H2O(l) + CO2(g) + Ca (aq) ​ (Getting rid of like species on either side of rx)​ ​ ​ ​ 14  -​ As CO2 escapes, mass of the rest of the system will decrease. -​ So rate could be expressed as. ​ ​ r= (open system) ​ ​ ​ ​ Δ time Note ​ ​ rate = slope of amount. vs. time graph Disregard sign of slope. -​ Slope will be negative if something is being consumed and positive if something is being produced. -​ Rate is just the Δamount/Δtime ) Note For a changing rate (slope) –which is more realistic -rate could be expressed over a certain interval. 15  Or rate at a certain point in time is the slope of the tangent at that point. ​ Homework: -​ Q. 6 – 9 pg 3 & 5. -​ Q. 18 – 19 pg. 11. S. 1.5 – Reaction Rates and Collision Theory. Collision theory​ -​ Explains rates on the molecular level Basic idea (basic premise) -​ Before molecules can react, they must collide. ​ H2​ +​ I2​ ​ ​ 2HI ​ ​ first​ ​ ​ ​ later​​ later still ​ ​ ​ ​ ​ successful collision ( reaction ) 16 ​ How collision theory explains : ​ Effect of concentration ​ low conc. both​ ​ high conc. blue​ ​ high conc. both ​ ​ ​ ​ ​ low conc. red ​ low chance​ ​ ​ higher chance​ ​ very high chance ​ ​ ​ of collision​ ​ ​ of collision​ ​ of collision (slow reaction) (faster reaction) (much faster reaction) ​ ​ Effect of temperature -​ When molecules move faster more collisions per unit time faster rate. -​ Also - when they move faster they collide with more kinetic energy. (hit harder) Homework: -​ Q. 20 – 22 (P. 12). ​ S. 1.6 – Enthalpy Changes in Chemical Reactions. (Enthalpy (H) & enthalpy change (ΔΗ )) Enthalpy ​ -​ The “heat content” of a substance. or -​ The total KE & PE of a substance at const. pressure. ​ Chemists are interested in enthalpy changes (ΔΗ ) 17  Equations and heat ​ ​ H2 + S ---> H2S ΔΗ = - 20 KJ ( -ive ΔΗ means exothermic) ​ 6C + 3H2 ---> C6H6​ ΔΗ = + 83 KJ ( +ive ΔΗ means endothermic)​ Thermochemical equations: -​ “Heat Term” is right in the equation. NO “ΔΗ” shown beside the equation! -​ “Heat term” shown on left side of arrow - endothermic (“it uses up heat like a reactant”)​ ​ eg. CH3OH + 201KJ C(s) + 2H2(g) + ½ O2(g) -​ “Heat term” shown on right side of arrow -exothermic ( “it gives off heat like a product”)​ ​ ​ eg. S(g) + O2(g) ---> SO2(g) + 296 kJ ​ ​ ​ ​ ​ ​ ​ ​ Homework: -​ Q. 24 – 29, pg. 14 – 28. S. 1.7 - Kinetic energy distributions -​ Some molecules have a very low kinetic energy level and some have a high kinetic energy level. 2 -​ Remember​ KE = ½ mv or = the Ea. In this case about 1/5th to 1/6th of the molecules have sufficient KE (the shaded region is about 1/5th to 1/6th the total area under the “Temperature T2 curve) ​ ​ Rule of thumb​ ​ ​ ​ ​ ​ ​ If the activation energy (threshold) is near the tail of the curve: -​ If the temperature is increased by 10oC reaction rate will about double. ​ (ie. about twice the number of molecules have sufficient KE for a successful collision.) 21 On the graph above, temperature T2 is 10°C higher that T1. The area under the T2 curve to the right of the Activation Energy is about twice the area under the T1 curve. This means that the number of molecules with sufficient KE at T2 is about double the number of molecules with sufficient KE at T1.​ ​ ​ ​ Note​ -​ If Activation Energy or ME is near the middle of the curve (or left side) -​ Reaction is already fast, so an increase in temperature has a less drastic effect on the reaction rate. ​ 22 S. 1.8 – Activation Energies – part 2: ​ ​ (back to collision theory.....) ​ Potential and Kinetic energy during a collision -​ As colliding molecules approach the repulsion slows them down so kinetic energy decreases. -​ As they push against the repulsive force potential energy increases (like compressing a spring). -​ so: Kinetic Energy ​ ​ Potential Energy ​ KE + PE = Total E (stays constant) 23 ​ ​ From Kinetic Energy diagrams to Potential Energy Diagrams: Three possible situations between molecules: 1.​ KE < Ea PE ∴ Reactants do not have enough energy to react --- Ineffective Collision. 2.​ KE = Ea PE ∴ Reactants have just enough energy to reach the top of the hill, come to a standstill and react. 3.