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In chemistry, a weak base is a chemical base that does not ionize fully in an aqueous solution. As 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: Since bases are proton acceptors, the base receives a hydrogen ion from water, H<sub>2</sub>O, and the remaining H<sup>+</sup> concentration in the solution determines the pH level. Weak bases will have a higher H<sup>+</sup> 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<sup>+</sup> concentration into the formula, a low pH level results. However, the pH level of bases is usually calculated using the OH<sup>-</sup> concentration to find the pOH level first. This is done because the H<sup>+</sup> concentration is not a part of the reaction, while the OH<sup>-</sup> concentration is.
By multiplying a conjugate acid (such as NH<sub>4</sub><sup>+</sup>) and a conjugate base (such as NH<sub>3</sub>) the following is given:
Since then,
By taking logarithms of both sides of the equation, the following is reached:
Finally, multipying 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 = pK<sub>w</sub> - pOH where pK<sub>w</sub> = 14.00.
Weak bases exist in chemical equilibrium much in the same way as weak acids do, with a Base Ionization Constant (K<sub>b</sub>) (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 K<sub>b</sub> will ionize more completely and are thus stronger bases. As stated above, the pH of the solution depends on the H<sup>+</sup> concentration, which is related to the OH<sup>-</sup> concentration by the Ionic Constant of water (K<sub>w</sub> = 1.0x10<sup>-14</sup>) (See article Self-ionization of water.) A strong base has a lower H<sup>+</sup> concentration because they are fully protonated and less hydrogen ions remain in the solution. A lower H<sup>+</sup> concentration also means a higher OH<sup>-</sup> concentration and therefore, a larger K<sub>b</sub>.
NaOH (s) (sodium hydroxide) is a stronger base than (CH<sub>3</sub>CH<sub>2</sub>)<sub>2</sub>NH (l) (diethylamine) which is a stronger base than NH<sub>3</sub> (g) (ammonia). As the bases get weaker, the smaller the K<sub>b</sub> values become. The pie-chart representation is as follows:
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 represents the base.
In this formula, [B]<sub>initial</sub> is the initial molar concentration of the base, assuming that no protonation has occurred.
Calculate the pH and percentage protonation of a .20 M aqueous solution of pyridine, C<sub>5</sub>H<sub>5</sub>N. The K<sub>b</sub> for C<sub>5</sub>H<sub>5</sub>N is 1.8 x 10<sup>-9</sup>.
First, write the proton transfer equilibrium:
The equilibrium table, with all concentrations in moles per liter, is
This means .0095% of the pyridine is in the protonated form of C<sub>5</sub>H<sub>6</sub>N<sup>+</sup>.
Other weak bases are essentially any bases not on the list of strong bases.