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A2 chemistry 4.3
Edexcel chemistry - rates
| Question | Answer |
|---|---|
| List the methods for determining the rate of reaction | Titration (chemical method), colorimetry, IR, polarimetry, conductivity, pH, product quantity, clock reactions |
| How can you use a titration to determine reaction rate? | Mix reactants (known conc) + start timer, at regular intervals withdraw sample + quench e.g. cold water, NaHCO3. Titrate with standard solution - acid/base, I2/Na2S2O3, halide/AgNO3 titre volume proportional to concentration of substance. Conc-time graph. |
| How can you use colorimetry to determine reaction rate? | Use if 1 reactant or product is coloured. Place reaction mixture in SPECTROPHOTOMETER / COLORIMETER with filter. At regular intervals, measure how much light of a particular frequency is absorbed by the chemical, depends on its concentration |
| What are the disadvantages of this method? | Heat from the lamp may increase the reaction rate, light may be absorbed by reaction tube |
| How can you use polarimetry to determine reaction rate? | Use if there is a change in optical activity e.g. dextro-> laevorotatory. Place reaction mixture in polarimeter. At regular intervals, measure angle + direction of rotation of plane polarised light, angle depends on concentration of chemical |
| How can you use IR to determine reaction rate? | Place reaction mixture in reaction tube. At regular intervals, measure how much IR radiation of varying frequencies is absorbed by the chemical (height of peaks) which depends on concentration |
| How can you measure product quantity to determine reaction rate? | At regular intervals: measure mass of solid, volume of gas (in a gas syringe e.g. NH3, Cl2 or in a measuring cylinder over water) |
| What are clock reactions + how can they be used to determine reaction rate? | Measure time taken for a change to occur e.g. colour change. If < 10% of reactants are used up, rate is proportional to 1/time. Repeat with different concentrations of reactants. Plot graph of rate-concentration |
| How can you use pH to determine reaction rate? | Use if 1 reactant/product is acidic/ alkaline and reaction is aqueous. pH probe continuously measures change in pH over time. Problem: logarithmic quantity, needs very accurate equipment |
| Define: rate of reaction | Change in concentration of reactant/ product over time |
| Define: rate equation | Shows how concentrations of reactants/ intermediates/ catalysts affects the rate of reaction and thus shows which species are involved in RDS |
| Define: order and partial order | Order- sum of powers to which the concentrations of SPECIES are raised in rate equation. Partial order - power to which the concentration of a SPECIES is raised in rate equation. |
| Why must orders be determined experimentally | Rate equation only includes species in RDS, in a multi-step mechanism RDS is not the same as overall equation . Therefore, can't be determined directly from reaction stoichiometry |
| Give the units for: k, [], rate | k - must be determined for each calculation. [] - mol/dm3, rate - mol/dm3/s |
| What is meant by 0 order? Draw graphs | Rate of reaction is unaffected by the concentration of reactant, not involved in RDS, excluded from rate equation |
| What is meant by 1st order? Draw graphs | Rate of reaction is directly proportional to the concentration of the reactant, 1 mol in RDS |
| What is meant by 2nd order? Draw graphs | Rate of equation is directly proportional to the square of the concentration of the reactant, 2 mol in RDS |
| How can initial rates be used to determine orders of reaction? | Work out initial rate of reaction, repeat with different concentrations of reactants to see how it affects the rate of reaction |
| How can a concentration-time graph be used to determine orders of reaction? | 1. half lives (time taken for concentration to halve): 0-decrease, 1-constant, 2-increase geometrically eg. doubling, tripling. 2. gradients of tangents at 2 points, negative gradients correspond to positive rates, relationship between change in rate/conc |
| When can a volume-time graph be used? | If overall volume is constant, volume is proportional to concentration, so it can be treated as a concentratin-time graph |
| If neither reactant is in excess, what does the order calculated from a concentration-time graph tell you? | Overall order of reaction |
| If neither reactant is in excess, how would you work out partial orders? | Increase the concentration of a reactant involved in RDS and gradient of graph will increase |
| If a reactant is in excess, what will the order calculated form a concentration-time graph tell you and why? | Reactant in excess: decrease in its concentration is negligible relative to initial concentration so concentration remains fairly constant, doesn't affect rate. Order is partial order of reactant not in excess. Does NOT tell you overall order of reaction |
| How can a rate-concentration graph be used to determine orders? | Work out rate (1/time or conc/time). Plot rate against [A]. Straight line = 1st order. Extend line to origin. Not straight line, plot rate against [A]2. Straight line = 2nd order |
| Give another method for determining orders graphically | Gradient of straight line on graph of log(rate) against log(concentration) = order |
