2 quick questions:

I am stuck on this:
1) The concentration of hydrogen ions in a 0.1 mol dm3 solution of phosphoric acid, H3PO4?
I would be able to do it if I knew the correct equation that = H3PO4 but I can't work it out.

2) In titrations why would more than three significant figures not be appropriate for the final answer?

I worked 1) for you below. Look at your previous posts before posting a duplicate.

2) You are limited to 3 s.f. because of the 0.100. There are three in that. The other values are to four s.f. but you are allowed the least in a series of numbers.

Ahh ok, sorry I had been looking a lot, I just didn't see it so I thought I'd post it again.

Thank you!

For your first question, to determine the concentration of hydrogen ions in a 0.1 mol dm3 solution of phosphoric acid (H3PO4), we need to understand the dissociation of phosphoric acid in water. Phosphoric acid is a triprotic acid, meaning it can release three hydrogen ions (H+) in water.

The dissociation equation for phosphoric acid can be written as:

H3PO4 ⇌ H+ + H2PO4-
H2PO4- ⇌ H+ + HPO42-
HPO42- ⇌ H+ + PO43-

Now, let's analyze the equation step by step:

1) H3PO4 dissociates to form one hydrogen ion (H+) and one H2PO4- ion.
2) H2PO4- dissociates to form one additional hydrogen ion (H+) and one HPO42- ion.
3) HPO42- dissociates to form one more hydrogen ion (H+) and one PO43- ion.

So, in total, each molecule of phosphoric acid (H3PO4) can release three hydrogen ions (H+).

Regarding your second question about titrations, fewer significant figures are typically reported in the final result due to the precision limitations of the technique and measurements involved. Here's why:

1) Titrations involve the use of burettes, which typically have markings to the nearest 0.1 mL. This limits the precision to three decimal places (0.001 mL).
2) Measurements of the volume of the titrant (standard solution) and the volume of the analyte (solution being titrated) are prone to random errors and uncertainties.
3) Additional uncertainties come from the reaction stoichiometry and endpoint determination, which can also introduce errors.

Given these limitations, it is appropriate to express the final titration result with no more than three significant figures. Reporting more significant figures might imply a higher level of precision than is actually achieved in practice.

It's important to note that the number of significant figures in the final result should match the least precise measurement used in the calculation. Keeping consistent significant figures throughout helps maintain the accuracy and precision of scientific data.