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Weak base

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A weak base is a base that, upon dissolution in water, does not dissociate completely, so that the resulting aqueous solution contains only a small proportion of hydroxide ions and the concerned basic radical, and a large proportion of undissociated molecules of the base.

pH, Kb, and Kw

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Bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e., the solution is said to have a pH greater than 7.0 at standard conditions, potentially as high as 14 (and even greater than 14 for some bases). The formula for pH is:

Bases are proton acceptors; a base will receive a hydrogen ion from water, H2O, and the remaining H+ concentration in the solution determines pH. A weak base will have a higher H+ concentration than a stronger base because it is less completely protonated than a stronger base and, therefore, more hydrogen ions remain in its solution. Given its greater H+ concentration, the formula yields a lower pH value for the weak base. However, pH of bases is usually calculated in terms of the OH concentration. This is done because the H+ concentration is not a part of the reaction, whereas the OH concentration is. The pOH is defined as:

If we multiply the equilibrium constants of a conjugate acid (such as NH4+) and a conjugate base (such as NH3) we obtain:

As is just the self-ionization constant of water, we have

Taking the logarithm of both sides of the equation yields:

Finally, multiplying both sides by -1, we obtain:

With pOH obtained from the pOH formula given above, the pH of the base can then be calculated from , where pKw = 14.00.

A weak base persists in chemical equilibrium in much the same way as a weak acid does, 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:

A base that has a large Kb will ionize more completely and is thus a stronger base. As shown above, the pH of the solution, which depends on the H+ concentration, increases with increasing OH concentration; a greater OH concentration means a smaller H+ concentration, therefore a greater pH. Strong bases have smaller H+ concentrations because they are more fully protonated, leaving fewer hydrogen ions in the solution. A smaller H+ concentration means a greater OH concentration and, therefore, a greater Kb and a greater pH.

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.[1]

Percentage protonated

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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

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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.[3]

First, write the proton transfer equilibrium:

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
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+.

Examples

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Simple Facts

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  • 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

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References

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  1. ^ "Explanation of strong and weak bases]". ChemGuide. Retrieved 2018-03-23.
  2. ^ Howard Maskill (1985). The physical basis of organic chemistry. Oxford University Press, Incorporated. ISBN 978-0-19-855192-8.
  3. ^ "Calculations of weak bases". Mr Kent's Chemistry Page. Retrieved 2018-03-23.
  4. ^ Atkins, Peter, and Loretta Jones. Chemical Principles: The Quest for Insight, 3rd Ed., New York: W.H. Freeman, 2005.
  5. ^ Clark, Jim. "Strong and Weak Bases."N.p.,2002. Web.
  6. ^ a b c d 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.
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