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THEORY The pH of the exact equivalence point in acid-base titrations depends on the relative strengths of the acid and base. For a strong acid-strong base titration the pH at the equivalence point is 7.0. If either the acid or the base is weak, the pH at equivalence will not be 7.0 due to hydrolysis of the salt. If a weak acid is titrated with a strong base, the equivalence point pH will occur on the alkaline side of neutrality (pH > 7), whereas a strong acid titrated with a weak base will have an acidic (pH <7) equivalence point. An indicator is used in titrations to approximate the equivalence point. Note that the equivalence point and endpoint are different. The equivalence point occurs when an equal number of moles of each species are present, whereas the endpoint is the point at which the observed indicator colour change takes place. A well chosen indicator will change colour very close to the pH of the equivalence point of the titration - the closer the endpoint is to the equivalence point, the better the indicator is for that specific titration. Indicators are substances (generally weak organic acids) that change colour according to the pH of the solution. Thus, if we know the approximate pH at the equivalence point of a titration we can select a suitable indicator which will change colour at that pH. An indicator may be described by the equilibrium reaction, Hin H* + In where the dissociation constant, Kain), is given by This can be rearranged to give so that K ab [H] [In] [Hin] -log[H]--log. [K]-log[Hin]+log [In] pH=pK+ logo [In] [Hin] (1) G (2) (3) (4) At the endpoint, the concentrations of the two different coloured forms of the indicator (Hin and In') will be approximately equal, so that from equation (4), pH pKa(in). Thus, for a given titration, the best indicator will be one with a pKa(in) close to the equivalence point pH. S

- 67% + The properties of some useful indicators are given in the following table: Indicator Methyl Violet Thymol Blue Methyl Orange Bromophenol Blue Methyl Red Bromocresol Purple Bromothymol Blue Phenol Red Thymol Blue Phenolphthalein Acid (Hin) (1) (2) yellow red red yellow red yellow yellow yellow Colour yellow colourless Alkaline (in) blue yellow yellow blue yellow purple blue red blue red Approximate pH range 0.0-1.6 1.2-2.8 3.1-4.4 2.9-4.6 4.2-6.3 5.2-6.8 6.0-7.6 6.8-8.4 8.0-9.6 8.3-10.0 pka(in) 1.7 3.7 770 4.0 5.1 6.3 10 N A more sensitive method of determining the pH of a solution, and hence of detecting the equivalence point in an acid-base titration, is based on the fact that the electric potentials of certain electrodes depend upon the hydrogen ion concentration of the solution in which they are immersed. The most convenient of these electrodes is the glass electrode. This consists of a glass tube terminating in a thin-walled glass bulb enclosing a dilute solution of potassium chloride (KCI) and acetic acid (CH₂COOH) in which is immersed a platinum wire coated with Ag/AgCl. The potential of this electrode, when referred to a standard electrode (usually saturated calomel or silver/silver chloride) can be directly related to pH. Thus it is possible to monitor the pH throughout a given acid-base titration. PROCEDURE JA 7.0 7.9 8.9 9.4 Solutions of 0.1 M acetic acid (CH₂COOH), 0.1 M hydrochloric acid (HCI) and approximately 0.1 M standardised sodium hydroxide (NaOH) are provided. When reading the burette, use the correct number of significant figures-a burette can only be read to an accuracy of 10.05 ml. A. Indicators and Endpoints. Dispense a 20 ml aliquot of the CH₂COOH solution into a clean 250 ml conical flask Titrate with NaOH, using phenolphthalein as the indicator. No more than 4 drops of indicator should be added. (3) In order to observe the range of pH over which the indicator changes colour, carefully add NaOH until localized colour changes indicate the approaching endpoint Add further NaOH drop-wise until a just-detectable permanent colour change is observed. >

