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When it comes to substances that are classified as acids or bases, there are three major categories to consider. According to the Arrhenius definition, an acid produces H+ in solution and a base produces OH- in solution. Svante Arrhenius developed this theory in 1883, and it is still in use today. Theoretical proposals for two more sophisticated and general theories were made later. The Bronsted-Lowry and Lewis definitions of acids and bases, respectively, are presented here.
The hierarchy of acid-base theories is depicted in this diagram. A subclass of Bronsted acids and bases, which are themselves a subclass of Lewis acids and bases, is represented by the Arrhenius acid and base family.
Chemical reactions involving acids and bases are described by the Arrhenius theory, which is the simplest and least general description available. Acids such as HClO4 and bases such as NaOH or Mg(OH)2 are included in this theory. When acids and bases react with one another, water and salts are formed. This theory successfully describes this reaction.
However, it does not explain why some substances, such as F- and NO2-, which do not contain hydroxide ions, can form basic solutions in water when combined with water. The Bronsted-Lowry definition of acids and bases provides a solution to this dilemma.
As in the Arrhenius theory, an acid is a substance that has the ability to release a proton, whereas a base is a substance that can accept a proton (similar to the Arrhenius theory). Na+F-, for example, is a basic salt that generates hydrogen ions in water by removing protons from the water itself (to form HF):
Acid Dissociation: HA(aq) ⇌A–(aq) + H+(aq)
Ka = [A–] [ H+ ]/HA
Base dissociation: B+(aq) + H2O(l) ⇌ HB+(aq) + OH– (aq)
Kb = [HB+][ OH- ]/[ B+]
HClO4(aq) H+(aq)+ClO4–(aq)
HBr(aq)⟶H+(aq)+Br−(aq)
CH3O−(aq)+H2O(l)⟶CH3OH(aq)+OH−(aq)
NH−2(aq)+H2O(l)⟶NH3(aq)+OH−(aq)
The right-handed arrow (⟶) indicates that the reaction will be carried out to its conclusion. As a result, when you dilute HClO4 in water to 1.0 M, you actually get 1.0 M H+(aq) and 1.0 M ClO–4aq), with very little remaining undissociated HClO4.
In contrast, weak acids and bases such as acetic acid (CH3COOH) and ammonia (NH3) dissociate only slightly in water – typically a few percent, depending on their concentration and the values of Ka and Kb – and exist primarily as the undissociated molecules.
One important consequence of these equilibria is that every acid (HA) has a conjugate base (A-), and vice versa. This is referred to as the acid-base relationship. In the base dissociation equilibrium, the conjugate acid of base B is represented by the symbol HB+.
Each of these equilibria is linked together by the water dissociation equilibrium for a given acid or base:
H2O(l) ⇌ H+(aq) + OH−(aq)
Kw = [H+][OH−]
for which the equilibrium constant Kw is 1.00 x 10-14 at 25°C, is a function of the temperature. It is simple to demonstrate that the product of the acid and base dissociation constants Ka and Kb is the constant Kw (for water).
The idea that strong acids have weak conjugate bases and weak acids have strong conjugate bases is a common misconception. Keeping in mind that KaKb = Kw , it is straightforward to see that this is incorrect. We define a strong acid or base as one in which K > 1, i.e., in which the substance completely dissociates, according to our definition. We define a weak acid or base as one that has a pH of 1 > K > Kw . As a result, if Ka > 1 (strong), then Kb cannot be greater than Kw (weak).
Strong acids, such as HCl, dissociate to produce spectator ions, such as Cl-, which act as conjugate bases, whereas weak acids, such as acetic acid, produce weak conjugate bases. In the case of acetic acid and its conjugate base, the acetate anion. Acetic acid is a weak acid (Ka = 1.8 x 10-5), and acetate is a weak base (Kb=KwKa=5.6 x 1010). Acetic acid is a weak acid and acetate is a weak base.
It is possible to produce acetic acid’s conjugate base and hydronium by reacting it with water in a reversible reaction.
The strength of a conjugate acid/base is inversely proportional to the strength or weakness of its parent acid or base. To be technical, any acid or base can be classified as either a conjugate acid or a conjugate base; these terms are simply used to distinguish between different species in solution (i.e acetic acid is the conjugate acid of the acetate anion, a base, while acetate is the conjugate base of acetic acid, an acid).
