Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: defencegrid on January 31, 2017, 12:21:59 AM
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I performed a titration with NaOH and HCl. When I tested the pH of the NaOH/HCl solution which had reached the equivalence point, the pH was 1.1. Shouldn't it be about seven?
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Something gone wrong. Real pH of 0.1 M NaCl is ~6.9. May be you used water instead of NaOH?
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Something gone wrong. Real pH of 0.1 M NaCl is ~6.9. May be you used water instead of NaOH?
Yeah, that's what I thought. I must of done something wrong.
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I performed a titration with NaOH and HCl. When I tested the pH of the NaOH/HCl solution which had reached the equivalence point, the pH was 1.1. Shouldn't it be about seven?
Please elaborate on what you did, as what you wrote doesn't make much sense - how do you know you have reached the equivalence point, if that's typically done by controlling the pH, and your pH is so off?
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1M HCl in the burette and 10ml of NaOH in the conical flask with two drops of phenolphthalein. Titrated until it turned colourless.
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Check your pH-meter.
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Also, consider the possibility that you overshot the titration. You should w to the palest pink color, a colorless state that returns to the palest pink when swirled, then the tiniest drop prevents any color from returning -- that's the endpoint. Check it yourself, start with a neutral NaCl solution, take its pH, add 1 drop of HCl, is the pH still 7?