| What must be kept constant during rate experiments + why? | Temperature, affects rate of reaction - affects average KE of molecules + proportion of molecules with KE>Ea, and frequency of collisions |
| Define: RDS, molecularity | RDS - slowest step with highest Ea. Molecularity - number of species involved |
| What do you know about a reactant if it's not involved in the rate equation? | Must enter reaction after RDS |
| What do you know about the mechanism of nucleophilic substitution for primary halogenoalkanes from kinetics? | SN2 - 2nd order - rate = k[halogenoalkane][nucleophile]. 1 step. Carbon-halogen bond breaks, carbon-nucleophile bond forms, transition state |
| What do you know about the mechanism of nucleophilic substitution for tertiary halogenoalkanes from kinetics? | SN1 - 1st order - rate = k[halogenoalkane]. 2 steps. step 1(RDS): carbon-halogen bond breaks. Step 2: carbon-nucleophile bond forms |
| What is the Arrhenius equation, written in 2 ways? | k = Ae^ -Ea/RT. Ea in J/mol. R is gas constant, 8.31J/K/mol. T in K.................. lnk = lnA -Ea/RT. equivalent to y = mx + c |
| What factors affect the value of k and why? | Orientation factor-complexity of geometry of molecules affects likelihood that they will collide with correct orientation for a successful collision (A). T - increase T makes exponent for AE less negative, increasing k.... |
| Continued | Increase T: molecules have higher avg KE, greater proportion pf collisions are successful, increased frequency of collision. Ea: higher Ea makes exponent more negative, decreases k, fewer collisions have sufficient KE so less are successful |
| How can you work out the Ea? | Table: k, lnk, T, 1/T. Plot lnk against 1/T, work out gradient (must be negative), gradient = -Ea/R. Ea = gradient x -8.31. Ea: J/mol, +ve sign |
| Define activation energy | Minimum amount of KE that colliding molecules must possess for a successful collision |
| How do catalysts speed up reactions? | Catalysts provide an alternative pathway with a lower Ea, larger proportion of molecules have KE > Ea cat than Ea uncat so a greater proportion of collisions involving the catalyst are successful. |
| Describe a heterogeneous catalyst | Different phase to reactants: gaseous reactant adsorb onto solid catalyst active sites, bonding reactant on a surface brings them into close contact, reduces Ea because it weakens intramolecular forces in reactants, products desorb |
| What are the advantages/ disadvantages of using a heterogeneous catalyst? | Adv: easy to separate from product mixture. Disadv: rate may be slowed down by availability of active sites |
| Describe a homogeneous catalyst | Same phase as reactants: catalyst reacts with a reactant to form an intermediate compound, which reacts with another reactant to reform the catalyst. |
| Give examples of homogeneous and heterogeneous catalysts | Heterogeneous: iron in haber process, catalytic converters in cars (palladium, platinum, rhodium). Homogeneous: Fe2+ in reaction between I2 and S2O82- as naturally reactants repel each other, O* radicals |
| Autocatalysis: product is a catalyst, so rate suddenly increases. Prove by adding product at the start, initial rate will be greater | |
| Write the equation for the reaction between iodide and iodate ions, and state how you would measure the rate of reaction | IO3- + 5I- + 6H+ -> 3H2O + 3I2. Add Na2S2O3 as a time delay, reduces I2 back to I-. After all Na2S2O3 has reacted, starch + I2 will turn blue-black. Clock reaction: measure time taken for colour change from colourless -> blue-black |
| Draw the mechanism for the reaction between iodine and propanone, write the chemical equation and rate equation | CH3COCH + I2 -> (H+) -> CH2ICOCH3 + HI. Rate = k [CH3COCH3] [H+] |
| How would you measure the rate of this reaction? | 1) colorimeter- measure colour change from brown -> colourless. 2) At regular intervals, withdraw sample of reaction mixture + quench with NaHCO3, titrate with Na2S2O3 until straw yellow, starch, titrate till blue-black -> colourless, conc-time graph |
| Define: collision theory | For substances to react, molecules must collide with the correct orientation and KE greater than or equal to Ea |
| How do pressure + concentration affect the rate of reaction + in what situations? | Pressure (homogeneous reaction of gases), concentration (solutions). Increase P/C = increased # reactant molecules per unit volume = increased frequency of collision = by probability, increased frequency of successful collision = faster rate |
| How does temperature affect the rate of reaction? | Increase T increase rate. 1) molecules have higher average KE, so a greater proportion of collisions have KE greater than or equal to Ea and are successful. 2) increased frequency of collision. 3) makes exponent for Arrhenius equation less negative |
| How does surface area affect the rate of reaction and in what situations? | Heterogeneous reaction involving a solid. Increased surface area of solid e.g. powder = increased number of reactant particles exposed at a time = increased frequency of collision = by probability increased frequency of successful collision = faster rate |
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