Weak base
A weak base is a base that upon dissolution in water, does not dissociate completely and the resulting aqueous solution contains hydroxide ions and the concerned basic radical in a small proportion, along with a large proportion of undissociated molecules of the base.
Acids and bases |
---|
Acid types |
Base types |
pH, K_{b}, and K_{w}
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, H_{2}O, 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 NH_{4}^{+}) and a conjugate base (such as NH_{3}) the following is given:
Since then,
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 = pK_{w} - pOH where pK_{w} = 14.00.
Weak bases exist in chemical equilibrium much in the same way as weak acids do, with a base dissociation constant (K_{b}) 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_{b} 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 (K_{w} = 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 K_{b}.
NaOH (s) (sodium hydroxide) is a stronger base than (CH_{3}CH_{2})_{2}NH (l) (diethylamine) which is a stronger base than NH_{3} (g) (ammonia). As the bases get weaker, the smaller the K_{b} values become.[1]
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.[2]
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, C_{5}H_{5}N. The K_{b} for C_{5}H_{5}N is 1.8 x 10^{−9}.[3]
First, write the proton transfer equilibrium:
The equilibrium table, with all concentrations in moles per liter, is
C_{5}H_{5}N | C_{5}H_{6}N^{+} | 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 | |
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 = pK_{w} - 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 C_{5}H_{5}NH^{+}.
Examples
- Alanine
- Ammonia, NH_{3}
- Methylamine, CH_{3}NH_{2}
- Ammonium hydroxide, NH_{4}OH
Simple Facts
- An example of a weak base is ammonia. It does not contain hydroxide ions, but it reacts with water to produce ammonium ions and hydroxide ions.[4]
- The position of equilibrium varies from base to base when a weak base reacts with water. The further to the left it is, the weaker the base.[5]
- When there is a hydrogen ion gradient between two sides of the biological membrane, the concentration of some weak bases are focused on only one side of the membrane.[6] Weak bases tend to build up in acidic fluids.[6] Acid gastric contains a higher concentration of weak base than plasma.[6] Acid urine, compared to alkaline urine, excretes weak bases at a faster rate.[6]
See also
References
- "Explanation of strong and weak bases]". ChemGuide. Retrieved 2018-03-23.
- Howard Maskill (1985). The physical basis of organic chemistry. Oxford University Press, Incorporated. ISBN 978-0-19-855192-8.
- "Calculations of weak bases". Mr Kent's Chemistry Page. Retrieved 2018-03-23.
- Atkins, Peter, and Loretta Jones. Chemical Principles: The Quest for Insight, 3rd Ed., New York: W.H. Freeman, 2005.
- Clark, Jim. "Strong and Weak Bases."N.p.,2002. Web.
- Milne, M.D.; Scribner, B.H.; Crawford, M.A. (1958). "Non-ionic diffusion and the excretion of weak acids and bases". The American Journal of Medicine. 24 (5): 709–729. doi:10.1016/0002-9343(58)90376-0.