The weak bases they are species with little tendency to donate electrons, dissociate in aqueous solutions, or accept protons. The prism with which its characteristics are analyzed is governed by the definition emerged by the studies of several famous scientists.
For example, according to the Bronsted-Lowry definition, a weak base is one that accepts in a very reversible (or null) a hydrogen ion H+. In water, its H moleculetwoO is the one who donates an H+ to the surrounding base. If instead of water it were a weak acid HA, then the weak base could hardly neutralize it.
A strong base would not only neutralize all acids in the environment, but could also participate in other chemical reactions with adverse (and deadly) consequences..
It is for this reason that some weak bases, such as milk magnesia, or phosphate salts or sodium bicarbonate tablets, are used as antacids (top image).
All weak bases have in common the presence of an electron pair or a stabilized negative charge on the molecule or ion. Thus, the CO3- is a weak base against OH-; and the base that produces less OH- in its dissociation (Arrenhius definition) it will be the weakest base.
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A weak base can be written as BOH or B. It is said to undergo dissociation when the following reactions occur with both bases in liquid phase (although it can occur in gases or even solids):
BOH <=> B+ + Oh-
B + HtwoOR <=> HB+ + Oh-
Note that although both reactions may seem different, they have in common the production of OH-. Furthermore, the two dissociations establish an equilibrium, so they are incomplete; that is, only a percentage of the base actually dissociates (which does not happen with strong bases such as NaOH or KOH).
The first reaction "sticks" more closely to the Arrenhius definition for bases: dissociation in water to give ionic species, especially the hydroxyl anion OH-.
While the second reaction, obeys the Bronsted-Lowry definition, since B is being protonated or accepts H+ of the water.
However, the two reactions, when they establish an equilibrium, are considered weak base dissociations.
Ammonia is perhaps the most common weak base of all. Its dissociation in water can be outlined as follows:
NH3 (ac) + HtwoO (l) <=> NH4+ (ac) + OH- (ac)
Therefore, the NH3 falls into the category of bases represented with 'B'.
The dissociation constant of ammonia, Kb, is given by the following expression:
Kb = [NH4+] [OH-] / [NH3]
Which at 25 ° C in water is about 1.8 x 10-5. Then calculating its pKb you have:
pKb = - log Kb
= 4.74
In the dissociation of NH3 This receives a proton from water, so water can be considered an acid according to Bronsted-Lowry.
The salt formed on the right hand side of the equation is ammonium hydroxide, NH4OH, which is dissolved in water and is nothing more than aqueous ammonia. It is for this reason that the Arrenhius definition for a base is fulfilled with ammonia: its dissolution in water produces NH ions4+ and OH-.
NH3 is capable of donating a pair of unshared electrons located on the nitrogen atom; This is where the Lewis definition for a base comes in, [H3N:].
The concentration of the aqueous solution of the weak base methylamine (CH3NHtwo) is as follows: [CH3NHtwo] before dissociation = 0.010 M; [CH3NHtwo] after dissociation = 0.008 M.
Calculate Kb, pKb, pH and percentage of ionization.
First the equation of its dissociation in water must be written:
CH3NHtwo (ac) + HtwoO (l) <=> CH3NH3+ (ac) + OH- (ac)
Following the mathematical expression of Kb
Kb = [CH3NH3+] [OH-] / [CH3NHtwo]
In equilibrium it is satisfied that [CH3NH3+] = [OH-]. These ions come from the dissociation of CH3NHtwo, so the concentration of these ions is given by the difference between the concentration of CH3NHtwo before and after dissociating.
[CH3NHtwo]dissociated = [CH3NHtwo]initial - [CH3NHtwo]Balance
[CH3NHtwo]dissociated = 0.01 M - 0.008 M
= 0.002 M
So, [CH3NH3+] = [OH-] = 2 ∙ 10-3 M
Kb = (2 ∙ 10-3)two M / (8 ∙ 10-two) M
= 5 ∙ 10-4
Calculated Kb, it is very easy to determine pKb
pKb = - log Kb
pKb = - log 5 ∙ 10-4
= 3,301
To calculate the pH, since it is an aqueous solution, the pOH must first be calculated and subtracted from 14:
pH = 14 - pOH
pOH = - log [OH-]
And since the concentration of OH is already known-, the calculation is direct
pOH = -log 2 ∙ 10-3
= 2.70
pH = 14 - 2.7
= 11.3
To calculate it, it must be determined how much of the base has been dissociated. As this was already done in the previous points, the following equation applies:
([CH3NH3+] / [CH3NHtwo]°) x 100%
Where [CH3NHtwo]° is the initial concentration of the base, and [CH3NH3+] the concentration of its conjugated acid. Calculating then:
Percentage of ionization = (2 ∙ 10-3 / 1 ∙ 10-two) x 100%
= 20%
-The weak amine bases have a characteristic bitter taste, present in fish and which is neutralized with the use of lemon..
-They have a low dissociation constant, which is why they cause a low ion concentration in aqueous solution. Not being, for this reason, good conductors of electricity.
-In aqueous solution they originate a moderate alkaline pH, which is why they change the color of litmus paper from red to blue.
-They are mostly amines (weak organic bases).
-Some are the conjugate bases of strong acids.
-Weak molecular bases contain structures capable of reacting with H+.
-Methylamine, CH3NHtwo, Kb = 5.0 ∙ 10-4, pKb = 3.30
-Dimethylamine, (CH3)twoNH, Kb = 7.4 ∙ 10-4, pKb = 3.13
-Trimethylamine, (CH3)3N, Kb = 7.4 ∙ 10-5, pKb = 4.13
-Pyridine, C5H5N, Kb = 1.5 ∙ 10-9, pKb = 8.82
-Aniline, C6H5NHtwo, Kb = 4.2 ∙ 10-10, pKb = 9.32.
The nitrogenous bases adenine, guanine, thymine, cytosine and uracil are weak bases with amino groups, which are part of the nucleotides of nucleic acids (DNA and RNA), where the information for hereditary transmission resides.
Adenine, for example, is part of molecules such as ATP, the main energy reservoir of living beings. Furthermore, adenine is present in coenzymes such as flavin adenyl dinucleotide (FAD) and nicotin adenyl dinucleotide (NAD), which are involved in numerous oxidation-reduction reactions.
The following weak bases, or that can fulfill a function as such, are ordered in decreasing order of basicity: NHtwo > OH- > NH3 > CN- > CH3COO- > F- > NO3- > Cl- > Br- > I- > ClO4-.
The location of the conjugate bases of the hydracids in the given sequence indicates that the greater the strength of the acid, the lower the strength of its conjugate base..
For example, the anion I- is an extremely weak base, while NHtwo is the strongest of the series.
On the other hand, finally, the basicity of some common organic bases can be arranged in the following way: alkoxide> aliphatic amines ≈ phenoxides> carboxylates = aromatic amines ≈ heterocyclic amines.
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