(4) Record the volume of the titre. Keep this first titration solution as a reference for subsequent titrations. (5) Repeat steps (1-4) until you obtain three concordant titres (+ 0.1 ml). (6) Repeat steps 1-5 using methyl orange as the indicator. It may be necessary to use 5 to 6 drops of indicator to give a convincing colour change. B. The pH Meter and Equivalence Points. Note: Great care must be exercised when handling the glass-calomel electrode as it is extremely fragile. Ensure that the electrode is always sitting in water - at no time should it be allowed to dry out. The pH electrode will need to be calibrated at the start of the experiment. Calibration of pH Electrode using Standard Buffers. Remove the pH electrode from distilled water beaker and thoroughly rinse several times with distilled water. Switch on pH meter. Carefully dry the electrode with a Kimwipe. Place the electrode in the flask labelled Buffer Solution Nº 1 pH = 4.0 and gently stirred solution. Wait a few minutes until pH has become constant and press calibration button and hold it down until the meter said OK (a few seconds). 1. 2. 67% +)| 3. 4 5. Repeat rinsing and drying of the electrode and then place pH electrode in flask labelled Buffer Solution N° 2 pH = 9.2 Wait a few minutes until reading has become constant and then hold calibration button until OK appears. Repeat rinsing and drying of pH electrode and placed it back in the calibration solution pH=4.0. The pH reading should eventually 4.00 1.02 and after rinsing and drying again the electrode is readily to use. NOTE: If the electrode does not read in the range 3.98-4.02 the entire procedure above must be repeated until it does. The apparatus consists of a pH meter with a glass-calomel electrode. The electrode may be left in the titration vessel during the experiment. However, it must be removed immediately the titration is complete, dosed.thoroughly under the tap and stored in distilled water. If the electrode is still reads above 7.0, place the electrode in a solution containing a small amount of Buffer No 1 (pH 4) has been added. Once the pH reading is below 5 remove electrode. rinsed again with distilled water. If pH is still not below 7 in the distilled water, cepeat procedure (see demonstrator). The electrode must not be allowed to dry out. The demonstrator will explain the operation of the pH meter.

31 Determine equivalence point (1) The pH meter is used to follow the pH throughout the course of the titration of 0.1 M acetic acid with 0.1 M NaOH. (2) (3) 67% + (5) Dispense a 20 ml aliquot of the acetic acid solution into a clean 100 ml beaker. Carefully rinse the glass electrode with distilled water and then immerse it in the titration vessel. Add sufficient distilled water to cover the electrode. Why doesn't this addition of water affect the titration results? (4) Measure the pH of the solution for a series of volumes of added NaOH, carefully stirring the solution after each addition. When approaching the equivalence point it is essential that the solution is thoroughly mixed and that you allow 10 to 15 seconds for the pH reading to stabilize due to the fluctuating needle. As a guide, NaOH may be added in the following way: 2 ml additions until 16 ml has been added. (1). (ii). Two or three additions of 0.5 ml. (ii) 0.2 ml additions until the pH is approximately 11. (iv). 2 ml additions until approximately 40 ml of NaOH has been added. (6) A computer will be used to prepare a graph of pH against volume of NaOH added. Your demonstrator will explain the operation of the computer program. (7) Using the pH meter as described above follow the course of the titration of 0.1 M HCI with 0.1 M NaOH, again plotting the titration curve usin the computer. Clean-up: a. Burettes need to be rinsed with distilled water, inverted and left with taps open b. Make sure that the glass electrode is rinsed and left immersed in distilled water c. Glassware must be returned to the correct drawers GRAPHICAL ANALYSIS (G1) Mark the equivalence point on each plotted graph. Record the pH and the volume of NaOH added at the equivalence point. (G2) On the CH.COOH v NaOH plot, mark the volume of NaOH added at the endpoint of your titrations with (1) phenolphthalein and (ii) methyl orange in Part A Record the corresponding pH at which the indicator changes colour (the endpoint). (G3) On the HCI curve, mark the point corresponding to the pH determined in step (G2) for each indicator. Record the volume of NaOH added to achieve this pH. O 3

Worksheet: Experiment 11 Acid-Base Indicators 120 USE THE TABLE IN THE LAB MANUAL (p57), YOUR TITRATION CURVE AND ANALYSIS RESULTS FOR REACTION BETWEEN HCI AND NaOH ANSWERS THE FOLLOWING QUESTIONS Q9-11 Q9. Comment on the suitability of the two indicators for giving an accurate endpoint? Why? 1 mark Q10. Is there a another more suitable indicator for the titration? Use the table of indicators and give a reason for your choice. 1 mark 2020 Online Expt 11 Modified by Ledia Hidalgo for CHEF

Worksheet: Experiment 11 Acid-Base Indicators Q11. Is the equivalence point, for this reaction, at pH 7? Why? 120 2020 Celine Expt 11: Modified by Leda dago for CHESAPL 1 mark Q12. Sketch the expected shape of a titration curve for the standardization of 0.1 M NH3 with 0.1 M HCI, with volume of HCI added on the X-axis and pH on the Y-axis Indicate the approximate pH at the start and end of the titration. Mark on your curve the equivalence point and its approximate pH. Suggest a suitable indicator for this titration, giving a reason for your choice. Take a photo of your drawing and insert it into your work document. 2 marks Focus ENG US