DrBob222 thank you for your previous response. I am reposting the entire question.
Two different 1.2 L buffered solutions were prepared using HOBr and LiOBr. Both buffered solutions had a pH of 5.2 at 25C. After 0.17 mols of HI were added to each of the solutions, it was found that the pH of one solution had dropped to 4.9 and the pH of the other had dropped to 3.1.
a. What is the balanced net ionic equation for the reaction that occurred when HI was added to these buffered solutions?
b. What trend must be true when comparing the concentrations of HOBr and OBr- in the two solutions is they shared the same pH before the HI was added?
Thank you!
a. The balanced net ionic equation for the reaction that occurred when HI was added to these buffered solutions is: HOBr + HI → H2O + I- + Br-.
b. The trend that must be true when comparing the concentrations of HOBr and OBr- in the two solutions is that the solution with the lower pH (3.1) must have had a higher concentration of OBr- than the solution with the higher pH (4.9).
To find the balanced net ionic equation for the reaction when HI is added to the buffered solutions, we need to first examine the initial components and their formulas.
The buffered solutions consist of HOBr (hypobromous acid) and LiOBr (lithium hypobromite). When HI (hydroiodic acid) is added to the solutions, it reacts with the components of the buffer according to the following net ionic equation:
HI + OBr- → HBr + IO-
In this equation, one molecule of hydroiodic acid (HI) reacts with one ion of hypobromite (OBr-) to produce one molecule of hydrogen bromide (HBr) and one ion of iodite (IO-).
Now, let's move on to the second part of the question.
To compare the concentrations of HOBr and OBr- in the two solutions before the addition of HI, we need to consider the pH values and the effect of adding HI on the pH.
A lower pH value indicates a higher concentration of hydrogen ions (H+), which in turn means a higher concentration of the corresponding acid. In this case, a pH drop from 5.2 to 4.9 suggests that more H+ ions were generated. Since the pH decrease occurred in both solutions after adding HI, we can conclude that the concentration of the HBr (from the reaction of HI and OBr-) increased.
Therefore, if the two buffered solutions had the same pH before HI was added, it means that the concentration of HOBr must have been greater than the concentration of OBr- in both solutions. This is because the addition of HI would have led to the formation of more HBr, causing a decrease in pH.
In summary, the trend for comparing the concentrations of HOBr and OBr- in the two solutions, given the change in pH after adding HI to both solutions, is that the concentration of HOBr is greater than the concentration of OBr-.
a. To determine the balanced net ionic equation for the reaction that occurred when HI was added to these buffered solutions, we need to consider the acid-base reaction between HI and the components of the buffer solution.
The balanced net ionic equation for the reaction can be written as:
HI + OBr- -> HOBr + I-
b. To determine the trend in the concentrations of HOBr and OBr- in the two solutions before the addition of HI, we need to consider Le Chatelier's principle.
In both buffered solutions, the initial pH was 5.2. When HI is added to the solution, it acts as a strong acid that donates protons (H+) to the solution. The increase in the concentration of H+ ions will shift the equilibrium of the reaction (HI + OBr- -> HOBr + I-) to the left, favoring the formation of more OBr- ions and reducing the concentration of HOBr.
Since the pH of one solution dropped to 4.9 and the other dropped to 3.1, it indicates that the solution with a pH drop to 4.9 had a higher initial concentration of OBr- ions compared to the solution with a pH drop to 3.1. Therefore, the trend is that the solution with a higher initial concentration of OBr- will have a smaller drop in pH when HI is added, indicating a higher buffering capacity and a higher concentration of OBr-.