In accordance with the simple rule discovered by Linus Pauling, neutral oxyacids (H2SO4, H3PO4, HNO3, HClO2, and others) can be divided into two groups: strong and weak. In acids, if the number of oxygen atoms is two or more times greater than the number of hydrogen atoms, the acid is considered strong; otherwise, the acid is classified as weak. For example, the strong acids HClO4 and HClO3, where the difference is 3 and 2, respectively, are both formed by the reaction of HClO4. Neither HNO2 nor HClO2 are particularly strong because the difference between them is one in both cases. When it comes to weak acids, the relative strength is determined by the difference between the two (for example, HClO2 is a stronger weak acid than HOCl) and the electronegativity of the central atom (for example, HClO2 is a stronger weak acid than HOCl) (HOCl is stronger than HOI).
Acids that have the ability to donate more than one proton are referred to as polyprotic acids. Taking the acidity of sulfuric acid, H2SO4, as an example, it is a strong acid with a conjugate base that is actually a weak acid in and of itself. This means that for every mole of H2SO4 present in aqueous solution, more than one mole of protons is donated. Carbonic acid (H2CO3) and phosphoric acid (H3PO4) are both polyprotic acids with a low pH of 3.5. The sequential pKa’s of polyprotic acid are typically separated by about 5 pH units, because it becomes increasingly difficult to remove protons as the ion becomes more negatively charged as the pH of the solution increases. Examples include phosphoric acid, which has three different pH values: 2.15, 7.20, and 12.35.
Among the other types of amphoteric compounds are oxides and hydroxides of elements that are found on either side of the boundary between the metallic and nonmetallic elements in the periodic table. Al(OH)3 ions, for example, are insoluble at neutral pH, but they can accept protons in acid to form [Al(H2O)6]3+ ions or accept an OH- ion in base to form Al(OH)4- ions when exposed to acid. As a result, aluminium oxide is soluble in both acid and base solutions, but not in neutral water. Other amphoteric oxides include BeO, ZnO, Ga2O3, Sb2O3, and PbO, to name a few examples. The acidity of a metal’s oxide is increased by increasing the oxidation state of the metal, which is accomplished by removing electron density from the oxygen atoms. As a result, while Sb2O5 is acidic, Sb2O3 is amphoteric.
Chemical reactions involving acids and bases are described by the Arrhenius theory, which is the simplest and least general description available. Acids such as HClO4 and bases such as NaOH or Mg(OH)2 are included in this theory. When acids and bases react with one another, water and salts are formed. This theory successfully describes this reaction.
However, it does not explain why some substances, such as F- and NO2-, which do not contain hydroxide ions, can form basic solutions in water when combined with water. The Bronsted-Lowry definition of acids and bases provides a solution to this dilemma.
As in the Arrhenius theory, an acid is a substance that has the ability to release a proton, whereas a base is a substance that can accept a proton (similar to the Arrhenius theory). Na+F-, for example, is a basic salt that generates hydrogen ions in water by removing protons from the water itself (to form HF):
F−(aq) +H2O(l) ⇌ HF(aq) + OH−
In contrast, when a Bronsted acid dissociates, it causes an increase in the concentration of hydrogen ions in the solution, [H+]; in contrast, when a Bronsted base dissociates, it causes a proton to be taken from the solvent (water), resulting in [OH–].Acid Dissociation: HA(aq) ⇌A–(aq) + H+(aq)
Ka = [A–] [ H+ ]/HA
Base dissociation: B+(aq) + H2O(l) ⇌ HB+(aq) + OH– (aq)
Kb = [HB+][ OH- ]/[ B+]
Strong and weak acids and bases
Acids and bases that completely dissociate are referred to as strong acids and bases:HClO4(aq) H+(aq)+ClO4–(aq)
HBr(aq)⟶H+(aq)+Br−(aq)
CH3O−(aq)+H2O(l)⟶CH3OH(aq)+OH−(aq)
NH−2(aq)+H2O(l)⟶NH3(aq)+OH−(aq)
The right-handed arrow (⟶) indicates that the reaction will be carried out to its conclusion. As a result, when you dilute HClO4 in water to 1.0 M, you actually get 1.0 M H+(aq) and 1.0 M ClO–4aq), with very little remaining undissociated HClO4.
In contrast, weak acids and bases such as acetic acid (CH3COOH) and ammonia (NH3) dissociate only slightly in water – typically a few percent, depending on their concentration and the values of Ka and Kb – and exist primarily as the undissociated molecules.
Conjugate acids and bases
Conjugate acids and bases are acids and bases that differ by a proton (H+).One important consequence of these equilibria is that every acid (HA) has a conjugate base (A-), and vice versa. This is referred to as the acid-base relationship. In the base dissociation equilibrium, the conjugate acid of base B is represented by the symbol HB+.