​ KE > Ea PE ∴ Reactants will lock together in space and will change into products. Potential energy diagrams 24 25 If colliding molecules don’t have enough KE to convert to PE to make it “over the Activation Energy Barrier”, it is an UNSUCCESSFUL collision and there is NO reaction. The molecules will just bounce off of each other unchanged. ACTIVATION ENERGY (Ea) -​ The minimum energy required for a successfull collision. (or) The minimum energy reacting molecules must have in order to form the Activated Complex. Note: The Activation Energy (Ea) is fixed by the nature of the reactants (#’s and strengths of bonds in reactants.) Ea is NOT affected by Δtemperature or Δ concentration. Activated Complex: A very short-lived, unstable combination of reactant atoms that exists before products are formed. -​ Favourable alignment between reacting molecules. -​ Bonds are being broken and formed. eg. for the rx.​ A2 + B2 2AB: ​ ​ the above collision has unfavourable alignment ​ (need higher energy for collision to be effective) 26 ​ In the above collision, the reactants have favourable alignment ​ ​ (less energy needed for an effective collision) ​ Potential energy diagram To Summarize Collision Theory so far: ​ For any successful collision (one resulting in a reaction): ​ ​ ​ ​ 3 Requirements:​ 1. Particles must collide ​ ​ ​ ​ 2. They must collide with sufficient energy > Ea ​ ​ ​ ​ 3. They need to have correct alignment (collision geometry)​ ​ ​ ​ ​ (to keep Ea as low as possible) Temperature’s role (Sufficient KE) -​ The temperature determines how many (or what fraction of the) molecules will have energy > Ea (to make it over the barrier & have a successful collision) ​ 27 Recall KE distributions: At a LOW temperature. 28  Note: In the diagrams on the previous page and above, that only a small fraction of the molecules had enough energy to overcome the Activation Energy barrier. Now, at a Higher Temperature: At the higher temperature, a greater fraction of the molecules have sufficient energy to “make it over” the Activation Energy barrier. (ie. a greater fraction of the molecules posses enough energy to form the Activated Complex): Looking at the diagram above, you can see that at a higher temperature, a greater fraction of the molecules have sufficient energy to make it over the barrier. Therefore the reaction is faster. 29  Increasing the temperature increases the fraction of molecules which have sufficient ​ energy to form the Activated Complex (ie. sufficient energy to “make it over” the ​ activation energy barrier.) NOTICE: Change in temperature does NOT change the Potential Energy diagram at all. Temperature does NOT affect the Activation energy or the ΔΗ ! Question: Consider two reactions at the same temperature: Which reaction is faster? ________________ Explain why. Ea, ΔΗ and bond strengths for forward and reverse reactions 30 Reactants Products -​ Ea(f) = forward Ea. -​ Ea(r) = reverse Ea. Note: Ea is endothermic ∴ Energy must always be added to the system to get the reacting molecules to the top of the hill. Endothermic reactions: -​ Ea(f) = Ea(r) + Δ H (Note Δ H > 0) -​ Ea(f) > Ea(r) Exothermic reactions: 31  -​ Ea(f) = Ea(r) + Δ H (Note Δ H < 0) -​ Ea(f) < Ea(r) Homework: -​ Pg. 23-24, # 33 – 40. -​ Pg. 25, # 41 – 45. S. 1.9 – Reaction Mechanisms Looking at reactions ​ Multiple reactants ​ - + eg. 5C2O42- + 2MnO4 + 16H -​ Involves 23 reacting particles. -​ Chances of this taking place in one step are almost “0”. ​ Even a 3 particle collision ​ ​ eg. 2H2(g) + O2(g) 32 ​ -​ 1,000 times less probable to occur in single reaction step than a 2 particle collision. ​ Note.​ ​ ​ ​ ​ ​ ​ ​ ​ + - -​ Most reactions (other than simple 2 particle collisions eg. Ag + Cl AgCl(s) ) take place in a series of simple steps. -​ Each step depends on the others before it. Reaction Mechanisms. -​ The series (sequence) of steps by which a reaction takes place. -​ Each step is called an Elementary process. -​ Cannot be determined by just looking at the overall reaction. Example ​ ​ For the overall reaction:​ 4HBr + O2 🡪 2H2O + 2Br2 ​ ​ ​ ​ ​ ​ ​ ​ ​ 5 reactant particles. Doesn’t take place in a single step! ​ Mechanism ​ step 1:​​ HBr + O2 HOOBr (found to be slow) step 2:​​ HBr + HOOBr 🡪 2HOBr (fast) step 3:​​ HOBr + HBr 🡪 H2O + Br2 (very fast) Rate determining step (RDS). -​ The slowest step in the mechanism. -​ The overall reaction can never be faster than the RDS -​ The only way to speed up an overall reaction is to speed up the RDS (eg. by increasing. The concentration of a reactant in the RDS) eg. In the case of 4HBr + O2 🡪 2H2O + 2Br2. RDS -​ Step 1, since it is the slowest step. Speed up rx. -​ Increase [reactants] ([HBr] or [O2]) in step 1 will increase the rx. Note. -​ Speeding up a fast step (not RDS) will have no effect on the overall rate. (eg. Adding HOOBr or HOBr has no effect). ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ Determining overall reaction given steps (mechanism).​​ -​ Cancel stuff which is identical on both sides - add up what’s left. 33 ​ eg.​ 1.)​ A + 2X AX2 ​ ​ 2.)​ AX2 + X AX + X2 ​ ​ 3.)​ AX + A A2 + X ​ _____________________________________________________________________ overall rx: Determining overall rx step given overall rx and all but one rx step. -​ Place overall rx at bottom of rx mechanism. -​ Cross out like species and place missing species into missing step. ​ eg. The following reaction occurs in a 3 step mechanism: ​ Overall rx: 2A4+ + B+ 2A3+ + B3+ Based on the overall rx and the following information, find step 3. 4+ 2+ 3+ 3+ ​ step 1:​ A +C C +A 4+ 3+ 4+ 3+ ​ step 2:​ A +C C +A ​ ​ step 3:​ ​ ​ ​ ​ ​ Question: Consider the following reaction for the formation of HCl in the presence of light. ​ ​ Overall rx: Cl2 + CHCl3 HCl + CCl4 ​ Based on the overall rx and the following information, find step 2. ​ Step 1:​Cl2 Cl + Cl ​ Step 2:​ ​ Step 3:​Cl + CCl3 CCl4 34 ​ Reaction intermediate -​ A species (atom, molecule or ion) which is produced in one step and used up in a later step. (appears on right & also lower on left) ​ eg. For the mechanism: ​ 1)​ HBr + O2 HOOBr ​ 2)​ HBr + HOOBr 2HOBr ​ 3)​ 2HBr + 2HOBr 2H2O + 2Br2 ​ ​ intermediates are ___ & ______________ ​ ​ ​ ​ ​ ​ ​ ​ ​ ​ Intermediates -​ Doesn’t accumulate (like a product) because as soon as it is formed, it gets used up again (like money). -​ Not necessarily unstable. (in other circumstances, they may last a while). Activated complex. -​ Very unstable and short-lived. It doesn’t usually obey bonding “rules”.​ (see diagrams p. 26 & 27) (very high PE, temporary arrangement) Homework: -​ Pg. 28 # 46 – 53. S. 1.10 – Energy Diagram of a Reaction Mechanism. From KE diagram to PE diagram: PE diagram for a reaction mechanism 35 Notes:​ -​ Each “bump” is a step. -​ The higher the bump, (greater Ea) the slower the step. -​ The highest bump (from the reactants level) is for the RDS. -​ AC’s at top of bumps, intermediates in middle “valleys”, products in the final “valley”. -​ The Ea for the forward overall rx. is vertical distance from reactants to top of highest bump. -​ Ea for the overall forward reaction is the difference in energy between the reactants and the top of the highest peak. Question: Given the following Potential Energy Diagram for a reaction mechanism: ​ 1. This mechanism has steps 2. Ea for overall rx = kJ ​ 3. Step is the RDS 4. Step is the fastest step. ​ 5. The overall rx. is thermic 6. ΔΗ = kJ ​ 7. ΔΗ for reverse rx. = kJ 8. Ea (reverse rx.) = kJ ​ 9. RDS for reverse rx. is step ​ ​ ​ Draw a Potential Energy Diagram for a reaction mechanism with 2 steps. The first step is fast and the second step is slow. The overall reaction is exothermic. With labeled arrows show the overall Activation Energy (Ea) and the ΔΗ for the forward reaction. 36 Homework: -​ Pg. 30, # 54 and 55. -​ Worksheets. Section 1.11 – The Effect of a Catalyst on the Activation Energy. Catalyst: An introduced substance which provides an overall reaction with an alternative mechanism having a lower activation energy​ 37 ​ Catalyized Reaction System: -​ Energy required (Ea) is less with the catalyst, so at the same temperature, more molecules can make it over the “barrier” so reaction rate speeds up (eg. lower standards for a pass, eg. 30% will let more students pass!) -​ Catalyzed reactions usually involve more steps but it’s highest Ea (highest bump) is never as high as the uncatalyzed reaction -​ A catalyst NEVER changes the PE of reactants of products - only the route between them. (no change in ΔΗ! ). -​ Uncatalyzed reaction still continues at its own slow rate when a catalyst is added. (usually insignificant compared to catalyzed rate). -​ If catalyst speeds up forward reaction, it also speeds up (reduces Ea for) the reverse reaction. 38 Catalysts sometimes work by: ​ eg.​ 2H2O2(l) 2H2O(l) + O2(g)​ (very slow uncatalyzed) - -​ Add some KI (I ) S. 1.12 – The Effect of a Catalyst on the Reaction Mechanism. ​ Catalyzed Mechanism: ​ ​ Reading: -​ See examples of catalyzed Mechanisms in textbook on pg.’s 32 – 33. -​ Read pg.’s 34 – 35 (S. 1.13 – Some uses of Catalysts.) Homework: -​ Pg. 34, # 56 – 61. -​ Pg. 36, # 62 – 63. 39