pH and pOH

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.