Introduction

• In pure water, [H3O+] and [OH–] are equal to each other.

• When an acid or base is dissolved in water, [H3O+] and [OH–] are not equal to each other.

• Kw = [H3O+] [OH–] = 1.0 × 10–14

• [H3O+] and [OH–] are inversely related to each other.

• Acidic solution: if [H3O+] > 1.0 × 10–7, then  [OH–] < 1.0 × 10–7

  • At room temperature

• Basic solution: if [OH–] > 1.0 × 10-7, then [H3O+] < 1.0 × 10–7

pH

• The acidity and basicity of a solution can be expressed in terms of its [H3O+] or [OH–].

• These concentrations can span many orders of magnitude, therefore logarithmic values are commonly used.

• pH = -log [H+]

• pOH = – log [OH-]

• [H+][OH-]  =  1.0 x 10-14

  • Now taking the log of both sides gives us

  • Log[H+]  +  Log [OH-]  =  -14.    Multiply by -1,  and

  • -log[H+]   - Log [OH-]   =  14.   But – log{H+] is pH!

  • So,  pH +  p OH  MUST = 14   (at room temp.)

pH and pOH in a Neutral Solution

• Neutral solution: if [H3O+] = [OH–] = 1.0 × 10–7

  • At room temperature.

• Like ANY equilibrium constant, the Kw is temperature dependent.  Since ionization is ENDOTHERMIC, the Kw increases as temperature increases.

• If the Kw is NOT 1.0x10-14 , then the pH of a neutral solution will NOT be 7.