pH of HF
5.3*10^-2 M
Since HF is a weak acid, I'm not sure which formula to use.
Which formula to use for what?
to find pH
Go through the Ka for a weak acid, solve for H^+, then convert to pH.
There is no Ka value that I see, so I looked it up hopefully on a reliable website. Could you make sure this is correct? 7.2*10^-4
so would this problem be:
7.4*10^-4= (x^2)/ (5.3*10^-2-x)
On a calculator I get this number 0.006262587 then -log to get 2.20
It is 7.2E-4 = Ka in my text (about 15 years old). Yes, that's what you do but I didn't end up with your numbers. I solved the quadratic and obtained 0.00583M for x and the pH = 2.23
If you neglect x (which is what you did) you should obtain 6.18E-3. Your value is a little higher than that because you punched in 7.4 and not 7.2.
To find the pH of a weak acid like HF, we need to use the formula for calculating the pH of weak acids.
HF is a weak acid because it only partially dissociates in water. Its dissociation equation can be written as:
HF (aq) ⇌ H+ (aq) + F- (aq)
To find the pH of HF, we need to determine the concentration of H+ ions (also known as the hydrogen ion concentration).
Given that the concentration of HF is 5.3 * 10^-2 M, we can assume that the initial concentration of H+ ions is negligible compared to the concentration of HF.
Since HF is a weak acid, we can use the equilibrium expression for its dissociation reaction:
[H+][F-] / [HF]
Remember that the concentration of H+ ions and F- ions will be equal since the stoichiometric coefficient in the balanced equation is 1:1.
So, let's represent the concentration of H+ ions as x and the concentration of HF as (5.3 * 10^-2 - x) since it will lose some HF molecules to the dissociation reaction. Similarly, the concentration of F- ions will also be x.
Now, we can substitute the concentrations into the equilibrium expression and solve for x:
x * x / (5.3 * 10^-2 - x) = Ka (Ka is the acid dissociation constant for HF)
The value of the acid dissociation constant (Ka) for HF is approximately 6.8 * 10^-4 at 25°C.
Solving this quadratic equation will provide us with the value of x, which represents the concentration of H+ ions. Once we have that value, we can calculate the pH using the formula:
pH = -log[H+]
However, solving this quadratic equation can be quite complex without the use of a calculator or computer software. Thus, we typically rely on approximations or use numerical methods to find the value of x and subsequently the pH.
In this case, assuming that the initial concentration of HF is much larger than the concentration of H+ ions at equilibrium (which is a reasonable approximation for weak acids), we can simplify this equation:
x * x / (5.3 * 10^-2) ≈ Ka
Simplifying further:
x * x ≈ Ka * (5.3 * 10^-2)
Solving this equation, we can find the value of x, which represents the concentration of H+ ions. Once we have x, we can calculate the pH using the formula:
pH = -log[H+]
Using this simplified approach, we can find an approximate value for the pH of HF without extensive calculations.