I'm having trouble getting the net ionic equation when HI is added to a buffered solution of HOBr. And the net ionic when HI is added to LiOBr.
From those results I'm supposed to compare why one pH drops more than the other.
Here is the whole 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!
To determine the net ionic equation when HI is added to a buffered solution of HOBr, we need to start by writing the balanced equation for the reaction between HI and HOBr. This balanced equation will help us identify the species involved in the net ionic equation.
The balanced equation for the reaction between HI and HOBr is:
HI + HOBr -> HBr + HIO
In a buffered solution, there is usually an equilibrium between an acid and its conjugate base. In this case, HOBr can act as an acid and OBr- can act as its conjugate base. The net ionic equation occurs only between the species involved in the reaction, excluding spectator ions that appear on both sides of the equation.
To write the net ionic equation, we need to identify the ions that are directly involved in the reaction. In this case, the net ionic equation for the reaction between HI and HOBr can be written as:
H+ + OBr- -> HBr + IO-
Now let's consider the reaction between HI and LiOBr to determine its net ionic equation.
The balanced equation for the reaction between HI and LiOBr is:
HI + LiOBr -> HBr + LiI + OBr-
In this case, the net ionic equation can be written as:
H+ + OBr- -> HBr + IO-
Now, to compare why one pH drops more than the other when HI is added to these buffered solutions, we need to consider the relative strength of the acids involved.
HOBr is a weak acid, while HI is a strong acid. When HI is added to the buffered solution of HOBr, the pH drops more significantly because HI is a strong acid, meaning it dissociates completely in solution to release more H+ ions. This increase in H+ concentration leads to a larger decrease in pH.
On the other hand, LiOBr is a salt of a weak acid (HOBr) and a strong base (LiOH). When HI is added to the buffered solution of LiOBr, the pH drops to a lesser extent compared to the buffered solution of HOBr because the weak acid nature of HOBr already partially neutralizes the added H+ ions from HI. Therefore, there is less of a change in pH when HI is added to the LiOBr buffered solution compared to the HOBr buffered solution.
In summary, the difference in the pH drop between the two solutions can be attributed to the relative strengths of the acids and the buffering capacity provided by the weak acid in the HOBr solution.
I'm not certain I understand; however, if you have a buffered solution with HOBr/OBr^-, then,
OBr^- + H^+ ==> HOBr
I assume the second one is not buffered so you are adding a strong acid to the salt of a weak acid and it's the same net ionic equation as the first one.
If you posted the entire question we could help better.