Why was C chosen for this question?
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Title: Using the same NaOH solution to titrate solutions of sulfuric acid and acetic acid of equal volume, the same volume of NaOH solution is consumed in each case. This indicates that in the solutions of sulfuric acid and acetic acid, ( ) A. the hydrogen ion concentrations are equal (unit: mol/L, the same below) B. the concentrations of sulfuric acid and acetic acid are equal C. the concentration of sulfuric acid is half that of acetic acid D. the degrees of ionization of sulfuric acid and acetic acid are equal E. the ranges of pH shifts during the titration are the same
If they are all strong acids, then C is the correct answer; that’s what I think
Reply to 2# Juliet: If they are all strong acids, then is E also correct? The concept of the range beyond which something occurs is unclear; please provide guidance!
Classmate Juliet, it would be best for you to take a look at the introductory course on analytical chemistry; it provides quite clear explanations regarding titration. If Brother Romeo has studied analytical chemistry, he can explain it to you very clearly. Furthermore, sulfuric acid is a dibasic acid; at the same molar concentration, it can provide twice as many titratable hydrogen ions as acetic acid. Acetic acid is a weak acid; therefore, the potential or pH shift at the end point of titration is smaller compared to that of strong acids, and the time at which this shift occurs also shifts slightly toward higher pH values.
Write out the reaction equations for the two chemical formulas and take a look for yourself
To be honest, I was a bit confused while working on this problem, but upon careful consideration, if we simplify things by assuming that each reaches its own titration endpoint, then the concentrations can be understood using the concepts from middle school chemistry. Of course, sulfuric acid can’t exist at a “half concentration”,,, :P It’s that simple; I don’t think it’s necessary to get stuck on such details.
Using the process of elimination, I choose C; the specific reason isn’t very clear
I have an explanation for this question in the next issue of Daily Question; was it not written clearly?
Sulfuric acid is a dibasic acid; at the same molar concentration, it can provide twice as many titratable hydrogen ions as acetic acid. Acetic acid is a weak acid; therefore, the potential or pH shift at the end point of titration is smaller compared to that of strong acids, and the time at which this shift occurs also shifts slightly toward higher pH values.
Reply to 8# *ngguanglueguo: Hello, that’s what I think. If the concentration of sulfuric acid is exactly half that of acetic acid, then when sulfuric acid and the added sodium hydroxide neutralize each other completely, the titration endpoint is reached and sodium sulfate is formed. On the other hand, if the same amount of sodium hydroxide is added to the acetic acid solution, a reaction occurs according to the stoichiometric ratio, resulting in sodium acetate being formed; it does not become neutral in this case! Has the acetic acid solution reached the endpoint of titration at this point? I never managed to figure it out, so I ruled out option C at that time. Since the same amount of sodium hydroxide was added, it means that the hydrogen ion concentration in the solutions is the same; it’s just that the hydrogen ions in acetic acid are ionized continuously as the titration proceeds. (Thinking about it later, the hydrogen ion concentration is the acid concentration multiplied by the degree of ionization; it seems they are not the same.)
Reply to 10# Juliet: Acetic acid is titrated using phenolphthalein as an indicator; its color change range corresponds to a weakly alkaline environment. After the titration of acetic acid is complete, the solution remains weakly alkaline. You can calculate the pH value at the endpoint, which falls within phenolphthalein’s color change range.