Calculate the [OH−] and the pH of a solution with an [H+]=5.3×10−10 M at 25 °C. [OH−]=
pH=

Calculate the [H+] and the pH of a solution with an [OH−]=3.6×10−5 M at 25 °C .
[H+]=
pH=

Calculate the [H+] and the [OH−] of a solution with a pH=6.92 at 25 °C.
[H+]=
[OH−]=

Please help quickly! I am being timed!
Thank you!

Answerthis solution will not form a precipitateexplanationsolubility equilbrium:         pbf₂(s)  ⇄ pb²⁺(aq) + 2 f⁻(aq)the solubility product expression is    (we exclude solids from the equilibrium expression).now suppose we have the trial ion product q, which we can define as    then, we can interpret  q as the product of the concentration of the ions that we have and ksp as the product of the ion concentrations that are needed to form a saturated solution.if q > ksp, then that implies that there will be a surplus of ions that exceed the amount that is required to form a saturated solution. a precipitate will then form (the precipitate continually  forms until the surplus ions are removed and q is decreased to ksp's value).if q = ksp, the saturated solution that forms is barely saturated and the precipitate will be minimal.if q < ksp, then the ion concentration will not be enough to form a saturated solution (so no precipitate).using our information, we calculate the trial ion product:     since q = 2.7  × 10⁻⁸ < ksp =  4  ×  10⁻⁸, a precipitate will not form.

We can't until we know what we are working with or on

The answer to your question is c. mitochondria.