What are the values of [H3O+] and [OH−] for a solution with pH = 3.86?
pH = 3.86
pH + pOH = pKw = 14
You know pH and pKw, solve for pOH.
Then pH = -log(H^+). Solve for (H^+).
and pOH = -log(OH^-). Solve for (OH^-).
To determine the values of [H3O+] and [OH−] for a solution with pH = 3.86, we need to use the concept of the pH scale and the relationship between [H3O+] and [OH−].
The pH scale is a logarithmic scale that measures the concentration of hydronium ions ([H3O+]) in a solution. It is given by the equation:
pH = -log [H3O+]
To find the value of [H3O+], we rearrange the equation:
[H3O+] = 10^(-pH)
Now, substituting the given pH value of 3.86 into the equation:
[H3O+] = 10^(-3.86)
Calculating this using a calculator or by hand, the value of [H3O+] is approximately 1.41 × 10^(-4) M.
Since we know that water is a neutral substance, the product of the [H3O+] and [OH−] concentrations is always equal to 1 × 10^(-14) at 25°C. Therefore, we can find the value of [OH−] by dividing this constant by the value of [H3O+]:
[OH−] = 1 × 10^(-14) / [H3O+]
Substituting the given value of [H3O+] into the equation:
[OH−] = 1 × 10^(-14) / (1.41 × 10^(-4))
Calculating this, the value of [OH−] is approximately 7.09 × 10^(-11) M.
Therefore, for a solution with pH = 3.86, the values of [H3O+] and [OH−] are approximately 1.41 × 10^(-4) M and 7.09 × 10^(-11) M, respectively.