1. Rates of Reaction Section 6.1

2. Rate of Reaction The rate of reaction indicates how fast reactants are being converted to products during a chemical reaction It is the rate of formation of a product, or the rate of consumption of a reactant, divided by the corresponding coefficient in the stoichiometric equation Rate has units of mol dm -3 s -1

3. Rate Rate will equal the change in concentration divided by the change in time R → P, rate = Δ[P] = – Δ[R]_ Δt Δt Notice the negative sign for the reactants because the concentration decrease with time whereas the concentration of the products increases with time (by convention, rate is a + value)

4. Stoichiometric Consideration You have to consider the amount of each substance (number of moles) MNO 4 - (aq) + 8H + (aq) + 5Fe 2+ (aq) → Mn 2+ (aq) + 4H 2 O (l) + 5Fe 3+ (aq) The rate of appearance of Fe 3+ is five times as great as the the rate at which MNO 4 - is consumed The rate is usually considered to apply to a product that has a coefficient of 1

5. Continued Rate = - Δ[MNO 4 - ] = 1 Δ[Fe 3+ ] Δt 5 Δt A general way to write this for reaction A→B Rate = 1 Δ[B] = - 1 Δ[A] b Δt a Δt

7. Graphs Any property that differs between the reactants and the products can be used to measure the rate of the reaction The graph is drawn of that property against time The rate of reaction is proportional to the slope (gradient) of the curve or line (ignoring the sign)

8. Graphs Changes in the gradient illustrate the effect of changing conditions on the rate of reaction Usually the rate of reaction decreases with time because the concentration of the reactants decreases with time (the reaction rate usually depends on the reactant concentration)

9. Graphs It is common to compare initial rates (gradient of the tangent to the curve at t = 0)

10. Measuring Rates Basically any property that changes between the start and end of the reaction can be used Best if the property changes by a large amount Easier to use a characteristic that is directly proportional to the concentration of one or more of the components

11. Measuring Rates For instance, it is not recommended to use pH because it is a logarithmic scale Keep in mind that the units for a reaction rate are mol dm -3 s -1 It is important to keep the reaction mixture at a constant temperature because temperature affects the reaction rate Usually a water bath is used

12. Techniques for Measuring Rates Titration Collection of an evolved gas or increase in gas pressure Measurement of the mass of the reaction mixture Light absorption Electrical conductivity Clock techniques

13. Titration Remove small samples from the reaction mixture at different times Titrate the sample to determine the concentration of either a reactant or a product Only good for very slow reactions because the titration takes so much time A graph is made of concentration against time

14. Example of Titration Reaction H 2 O 2(aq) + 2H + (aq) + 2I - (aq) → 2H 2 O (l) + I 2(aq) The amount of iodine produced can be measured by titrating the mixture with aqueous sodium thiosulfate

15. Collection of an Evolved Gas/Increase in Gas Pressure The gas is collected in a gas syringe or in a graduated cylinder over water The volume collected at different times is recorded The gas must not be water soluble if it is collected over water Could monitor the gas pressure in a container of fixed volume

16. Example Reactions Zn (s) + 2H + (aq) → Zn 2+ (aq) + H 2(g) Na 2 CO 3(s) + 2HCl (aq) → 2NaCl (aq) + CO 2(g) + H 2 O (l)

17. Measurement of the Mass of the Reaction Mixture The total mass of the reaction mixture will only vary if a gas is evolved The gas should have a high molar mass for this technique to be effective (not H 2 ) CaCO 3(s) + 2H + (aq) → Ca 2+ (aq) + H 2 O (l) + CO 2(g)

18. Light Absorption Sometimes a reaction produces a precipitate which “clouds” the reaction mixture A mark can be made on a piece of paper that is then viewed through the reaction mixture When the mark is obscured, the reaction is complete Keep the depth of the liquid constant

19. Example S 2 O 3 2- (aq) + 2H + (aq) → H 2 O (l) + SO 2(g) + S (s) The yellow suspended sulfur will obscure the mark

20. Light Absorption If a colored reactant or product is involved, the intensity of the color can be used to monitor the concentration of the species You could compare the color with your eyes against a know set of standard solutions Could use a colorimeter or a spectrophotometer

21. Example Reaction between propanone and iodine to form iodopropanone The yellow-brown iodine is the only colored species involved CH 3 COCH 3(aq) + I 2(aq) → CH 3 COCH 2 I (aq) + H + (aq) + I - (aq) Use the complementary color of blue The intensity of blue light passing through the solution will increase with time as the iodine concentration falls

22. Electrical Conductivity The presence of ions allows a solution to conduct electricity If there is a large change in the concentration of ions during a reaction, the reaction rate can be found from the change in conductivity Use an instrument that measures A.C. resistance by placing two electrodes in the solution

23. Example for Electrical Conductivity PCl 3(aq) + 3H 2 O (l) → H 2 PO 3 - (aq) + 4H + (aq) + 3Cl - (aq) Electrical conductivity will increase as the number of ions increases (products)

24. Clock Techniques Some reactions occur in which the product produced can be further reacted with another substance to form another product The formation of the 2 nd product can usually be measured through a color change The time taken for the 2 nd product to appear is inversely proportional to the rate of the original reaction

25. Example H 2 O 2(l) + 2H + (aq) + 2I - (aq) → 2H 2 O (l) + I 2(aq) 2S 2 O 3 2- (aq) + I 2(aq) → S 4 O 6 2- (aq) + 2I - (aq) Blue color of the iodine-starch complex appears when all of the thiosulfate has been consumed This is inversely proportional to the rate