0
$\begingroup$

For school, I am doing a small study on a chemical reaction involving 2 M Sulfuric Acid (H₂SO₄) and Zinc (Zn). I have done some background research on the process to calculate the pH of a solution, but the expected value doesn't seem to match either the initial pH of the H₂SO₄, or the change in pH over time.

For reference, we calibrated our pH probe multiple times to an accuracy within ± 0.1 of the expected values of the calibration solutions (pH 4 and pH 7). Water displayed a pH of 7.06.

According to the probe, the pH of 2 molar H₂SO₄ is ≈ -0.25 ± 0.03, which I do not understand. Research indicates that the pH of a solution can be calculated using $-log_{10}(\ce{H^+})$. When completely dissociated in water, 2 molar sulfuric acid should have a H⁺ value of 4, so the pH should be equal to $-log_{10}(4)$, which is ≈ -0.602, over 0.35 off from the measured value of -0.25.

Also, as the reaction between Zinc and Sulfuric Acid progressed, the pH of the solution decreased (from ≈ -0.25 to ≈ -0.30). This also makes no sense to me, as my intuition would tell me that the pH would increase, not decrease. This is the main result I am seeking answers for.

If anyone could explain this to me, it would be incredibly helpful for my study.

Improve this question
$\endgroup$
7
  • 3
    $\begingroup$ HSO4^- is not strong acid. pH is given by H+ activity, not concentration. At such ionic strength I, activity coefficients raise with I. Dissolving Zn raises ionic strength, which may cause stronger activity coefficient raise than is the decrease of H+ concentration. Note that generally, calculations at these concentrations may be very misleading. $\endgroup$
    – Poutnik
    Aug 28 at 6:10
  • 2
    $\begingroup$ Chem+Math Expression formatting reference: MathJax Basics / Chem+Math expressions/formulas/equations / Upright vs italic / Math SE Mathjax tutorial // MathJax is preferred not to be used in CH SE Q titles. $\endgroup$
    – Poutnik
    Aug 28 at 9:27
  • 4
    $\begingroup$ Why you'd expect it to work well below 0 pH? You shouldn't even put glass electrodes in extreme solutions. If it was a base, you could even damage it. $\endgroup$
    – Mithoron
    Aug 28 at 13:37
  • 1
    $\begingroup$ I agree with Mithoron, I forgot to mention that. $\endgroup$
    – Poutnik
    Aug 29 at 8:51
  • 1
    $\begingroup$ @Matt The meaning of the "Extreme solution" is very much context dependent. For theoretical prediction of pH it IS an extreme solution as it is extremely difficult to calculate it. One has to use quite complicated models of thermodynamic activity of H+(aq). $\endgroup$
    – Poutnik
    Aug 31 at 13:53

2 Answers 2

Reset to default
1
$\begingroup$

The first consideration is the equipment, a calibration variance of 0.1 is unacceptable. The reference electrode should be a flowing double junction electrode, the outer chamber fill solution might need some research. Calibrate at pH7; set the slope at pH4; record all mv values. The difference between7 and 4 should be almost 180 mv. pH 2 buffer is available get some and check for linearity.

Accurately prepare 2M, 1M and 0.5M sulfuric acid and measure the mv readings and the pH of each. Find the Ka for bisulfate and try to calculate the possible [H3O+]. Make calibration curves for both concentrations. Then add an amount of Zn to give a calculated molarity change to each solution. Hopefully to change each solution into the lower one in sulfuric acid. If the values fit one of the calibration curves the activity coefficients are reasonably constant; if not they must be researched. Good luck the results should be very interesting. Please report back

Improve this answer
$\endgroup$
2
  • $\begingroup$ Unfortunately, our pH probe only supports two point calibration to my knowlege. Not only that, but we have no time left to collect any further results. Does this mean that our results, particularly the negative pH, are wrong and due to the pH probe measuring incorrectly? $\endgroup$
    – Matt
    Aug 29 at 2:03
  • 1
    $\begingroup$ The negatve is right you do not know how negative. pH meters are calibrated at only ONE point usually the isoelectric point near pH 7. The pH 4 point sets the slope done with the temperature control not the calibration control. Find the time and try to do the experiment properly. $\endgroup$
    – jimchmst
    Aug 29 at 21:25
0
$\begingroup$

I am also working with Matt on this experiment, we are unable to gain any further results, due to time constraints. As a school experiment, our teachers provided us with the option of using 2M Sulfuric acid and Zinc. From what I gathered from your responses, is that these reactants are not ideal at all. We are using the SparkVUE app paired with the PASCO pH probe. We calibrated the probe with two points, under the careful supervision of our teacher, and Lab Technicians at our school who confirmed that it is the correct way to calibrate. The results that we gathered are negative and decrease after the zinc is put into the sulfuric acid. What we are trying to find out is why the zinc is causing the pH to decrease, and what the trend of the data means. Thankyou for your helpful responses Poutnik, Mithoron, and jimchmst. average change in pH over time with zinc in soution

This is the graph of the change in pH over time for reference.

Improve this answer
$\endgroup$
5
  • $\begingroup$ @Menton, could you check the pH range in which your electrode is reliable? The pH value you are trying to measure seems too high. Can you try your experiment with a lower pH? For example, using a 0.1 M $\ce{H2SO4}$ solution. $\endgroup$
    – PAEP
    Aug 31 at 17:03
  • $\begingroup$ no, we are unable to do anymore tests. and we were not given an option for a lower concentration. A different group asked for 0.5M but were not given it. $\endgroup$
    – Brenton
    Sep 2 at 7:40
  • $\begingroup$ @Menton, I don't usually work with pH meters but I'm not sure that if you work very far out of the calibration range of the electrode, the pH meter readings can be very reliable. Maybe you need different calibration points to be able to measure such low pHs. $\endgroup$
    – PAEP
    Sep 2 at 16:39
  • $\begingroup$ You may find this document helpful. A guide to pH measurement (Metler-Toledo) : mt.com/mt_ext_files/Editorial/Generic/1/… $\endgroup$
    – PAEP
    Sep 2 at 16:40
  • $\begingroup$ I understand you might not be able to do so, but the right way to proceed would be a negative control: perform the exact same measurement but without adding the zinc powder. The pH decreasing during the reaction is clearly anomalous, and so most likely an experimental artefact. As several people have pointed out, regular pH electrodes were not designed for highly acidic solutions like this, so its' very possible that the membrane's properties slowly change, becoming more sensitive to protons over time. If so, you'd see the same in the blanco experiment, except more so as no hydrogen is evolved $\endgroup$
    – Randy Lambreghts
    Sep 5 at 13:50

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service and acknowledge that you have read and understand our privacy policy and code of conduct.

Not the answer you're looking for? Browse other questions tagged or ask your own question.