Each of these equilibria is linked together by the water dissociation equilibrium for a given acid or base:
H2O(l) ⇌ H+(aq) + OH−(aq)
Kw = [H+][OH−]
for which the equilibrium constant Kw is 1.00 x 10-14 at 25°C, is a function of the temperature. It is simple to demonstrate that the product of the acid and base dissociation constants Ka and Kb is the constant Kw (for water).
The idea that strong acids have weak conjugate bases and weak acids have strong conjugate bases is a common misconception. Keeping in mind that KaKb = Kw , it is straightforward to see that this is incorrect. We define a strong acid or base as one in which K > 1, i.e., in which the substance completely dissociates, according to our definition. We define a weak acid or base as one that has a pH of 1 > K > Kw . As a result, if Ka > 1 (strong), then Kb cannot be greater than Kw (weak).
Strong acids, such as HCl, dissociate to produce spectator ions, such as Cl-, which act as conjugate bases, whereas weak acids, such as acetic acid, produce weak conjugate bases. In the case of acetic acid and its conjugate base, the acetate anion. Acetic acid is a weak acid (Ka = 1.8 x 10-5), and acetate is a weak base (Kb=KwKa=5.6 x 1010). Acetic acid is a weak acid and acetate is a weak base.
It is possible to produce acetic acid’s conjugate base and hydronium by reacting it with water in a reversible reaction.
The strength of a conjugate acid/base is inversely proportional to the strength or weakness of its parent acid or base. To be technical, any acid or base can be classified as either a conjugate acid or a conjugate base; these terms are simply used to distinguish between different species in solution (i.e acetic acid is the conjugate acid of the acetate anion, a base, while acetate is the conjugate base of acetic acid, an acid).
In accordance with the simple rule discovered by Linus Pauling, neutral oxyacids (H2SO4, H3PO4, HNO3, HClO2, and others) can be divided into two groups: strong and weak. In acids, if the number of oxygen atoms is two or more times greater than the number of hydrogen atoms, the acid is considered strong; otherwise, the acid is classified as weak. For example, the strong acids HClO4 and HClO3, where the difference is 3 and 2, respectively, are both formed by the reaction of HClO4. Neither HNO2 nor HClO2 are particularly strong because the difference between them is one in both cases. When it comes to weak acids, the relative strength is determined by the difference between the two (for example, HClO2 is a stronger weak acid than HOCl) and the electronegativity of the central atom (for example, HClO2 is a stronger weak acid than HOCl) (HOCl is stronger than HOI).
Acids that have the ability to donate more than one proton are referred to as polyprotic acids. Taking the acidity of sulfuric acid, H2SO4, as an example, it is a strong acid with a conjugate base that is actually a weak acid in and of itself. This means that for every mole of H2SO4 present in aqueous solution, more than one mole of protons is donated. Carbonic acid (H2CO3) and phosphoric acid (H3PO4) are both polyprotic acids with a low pH of 3.5. The sequential pKa’s of polyprotic acid are typically separated by about 5 pH units, because it becomes increasingly difficult to remove protons as the ion becomes more negatively charged as the pH of the solution increases. Examples include phosphoric acid, which has three different pH values: 2.15, 7.20, and 12.35.
Compounds with Amphoteric properties
Some substances have the ability to act both as an acid and as a base. Water is a good example. H2O molecules have the ability to either donate or accept hydrogen ions. Because of this, water has the property of being an amphoteric solvent. When an acid dissociates in solution, water acts as a base, and the acid dissociates in solution. When bases dissociate, water, on the other hand, acts as an acid. OH– (aq) is the most powerful acid that we can make in water, and H+ (aq) is the most powerful base we can make in water (aq).Among the other types of amphoteric compounds are oxides and hydroxides of elements that are found on either side of the boundary between the metallic and nonmetallic elements in the periodic table. Al(OH)3 ions, for example, are insoluble at neutral pH, but they can accept protons in acid to form [Al(H2O)6]3+ ions or accept an OH- ion in base to form Al(OH)4- ions when exposed to acid. As a result, aluminium oxide is soluble in both acid and base solutions, but not in neutral water. Other amphoteric oxides include BeO, ZnO, Ga2O3, Sb2O3, and PbO, to name a few examples. The acidity of a metal’s oxide is increased by increasing the oxidation state of the metal, which is accomplished by removing electron density from the oxygen atoms. As a result, while Sb2O5 is acidic, Sb2O3 is amphoteric.
