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Introduction

1. fill burette with deionized water, NaOH, HCl + white tile 2. 1st beaker: 25 HCl + 10 water into beaker 3. 2nd beaker: 25 NaOH +10 water 4. 3rd beaker: 25 buffer + 2 drop of indicator Add more indicator if too pale No of drop 5. Add same no of drop of indicator to 1st and 2nd 6. Dilute HCl with water, vol same as 3rd Dilute NaOH with water, vol same as 3rd 7. Add NaOH 1cm3 at a time to 3rd, mix 8. Measure pH with pH meter just when color change (compare to 1st) 9. Add NaOH 1cm3 at a time to 3rd, mix 10. Measure pH until color change is complete (compare with 2nd ) 11. Repeat with other indicator Indicator pH ( color change start) pH (color change complete) Color 1 Color 2 S.K.H. Lam Woo Memorial Secondary School F.7 Chemistry Laboratory Report Name: Chan Ching Wai Class: F.7H (2) Experiment 4: Indicator Date of Experiment: 16-11-2010 Objective To determine the indicator range of some acid-base indicators Introduction In this experiment, the indicator ranges of some acid-base indicators were determined. Indicators are chemicals that would change color as the pH of the solution in which they are dissolved changes within the indicator range. Indicators are commonly used in acid-base equilibrium in order to determine the concentration of a solution. ...read more.

Middle

In this process, the solution in beaker 1 and beaker 2 could be diluted from time to time in order to give fair comparison with color of beaker 3 with the same volume. Results Indicator pH ( color change start) pH (color change complete) Color 1 Color 2 Litmus 6.8 9.9 Pink Blue Phenolphthalein 7.1 10.4 Colorless Pale pink Methyl red 3.3 5.3 Red Pale yellow Methyl orange 4.3 5.6 Red Pale yellow Bromophenol blue 2.6 3.8 Pale yellow Pale purple Discussion Measured in buffer solution The indicator range of indicator was determined by adding sodium hydroxide solution to buffer solution in beaker 3. Beaker 3 should be filled with buffer solution but not deionized water for easy measurement of pH at particular point. To begin with, a buffer solution is a solution which tends to resist pH change when it is diluted or a small amount of acid or bases is added to it. With the ions inside the buffer solution, the acid and base added would be converted into weak acid or weak bases which would only slightly ionized into water. This was because the buffer solution is made of weak acid or weak base with its salt. The salt in buffer solution completely ionized in water to give the salt of the acid or base and the other oppositely charged ion. ...read more.

Conclusion

Thus, the determination of color change is important. Yet, determination by human eyes is not accurate and quite objective. To improve this situation, the color of beaker 3 could be compared with beaker 1 and 2 as stated above. Also, white tile could be placed under the beakers to make the color change more conspicuous. Thirdly, equal number of drops of indicators was added to the three beakers to have fair comparison of the colors. If more drops of indicator are added, the color of the solution would be deeper and did not resemble the initial color of beaker 3. This may lead to an inaccurate determination of the point of color change and a measurement of pH at wrong time. Fourthly, as the sodium hydroxide solution was added 1 cm3 each time, the accuracy could only be corrected to 1 cm3. If color change occurred when an additional amount of 0.5 cm3 added to the solution, an extra amount of 0.5cm3of sodium hydroxide solution may already have been added. Thus, a higher pH would be measured. Therefore, sodium hydroxide solution could be added drop by drop when the color was about to change or stop to change. The pH meter should also be dipped immediately when color started to change or stopped to change so as to measure the pH exactly at these points. Conclusion In this experiment, the indicator range of the following indicators were measured. Indicator Indicator range Litmus 6.8-9.9 Phenolphthalein 7.1-10.4 Methyl red 3.3-5.3 Methyl orange 4.3-5.6 Bromophenol blue 2.6-3.8 ...read more.

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