Acids and bases

The acids and bases are basic compounds for the chemical reactions . They allow to regulate the pH necessary for many reactions to occur or for the conservation of different compounds and microorganisms. If you are interested in chemistry, you cannot fail to know what the differences are between  acids and bases, which, as you will see later, are two sides of the same coin.

Acid
Base
Definition according to Arrhenius (*)An acid is any chemical compound that when dissolved in water releases H + .A base is any chemical compound that when dissolved in water produces OH-.
Definition according to Bronstead-Lowry (*)Acids are substances capable of giving up protons (H + )Bases are substances capable of accepting protons (H + )
Definition according to Lewis (*)A Lewis acid is defined as a substance capable of sharing or accepting an electron pair.A Lewis base is a substance with the ability to share or give pairs of electrons.
pH 0<pH<7 7<pH<14
PowerfulThey dissociate completely when dissolved in water.They associate completely with H + ions when they dissolve in water, generating OH- ions.
Strong: ExamplesHCl, H2SO4NaOH, LiOH
WeakThey generate an acid-base balance when in water, partially dissociating according to the acid dissociation constant (Ka) of each compound.They generate an acid-base balance when in water, partially dissociating according to the acid dissociation constant (Ka) of each compound.
Weak: ExamplesH2CO3NH3,

What are acids

They are those substances that are capable of giving up protons (according to Bronstead-Lowry)   or of accepting or sharing electrons (according to Lewis). In water, they produce a pH lower than 7, since they lower it from the original pH of pure water: 7. They can be classified as strong or weak acids. The strong acids , completely dissociate such as HCl:

HCL —–> H+ Cl-

The weak acids are partially dissociated according to the acid dissociation constant (Ka) which is characteristic of each acid. One of them is, for example, carbonic acid that dissociates according to:

H2CO3 <—> H+ + HCO3-

HCO3- <—> H+ – CO32-

Where each dissociation has a different Ka.

What are the bases

They are those substances that are capable of accepting protons (according to Bronstead-Lowry)   or of giving up or sharing electrons (according to Lewis). In water, they produce a pH greater than 7 and less than 14, since they increase it from the original pH of pure water: 7. They can be classified as strong or weak bases. The strong bases , completely dissociate such as NaOH:

NaOH —–> Na+ + OH-

The weak bases partially dissociated according to the acid dissociation constant (Ka) which is specific to each base. One of them is, for example, ammonia that accepts protons according to:

NH3 + H3O+ <—> NH4+ + H2O

Acids and Bases: Evolution of Definitions

The definition of Arrehinius is the oldest as it defines acids and bases based on their ability to give up protons (H +) and hydroxide ions (OH-) respectively. It was questioned because there are substances that produce an increase in the concentration of OH- without containing them in their chemical composition. For example, NH3:

NH3 + H20 <—> NH4+ + OH-

A more satisfactory theory than Arrhenius’s is that formulated in 1923 by Brønsted and Lowry . This theory differentiates acids from bases by their ability to give up or accept protons (H +). According to this theory, ammonia (NH3) clearly classifies as a base given its ability to accept protons. They, in turn, classified the acids as strong or weak according to whether they accepted or gave up their (strong) protons completely or did so partially.

Finally, the Lewis theory goes even further than the variation of the pH and takes into account whether the substance accepts or gives up a pair of electrons. Note that electrons (-) and protons (+) are oppositely charged particles and are the basis of all acid-base reactions. The H + is called a proton because it is a hydrogen that has lost its electron and therefore, is composed only of its positively charged part. Therefore, when a compound accepts an H + ion it is accepting protons and giving up electrons, thus functioning as a Lewis base.

In this way, all the substances that for the theories of Arrhenius or Bronsted-Lowry were acids or bases, they are also for Lewis. However, Lewis extends the concept of acid beyond the previous theories and many Lewis acids are not Bronsted, such as BF3 that shares its free electron pair with other compounds:

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