||This article relies largely or entirely upon a single source. (April 2009)|
|Acids and bases|
In chemistry, a weak base is a chemical base that does not ionize fully in an aqueous solution. As Brønsted–Lowry bases are proton acceptors, a weak base may also be defined as a chemical base in which protonation is incomplete. This results in a relatively low pH compared to strong bases. Bases range from a pH of greater than 7 (7 is neutral, like pure water) to 14 (though some bases are greater than 14). pH has the formula:
Since bases are proton acceptors, the base receives a hydrogen ion from water, H2O, and the remaining H+ concentration in the solution determines pH. Weak bases will have a higher H+ concentration because they are less completely protonated than stronger bases and, therefore, more hydrogen ions remain in the solution. If you plug in a higher H+ concentration into the formula, a low pH results. However, pH of bases is usually calculated using the OH- concentration to find the pOH first. This is done because the H+ concentration is not a part of the reaction, while the OH- concentration is.
By multiplying a conjugate acid (such as NH4+) and a conjugate base (such as NH3) the following is given:
By taking logarithms of both sides of the equation, the following is reached:
Finally, multiplying throughout the equation by -1, the equation turns into:
After acquiring pOH from the previous pOH formula, pH can be calculated using the formula pH = pKw - pOH where pKw = 14.00.
Weak bases exist in chemical equilibrium much in the same way as weak acids do, with a base dissociation constant (Kb) indicating the strength of the base. For example, when ammonia is put in water, the following equilibrium is set up:
Bases that have a large Kb will ionize more completely and are thus stronger bases. As stated above, pH of the solution depends on the H+ concentration, which is related to the OH- concentration by the self-ionization constant (Kw = 1.0x10−14). A strong base has a lower H+ concentration because they are fully protonated and less hydrogen ions remain in the solution. A lower H+ concentration also means a higher OH- concentration and therefore, a larger Kb.
NaOH (s) (sodium hydroxide) is a stronger base than (CH3CH2)2NH (l) (diethylamine) which is a stronger base than NH3 (g) (ammonia). As the bases get weaker, the smaller the Kb values become.
Percentage protonated 
As seen above, the strength of a base depends primarily on pH. To help describe the strengths of weak bases, it is helpful to know the percentage protonated-the percentage of base molecules that have been protonated. A lower percentage will correspond with a lower pH because both numbers result from the amount of protonation. A weak base is less protonated, leading to a lower pH and a lower percentage protonated.
The typical proton transfer equilibrium appears as such:
B represents the base.
In this formula, [B]initial is the initial molar concentration of the base, assuming that no protonation has occurred.
A typical pH problem 
Calculate the pH and percentage protonation of a .20 M aqueous solution of pyridine, C5H5N. The Kb for C5H5N is 1.8 x 10−9.
First, write the proton transfer equilibrium:
The equilibrium table, with all concentrations in moles per liter, is
|change in normality||-x||+x||+x|
|equilibrium normality||.20 -x||x||x|
|Substitute the equilibrium molarities into the basicity constant|
|We can assume that x is so small that it will be meaningless by the time we use significant figures.|
|Solve for x.|
|Check the assumption that x << .20||; so the approximation is valid|
|Find pOH from pOH = -log [OH-] with [OH-]=x|
|From pH = pKw - pOH,|
|From the equation for percentage protonated with [HB+] = x and [B]initial = .20,|
This means .0095% of the pyridine is in the protonated form of C5H5NH+.
Other weak bases are essentially any bases not on the list of strong bases.
See also 
- Atkins, Peter, and Loretta Jones. Chemical Principles: The Quest for Insight, 3rd Ed., New York: W.H. Freeman, 2005.