4469.6 g of 10% H2SO4 mixed with 1047 g of 45% KOH and 70 L of water (total volume is ~75L), should give a pH of around 2.....however in real life I get a pH of around 1.5-1.6....what is going on here?
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4469.6 grams of 10% H2SO4 mixed with 1047 grams of 45% KOH and 70 L of water (so total volume is ~75L), calculated should give a pH of around 2.....however in real life I get a pH of around 1.5-1.6....what is going on here????
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The calculations are:
4469.6 grams of 10% H2SO4 = 446.95 g/98g per mole = 4.561 moles of H2SO4, 2 protons per mole = 9.122 moles of H+
1047 grams of 45% KOH = 471.15g/56 g per mole = 8.41 moles of OH-
9.122 – 8.41 = 0.712 moles of protons present
à Final volume of solution is around 75L
0.712/75 = 0.00949 moles of H+ present (per litre)
pH = -log[H+]
pH = 2.02
sure a small difference would occur between theory and the real world, but please, this is a pH calcuation
2007-09-26 04:54:45 · 4 answers · asked by Anonymous in Science & Mathematics ➔ Chemistry
are you just going to talk around the answer since you don;t know it? For these calculations, all the BS you mention is negligible. According to you, how can anyone calculate anything since there are always errors involved? Gawd, have you ever even been in a lab before?
2007-09-26 05:12:03 · update #1
You're not correcting your calculations for ionic strength, nor are you calibrating your pH meter using buffers of the same ionic strength.
pH is *not* -log[H+], it is -log{H+} and {H+} = a_H+[H+]
where the { } imply activity, not concentration, and a_H+ is the activity of the hydrogen ion. a_H+ is a strong function of ionic strength and these corrections are significant, especially at the ion concentrations in your solution.
Your solution has a relatively high ionic strength, from the H+ concentration alone. Plus you have a lot of counter ions in there. You need to do some serious correcting for activities in high ionic strength media to get reality and calculations to agree. This is in addition to getting buffers of approriate ionic strength since a pH meter *does* measure -log{H+} and the ionic strength of the buffer will make a difference.
Pick up a graduate level text on aquatic chemistry, Stumm and Morgan (Aquatic Chemistry) and Pankow (Aquatic Chemistry Concepts) are very good books, and they will explain how to do all this.
2007-09-26 05:54:28 · answer #1 · answered by gcnp58 7 · 0⤊ 0⤋
"but please, this is a pH calcuation"
What ? Are pH formulae sacrosanct and inviolable ? Just take a look at all the assumptions made while deriving the formulae you have used. Even the relation pH = -log[H+] is strictly wrong as it is really H3O+.
Also review your experiment to be sure that no errors crept in. Generally a deviation of 1 to 5 % is acceptable depending on the experiment's sensitivity.
2007-09-26 05:02:38 · answer #2 · answered by ag_iitkgp 7 · 0⤊ 0⤋
Did you calibrate your pH meter? You should bracket the targeted pH with standard buffer solutions (to measure a pH of 2.02, you should use pH 1 and 4 standard buffers)
What was the temperature of your solution when you measured pH? Actual readings drop off at very low and very high pH.
2007-09-26 05:04:50 · answer #3 · answered by JOhn M 5 · 0⤊ 0⤋
not extremely. I look greater outgoing in this community than in the real international. i'm afraid i could say some thing that would scare somebody while face 2 face. The distances in contact in this community i think as though that's not some thing i'd desire to fret approximately.
2016-12-17 10:50:02 · answer #4 · answered by hinokawa 4 · 0⤊ 0⤋