Aspirin equilibrium is : HC9H7O4(aq) + H2O (double arrow)C9H7O4^-(aq) + H3O^+
You have 2 tablets each containing 0.325g of active ingredient that you dissolve in water up to 225ml.
Calculate the pH.
Let's simplify aspirin formula to HA.
HA + H2O ==> H3O^+ + A^-
Ka = (H3O^+)(A^-)/(HA)
M = (2*0.325)/molar mass HA
Set up an ICE chart, substitute into the Ka expression and solve for x. Convert that to pH. I assume you know Ka or you can look it up.
To calculate the pH of the solution, we need to determine the concentration of H3O+ ions in the solution. We can start by finding the moles of the active ingredient (aspirin) in the tablets, and then calculate the concentration of H3O+ ions.
1. Find the moles of aspirin:
Given: Mass of aspirin in each tablet = 0.325 g
Molar mass of aspirin (C9H7O4) = 180.16 g/mol
Number of moles of aspirin = (mass of aspirin)/(molar mass of aspirin)
= 0.325 g / 180.16 g/mol
2. Calculate the concentration of H3O+ ions:
The reaction equation shows that for every 1 mole of aspirin (C9H7O4), 1 mole of H3O+ ions is produced.
The volume of the solution is 225 mL = 0.225 L
Molarity (M) = (moles of solute)/(volume of solution in liters)
= (moles of H3O+ ions)/(0.225 L)
3. Calculate the pH:
The pH is a measure of the concentration of H3O+ ions. It is given by the equation:
pH = -log[H3O+]
Now that we have the concentration of H3O+ ions, we can calculate the pH using the above equation.
Note: The equilibrium between undissociated aspirin (HC9H7O4) and the dissociated form (C9H7O4-) plays a role in determining the actual pH of the aspirin solution. However, since the problem statement does not provide any information about the equilibrium constant or any other necessary values, we are assuming that the equilibrium is already established and that the concentration of H3O+ can be calculated directly from the number of moles of aspirin.
Please let me know if there are any more details or clarifications needed to provide a more accurate answer.