1. TITRATION This involves removing small samples from the reaction mixture at different times and then titrating the sample to determine the concentration of either one of the reactants or one of the products at this time. Then we can draw concentration vs time graph.

2. This method is suitable for quite slow reactions so that the time taken to titrate the mixture is insignificant compared to the total time taken for the reaction. We can overcome this difficulty by quenching (stopping) the reaction before carrying out the titration. This can be done: 1.Cooling the reaction mixture rapidly to a very low temperature 2.Adding an excess amount of a compound that reacts rapidly with one of the reactants.

3. This method can be suitable for the 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. The reaction can be quenched by adding excess insoluble solid base to neutralize the acid that is required for the reaction. In analytical chemistry, the most important use comes from the fact that the thiosulfate anion reacts stoichiometrically with iodine, reducing it to iodide as it is oxidized to tetrathionate:stoichiometricallyiodineiodide tetrathionate 2 S 2 O 3 2− (aq) + I 2(aq) → S 4 O 6 2− (aq) + 2 I − (aq)I

4. TITRATION CALCULATIONS Titration is a technique which involves measuring the volume of one solution which just reacts completely with another solution. One of the solutions have an accurately known concentration (standard solution) and this is used to find the concentration of the other solution.

5. The volume of one of the solutions will be measured with a pipette and put into a conical flask and a few drops of an indicator is added. Conical flask pipette

6. Second solution is run in from a burette until the indicator just changes color. Burette Burette has a tap and is calibrated so as to measure a variable volume of a liquid. Indicator: a chemical substance that changes colour at the equivalence point. Indicator The equivalence point is the point where the moles of titrant added equals the moles of substance initially in the solution being titrated.

7. Using a Burette When reading a buret it is important that your line of sight be in a direction perpendicular to the buret column. You will read the volume from the bottom of the meniscus. The use of a buret card prevents the lightenning of the bottom of the meniscus by random reflections in the laboratory. A 50 mL buret can be read to ±0.01 mL

8. Doing a Titration Begin by preparing your buret. Your buret should be filled with titrant solution. You should check for air bubbles and leaks, before proceding with the titration. Take an initial volume reading and record it in your notebook.

9. Prepare the solution to be analyzed by placing it in a clean erlenmeyer flask and add the indicator. Use the buret to deliver titrant into erlenmayer flask. You will see the indicator change color when the titrant hits the solution in the flask, but the color change disappears upon stirring.

10. Approach the endpoint more slowly and watch the color of your flask carefully and be sure all titrant is mixed in the flask. Make sure you know what the endpoint should look like. For phenolphthalein, the endpoint is the first permanent pale pink. If you think you might have reached the endpoint, you can record the volume reading and add another partial drop.

11. When you have reached the endpoint, read the final volume in the buret and record it in your notebook. Subtract the initial volume to determine the amount of titrant delivered. Use this, the concentration of the titrant, and the stoichiometry of the titration reaction to calculate the number of moles of reactant in your analyte solution.

12. Example: It is found that 10.00 cm 3 of 0.200 mol/dm 3 aqueous sodium carbonate requires 25.00 cm 3 of hydrochloric acid to just neutralize it. What is the concentration of the hydrochloric acid? 1. Calculate the amount in the solution of known concentration Amount of sodium carbonate: 0.200 mol/dm 3 x 0.0100 dm 3 = 0.00200 mol 2. Write the balanced equation Na 2 CO 3 + 2HCl 2NaCl + CO 2 + H 2 O 1 mol 0.00200 mol 2 mol 2x0.00200 mol= 0.00400 mol [HCl] = 0.00400 mol/0.02500 dm 3 = 0.160 mol/dm 3 3. Calculate the concentration of the unknown solution

13. Titration with a pH meter Titration with a pH meter follows the same procedure as a titration with an indicator, except that the endpoint is detected by a rapid change in pH, rather than the color change of an indicator. Arrange the sample, stirrer, buret, and pH meter electrode so that you can read the pH and operate the buret with ease.

14. To detect the endpoint accurately, record pH vs. volume of titrant added and plot the titration curve(graph of pH vs volume of titrant added.

15. Titration curves for strong acid – strong base titrations

16. LIGHT ABSORPTION If a reaction produces a precipitate, then the time taken for the precipitate to obscure a mark on a pieces of paper under the recation vessel can be used as a measure of reaction rate. This is suitable for the reaction: S 2 O 3 2- (aq) + 2H + (aq) H 2 O (l) + S (s) which produces finely divided precipitate of sulfur.

17. If the reaction involves a colored reactant or product, then the intensity of the color can be used to monitor the concentration of that species. This can be done by comparing the color by eye againist a set of standard solutions. It is better to use an instrument that measures the absorbance.

18. Absorbance and Beer Lambert’s Law Because light absorption is a function of the concentration of the absorbing molecules, a more precise way of reporting intensity of absorption is by use of the Beer-Lambert Law: Absorbance = -log(I 1 / I 0 ) = acl where: a = molar absorptivity coefficient(how strongly a compound absorbs light) c = molar concentration of solute l = length of sample cell (cm) I 0 : initial intensity I 1 : Final intensity Absorbance is found directly by UV/visible spectrophotometers.

19. CLOCK TECHNIQUES In some reactions the product can be consumed with addition of another substance. When all of this substance is consumed an observable change will occur. The time taken for this corresponds to the time for a certain amount of product to have been formed. So rate α 1/time.

20. The classical reaction for this technique is: 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) The blue color of iodine-starch complex suddenly appears when all of the thiosulfate has been consumed.