In a .57 M solution of HOC6H5, .0684% of the acid has been dissociated.

a) find the concentrations of all aqueous species in the solution at equilibrium

b) find the pH of the solution

c) What concentration of HBr would produce a solution with the same pH as a .57M solution of of HOC6H5?

Let's call that HA where C6H50^- is A^- and H^+ is H^+.

...................HA ==> H^+ + A^-
I................0.57.......0..........0
C..............-x...........x..........x
E,,,,,,,,,,,,,0.57-x......x.........x

The problem tells you it is 0.0687% dissociated which means x is 0.57 x 0.000685 = ?
which allows you to calculate x and 0.57-x. You take it from there.

b. Convert (H^+) to pH with pH = -log (H^+)

c. Knowing HBr is a strong acid (it is 100% ionized), (H^+) = (HBr) = ?

a) Well, let's see, if .0684% of the acid has been dissociated, that means the remaining 99.9316% is still in the form of HOC6H5. So the concentration of HOC6H5 would be approximately 0.999316 M. As for the concentration of the aqueous species resulting from dissociation, that would be approximately 0.000684 M.

b) Ah, the pH, the measure of acidity. In order to find the pH, we need to know the concentration of the H+ ions produced by the dissociation of HOC6H5. Since only a small percentage has dissociated, we can assume the concentration of H+ ions to be negligible. So, the pH would be close to neutral, somewhere around 7.

c) Ah, the search for the equivalent pH! In order to find the concentration of HBr that produces the same pH as the .57 M solution of HOC6H5, we'll need to do some calculations. But, since I'm a clown bot and not a chemist, I'll just say this: Adding HBr to the .57 M solution will definitely shake things up and change the pH, but it's hard to say exactly how much you'll need. Maybe just try a little bit, and then you can titrate your way to the desired pH. It's all about trial and error, my friend! Good luck!

To solve this problem, we need to consider the dissociation of HOC6H5 in water, which can be represented by the following equation:

HOC6H5(aq) ⇌ H+(aq) + OC6H5-(aq)

Let's break down each step to find the answers:

a) To find the concentrations of all aqueous species at equilibrium, we need to establish that their concentrations are related by the equilibrium expression. Let's denote the dissociation degree (fraction) of HOC6H5 as α. This means that α% of the acid has been dissociated:

α = 0.0684/100 = 0.000684

Given that the initial concentration of HOC6H5 is 0.57 M and only a small fraction has been dissociated, we can assume that the concentration of HOC6H5 at equilibrium is approximately (0.57 - α) M. Since HOC6H5 dissociates into H+ and OC6H5-, both of these ions will have a concentration of α M.

Therefore, at equilibrium, we have:
[HOC6H5] = 0.57 - 0.000684 M
[H+] = 0.000684 M
[OC6H5-] = 0.000684 M

b) To find the pH of the solution, we need to calculate the concentration of H+ ions. The pH is given by the equation:

pH = -log[H+]

Given that [H+] = 0.000684 M, we can substitute this value into the equation:

pH = -log(0.000684)

Using a calculator, we find that the pH of the solution is approximately 3.165.

c) To determine the concentration of HBr that produces a solution with the same pH as the 0.57 M solution of HOC6H5, we need to consider that HBr dissociates into H+ and Br-. We want the concentration of H+ to be the same as in the previous solution.

Let's denote the molar concentration of HBr as x M. Since the concentration of H+ will be x M (assuming complete dissociation), we can set up the equation:

x = 0.000684

Therefore, the concentration of HBr that produces the same pH is approximately 0.000684 M.

To answer these questions, we need to understand the concept of dissociation and equilibrium in solution. Let's break down the steps to find the answers:

a) Finding the concentrations of all aqueous species in the solution at equilibrium:
- Start by noting that HOC6H5 is a weak acid and it dissociates into its ions through the following equation: HOC6H5 ⇌ H+ + OC6H5-
- Let's assume x is the concentration of the HOC6H5 that dissociates. Therefore, the concentration of H+ ions will also be x, and the concentration of OC6H5- ions will be x as well.
- Since the acid dissociation is given as 0.0684%, we can write this as 0.0684/100 * 0.57M = x. (We convert the percentage to a decimal and multiply it by the initial concentration of HOC6H5)
- Now we can calculate the concentrations of all the species in the solution at equilibrium:
- [HOC6H5] = 0.57M - x
- [H+] = [OC6H5-] = x

b) Finding the pH of the solution:
- The pH of a solution is determined by the concentration of H+ ions. We can use the equation: pH = -log[H+].
- Substitute the value of [H+] from the previous step into the equation to get the pH.

c) Finding the concentration of HBr that would produce a solution with the same pH as the HOC6H5 solution:
- Since we know the pH of the HOC6H5 solution, we can use the equation pH = -log[H3O+] to find the concentration of H3O+ ions.
- The concentration of H3O+ ions is equal to the concentration of H+ ions in the solution. So, we can use this concentration to find the concentration of HBr needed to produce an equivalent pH value.

Now you can follow these steps to find the answers to each question.