What is a conjugate base

A conjugate base is the chemical species that forms after a molecule of an acid is neutralized, either by the loss of a proton or by receiving an unpaired pair of electrons from a Lewis base . In other words, it is the product of an acid-base neutralization reaction that comes directly from the original acid. The acid and its conjugate base are collectively called a conjugate acid-base pair.

Consider the following Brønsted-Lowry dissociation reaction of a weak acid:

In this case, the acid is the reactant on the left, HA, while on the right are the proton released by the acid and the anion, A – , ^left over after the acid lost its proton.

The reason it is called a conjugate “base” is because all acid-base reactions are reversible, even those involving strong acids and bases (only their equilibrium constants are very large and the equilibria are far shifted toward the products). ). For this reason, what in one sense represents the ionization of an acid as in the previous equation, in the opposite sense represents the protonation of a base, in this case, the anion A ^– .

How to recognize a conjugate base

From the point of view of the Brønsted-Lowry concept of acids and bases, an acid is any substance that, when dissolved in water, is capable of ionizing and donating a proton. Since it is converted to its conjugate base in doing so, the only difference between an acid and its conjugate base is the absence of a proton.

In addition to this, because the proton is positive and takes its carb with it, the conjugate base always ends up with a lower electrical charge by one unit than the respective acid. This means that if the acid was neutral, then its conjugate base will be negative (charged -1), while if the acid is positive, then the conjugate base will be neutral, and so on.

Conjugate bases of polyprotic acids

Recognizing the conjugate base of a monoprotic acid is usually straightforward, however, in the cases of polyprotic acids, some confusion can arise. This is because we sometimes write dissociation reactions of acids like H ₂ SO ₄ as losing both protons in a single step. However, this is not what actually happens.

All polyprototic acids undergo successive ionization reactions, and in each reaction they are converted to a different conjugate base. The confusion arises from the fact that the first conjugate base of a polyprotic acid still retains protons, so in addition to conjugate bases, they are also acids that have their own conjugate base.

The following example will illustrate this more clearly:

Example of polyprotic acids and their conjugate bases: phosphoric acid

Perhaps one of the best examples to illustrate the equilibria of a polyprotic acid is phosphoric acid or H ₃ PO ₄ . This acid can lose a total of three protons according to the following reversible dissociation reactions:

In this case, phosphoric acid (H ₃ PO ₄ ) becomes the dihydrogen phosphate ion (H ₂ PO ₄ ^– ) after losing a proton, so this is its conjugate base. At the same time, H ₂ PO ₄ ^– is an acid that ionizes in the second reaction to become the hydrogen phosphate ion (HPO ₄ ^2- ), so the latter is the conjugate base of H ₂ PO ₄ ^– , but not from H ₃ PO ₄ . The same is true of the HPO ₄ ^2- ion , which is also an acid (in addition to being the conjugate base of H₂ OP ₄ ^– ). Upon dissociation, it becomes the phosphate ion, which is its conjugate base.

Relationship of the conjugate base to the acidity of the acid

The conjugate base structure can give clues about the acidity of any acid. Analyzing the stability of that chemical species and comparing it to the structural stability of the original acid helps explain why some acids are stronger than others.

Among the stability criteria that can be applied to the analysis of the structure of both the acid and its conjugate base are:

• Full octets: Lewis bond theory indicates that molecules with atoms that violate the octet rule are less stable than those in which all atoms have full octets.
• Resonance Structures: Molecules with more resonance structures are more stable than those with fewer.
• Aromaticity: Species that exhibit aromaticity tend to be much more stable than those that are not aromatic, and these are more stable than those that are antiaromatic.
• Lower Total Charge: In general, neutral species tend to be more stable than ionic species, and when comparing ions, those with less net charge tend to be more stable than those with more.
• Separation of charges: when comparing two structures with the same net charge, the one with fewer formal charges separated between several atoms is more stable than those with more formal charges.
• Location of formal charges: between two molecules that have the same formal charges, the one with the negative charges on the more electronegative atoms and the positive ones on the less electronegative atoms will be more stable.

Comparing the acid to its conjugate base based on these stability criteria allows you to determine whether the acid will prefer to be in its protonated (like HA, for example) or ionized (like A – , for example) ^form .

If the conjugate base is more stable than the acid, then the acid will tend to dissociate and be stronger, whereas if the opposite is true, it will be a weak acid.

Examples of acid:conjugate base pairs

Here are some additional examples of different acids and their respective conjugate bases:

• Hydrochloric acid and chloride anion (HCl and Cl ^– )
• The bicarbonate anion and the carbonate anion (HCO ₃ ^– and CO ₃ ^2- )
• The ammonium cation and ammonia (NH ₄ ⁺ and NH ₃ )
• Sulfuric acid and bisulfate (H ₂ SO ₄ and HSO ₄ ^– )

References

• assolea. (2020, May 2). 7.6: The pH Scale . Retrieved from https://assolea.org/es/7-6-la-escala-de-ph/
• Brown, T. (2021). Chemistry: The Central Science (11th ed.). London, England: Pearson Education.
• Chang, R., Manzo, Á. R., Lopez, PS, & Herranz, ZR (2020). Chemistry (10th ed.). New York City, NY: MCGRAW-HILL.
• Flowers, P., Theopold, K., Langley, R., & Robinson, W. (2019a). 14.1 Brønsted-Lowry Acids and Bases – Chemistry 2e . Retrieved from https://openstax.org/books/chemistry-2e/pages/14-1-bronsted-lowry-acids-and-bases
• pH and pOH . (2020, October 30). Retrieved from https://espanol.libretexts.org/@go/page/1911