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Weak base

In chemistry, a weak base is a chemical base that does not ionize fully in an aqueous solution. As Bronsted-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 level 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). The pH level has the formula:

\mbox{pH} = -\log_{10} \left[ \mbox{H}^+ \right]

Since bases are proton acceptors, the base receives a hydrogen ion from water, H2O, and the remaining H+ concentration in the solution determines the pH level. 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 level results. However, the pH level of bases is usually calculated using the OH- concentration to find the pOH level first. This is done because the H+ concentration is not a part of the reaction, while the OH- concentration is.

\mbox{pOH} = -\log_{10} \left[ \mbox{OH}^- \right]

By multiplying a conjugate acid (such as NH4+) and a conjugate base (such as NH3) the following is given:

K_a \times K_b = {[H_3O^+][NH_3]\over[NH_4^+]} \times {[NH_4^+][OH^-]\over[NH_3]} = [H_3O^+][OH^-]

Since Kw = [H3O + ][OH] then, K_a \times K_b = K_w

By taking logarithms of both sides of the equation, the following is reached:

logKa + logKb = logKw

Finally, multipying throughout the equation by -1, the equation turns into:

pKa + pKb = pKw = 14.00

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 Ionization Constant (Kb) (or the Base Dissociation Constant) 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, the pH of the solution depends on the H+ concentration, which is related to the OH- concentration by the Ionic Constant of water (Kw = 1.0x10-14) (See article Self-ionization of water.) 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. The pie-chart representation is as follows:

  • purple areas represent the fraction of OH- ions formed
  • red areas represent the cation remaining after ionization
  • yellow areas represent dissolved but non-ionized molecules.


Percentage protonated

As seen above, the strength of a base depends primarily on the pH level. 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 level 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(aq) + H_2O(l) \leftrightarrow HB^+(aq) + OH^-(aq)

B represents the base.

Percentage\ protonated = {molarity\ of\ HB^+ \over\ initial\ molarity\ of\ B} \times 100\% = {[{HB}^+]\over [B]_{initial}} {\times 100\%}

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:

\mathrm{H_2O(l) + C_5H_5N(aq) \leftrightarrow C_5H_5NH^+ (aq) + OH^- (aq)}
K_b=\mathrm{[C_5H_5NH^+][OH^-]\over [C_5H_5N]}

The equilibrium table, with all concentrations in moles per liter, is

C5H5N C5H6N+ OH-
initial normality .20 0 0
change in normality -x +x +x
equilibrium normality .20 -x x x

Substitute the equilibrium molarities into the basicity constant K_b=\mathrm {1.8 \times 10^{-9}} = {x \times x \over .20-x}
We can assume that x is so small that it will be meaningless by the time we use significant figures. \mathrm {1.8 \times 10^{-9}} \approx {x^2 \over .20}
Solve for x. \mathrm x \approx \sqrt{.20 \times (1.8 \times 10^{-9})} = 1.9 \times 10^{-5}
Check the assumption that x << .20 \mathrm 1.9 \times 10^{-5} \ll .20; so the approximation is valid
Find pOH from pOH = -log [OH-] with [OH-]=x \mathrm pOH \approx -log(1.9 \times 10^{-5}) = 4.7
From pH = pKw - pOH, \mathrm pH \approx 14.00 - 4.7 = 9.3
From the equation for percentage protonated with [HB+] = x and [B]initial = .20, \mathrm percentage \ protonated = {1.9 \times 10^{-5} \over .20} \times 100\% = .0095\%

This means .0095% of the pyridine is in the protonated form of C5H6N+.


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.
    This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Weak_base". A list of authors is available in Wikipedia.
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