Acid dissociation constant

In <a href="/facts/Chemistry/SrNqkGVm">chemistry</a>, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted ⁠
 
 
 
 
 K
 
 a
 
 
 
 
 {\displaystyle K_{a}}
 
⁠) is a <a href="/facts/Quantitative_property/C1t5h075">quantitative</a> measure of the <a href="/facts/Acid_strength/vJ4rj6Lu">strength</a> of an <a href="/facts/Acid/FySU18n1">acid</a> in <a href="/facts/Solution_(chemistry)/MGchhMQV">solution</a>. It is the <a href="/facts/Equilibrium_constant/yMTuuBbX">equilibrium constant</a> for a <a href="/facts/Chemical_reaction/GpqWlKrz">chemical reaction</a>

HA
 
 
 
 
 
 
 ↽
 
 
 
 
 −
 
 
 
 
 
 
 
 
 −
 
 
 
 
 ⇀
 
 
 
 
 
 
 
 A
 
 −
 
 
 +
 
 H
 
 +
 
 
 
 
 
 {\displaystyle {\ce {HA <=> A^- + H^+}}}

known as <a href="/facts/Dissociation_(chemistry)/gDRGdeSN">dissociation</a> in the context of <a href="/facts/Acid%E2%80%93base_reaction/z8pbL0X8">acid–base reactions</a>. The <a href="/facts/Chemical_species/A8BxM3oW">chemical species</a> HA is an <a href="/facts/Acid/FySU18n1">acid</a> that dissociates into A−, called the <a href="/facts/Conjugate_base/hohgQ07M">conjugate base</a> of the acid, and a <a href="/facts/Hydron_(chemistry)/v5Qkumvm">hydrogen ion</a>, H+. The system is said to be in <a href="/facts/Chemical_equilibrium/y8ToYXHp">equilibrium</a> when the concentrations of its components do not change over time, because both forward and backward reactions are occurring at the same rate.
The dissociation constant is defined by

K
 
 a
 
 
 =
 
 
 
 [
 
 A
 
 −
 
 
 ]
 [
 
 H
 
 +
 
 
 ]
 
 
 [
 H
 A
 ]
 
 
 
 ,
 
 
 {\displaystyle K_{\text{a}}=\mathrm {\frac {[A^{-}][H^{+}]}{[HA]}} ,}
 
 or by its <a href="/facts/Logarithm/QeXaV97q">logarithmic</a> form

p
        
        
          K
          
            
              a
            
          
        
        =
        −
        
          log
          
            10
          
        
        ⁡
        
          K
          
            a
          
        
        =
        
          log
          
            10
          
        
        ⁡
        
          
            
              
                [
                HA
                ]
              
            
            
              [
              
                
                  A
                  
                    −
                  
                
              
              ]
              [
              
                
                  H
                  
                    +
                  
                
              
              ]
            
          
        
      
    
    {\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log _{10}K_{\text{a}}=\log _{10}{\frac {{\ce {[HA]}}}{[{\ce {A^-}}][{\ce {H+}}]}}}

where quantities in square brackets represent the <a href="/facts/Molar_concentration/P2WL5b9q">molar concentrations</a> of the species at equilibrium. For example, a hypothetical weak acid having Ka = 10−5, the value of log Ka is the exponent (−5), giving pKa = 5. For <a href="/facts/Acetic_acid/uyBKA2Pd">acetic acid</a>, Ka = 1.8 x 10−5, so pKa is 4.7. A higher Ka corresponds to a stronger acid (an acid that is more dissociated at equilibrium). The form pKa is often used because it provides a convenient <a href="/facts/Logarithmic_scale/CvIGK5TQ">logarithmic scale</a>, where a lower pKa corresponds to a stronger acid.

Acid dissociation constant open-in-new

Acid dissociation constant