True or false question

The pH of 0.1mol/L acetic acid is 1.
So the amswer is False, the pH is greater than 1.
How is pH-log(0.1)=1 wrong?

I researched and found this solution but I don't know why I cant use pH=-log ([H+])

Ka=1.8×10^-5
1.8=x^2÷x
I get x=1.34 x 10^-3
pH=-log (1.34 x 10^-3)=2.87

I meant to type 1.8x10^-5=x^2÷x

I applaud your persistence as well as you resourcefulness. The answer you obtained for (H^+) of 1.34E-3 is correct for a 0.1M solution of acetic acid. You CAN(you really can) use pH = -log(H^+). What you aren't getting is that the (H^+) is 1.34E-3M and NOT (IT IS NOT, REPEAT, NOT) 0.1M The 0.1M is the concentration of the acetic acid and NOT the concentration of (H^+). That's because acetic acid is a WEAK acid and it is not ionized 100%. Solutions of 0.1M HCl are pH = 1 because (H^+) = 0.1. Weak acids don't ionize 100% so their (H^+) is less than 0.1 M for 0.1M solutions. From these data you can see that 0.1M acetic acid solutions only ionize 1.34%. Again, there is nothing wrong with pH = -log(H^+). The problem is that you are using (HAc) for (H^+) and for weak acids that isn't true. Hope this helps.

Using the equation pH = -log[H+], where [H+] represents the concentration of hydrogen ions, is a common way to calculate the pH of a solution. However, it's important to ensure you are using the correct concentration value for [H+].

In the case of acetic acid (CH3COOH), it is a weak acid and partially dissociates in water. The dissociation reaction is CH3COOH ⇌ H+ + CH3COO-. The concentration of [H+] is not equal to the concentration of the acetic acid solution (0.1 M), but rather depends on the equilibrium constant Ka and the initial concentration of the acid.

To find the concentration of H+ in the solution, you need to use the Ka value for acetic acid, which you mentioned is 1.8x10^-5. By setting up the equilibrium expression: Ka = [H+][CH3COO-]/[CH3COOH], where the concentration of CH3COOH is known (0.1 M), you can solve for [H+].

Plugging in the values, you get 1.8x10^-5 = [H+](0.1 - [H+]). Rearranging the equation and solving for [H+], you find that [H+] = 1.34x10^-3 M.

Now, you can use the equation pH = -log[H+] to calculate the pH. Plugging in the value of [H+], you correctly found that the pH is approximately 2.87, which is greater than 1.

Therefore, the statement "The pH of 0.1 M acetic acid is 1" is false, and you correctly determined that the pH is greater than 1. The discrepancy might have been due to mistakenly taking the concentration of acetic acid as the concentration of H+.