# Weak base

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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: :$mbox\left\{pH\right\} = -log_\left\{10\right\} left \left[ mbox\left\{H\right\}^+ ight\right]$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.:$mbox\left\{pOH\right\} = -log_\left\{10\right\} left \left[ mbox\left\{OH\right\}^- ight\right]$

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

:$K_a imes K_b = \left\{ \left[H_3O^+\right] \left[NH_3\right] over \left[NH_4^+\right] \right\} imes \left\{ \left[NH_4^+\right] \left[OH^-\right] over \left[NH_3\right] \right\} = \left[H_3O^+\right] \left[OH^-\right]$

Since $\left\{K_w\right\} = \left[H_3O^+\right] \left[OH^-\right]$ then, "$K_a imes K_b = K_w$"

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

:$logK_a + logK_b = logK_w$

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

:$pK_a + pK_b = pK_w = 14.00$

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:

:$mathrm\left\{K_b=\left\{ \left[NH_4^+\right] \left[OH^-\right] over \left[NH_3\right]$

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:

:$B\left(aq\right) + H_2O\left(l\right) leftrightarrow HB^+\left(aq\right) + OH^-\left(aq\right)$

B represents the base.

:$Percentage protonated = \left\{molarity of HB^+ over initial molarity of B\right\} imes 100% = \left\{ \left[\left\{HB\right\}^+\right] over \left[B\right] _\left\{initial \left\{ imes 100%\right\}$

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:

:$mathrm\left\{H_2O\left(l\right) + C_5H_5N\left(aq\right) leftrightarrow C_5H_5NH^+ \left(aq\right) + OH^- \left(aq\right)\right\}$

:$K_b=mathrm\left\{ \left[C_5H_5NH^+\right] \left[OH^-\right] over \left[C_5H_5N\right] \right\}$

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

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

Examples

*Alanine, C3H5O2NH2
*Ammonia, NH3
*Methylamine, CH3NH2
*Pyridine, C5H5N

Other weak bases are essentially any bases not on the list of strong bases.

ee also

* Strong base
* Weak acid

References

*Atkins, Peter, and Loretta Jones. Chemical Principles: The Quest for Insight, 3rd Ed., New York: W.H. Freeman, 2005.

*http://wine1.sb.fsu.edu/chm1046/notes/AcidBase/WeakBase/WeakBase.htm
*http://www.chemguide.co.uk/physical/acidbaseeqia/bases.html
*http://bouman.chem.georgetown.edu/S02/lect16/lect16.htm
*http://www.intute.ac.uk/sciences/reference/plambeck/chem1/p01154.htm

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