`color{purple}(✓✓)color{purple} " DEFINITION ALERT"`
According to Brönsted-Lowry theory, acid is a substance that is capable of donating a hydrogen ion `color{red}(H^+)` and bases are substances capable of accepting a hydrogen ion, `color{red}(H^+)`. In short, acids are proton donors and bases are proton acceptors.
Consider the example of dissolution of `color{red}(NH_3)` in `color{red}(H_2O)` represented by the following equation:
`color{red}(undersettext(base)(NH_3(aq)) +undersettext(acid)(H_2O(I)) ⇌ undersettext(conjugate acid)(NH_4^(+)(aq)) + undersettext(conjugate base)(OH^(-) (aq)))`
`color{red}(NH_3(aq) oversettext(adds proton)→ undersettext(conjugate acid)(NH_4^(+)(aq)))`
`color{red}(undersettext(acid)(H_2O(l)) oversettext(loses proton) → undersettext(conjugate base)(OH^(-)(aq)))`
•The basic solution is formed due to the presence of hydroxyl ions. In this reaction, water molecule acts as proton donor and ammonia molecule acts as proton acceptor and are thus, called Lowry-Brönsted acid and base, respectively.
•In the reverse reaction, `color{red}(H^+)` is transferred from `color{red}(NH_4^(+))` to `color{red}(OH^–)`. In this case, `color{red}(NH_4^+)` acts as a Bronsted acid while `color{red}(OH^–)` acted as a Brönsted base.
•The acid-base pair that differs only by one proton is called a conjugate acid-base pair. Therefore, `color{red}(OH^–)` is called the conjugate base of an acid `color{red}(H_2O)` and `color{red}(NH_4^+)` is called conjugate acid of the base `color{red}(NH_3)`.
•If Brönsted acid is a strong acid then its conjugate base is a weak base and viceversa. It may be noted that conjugate acid has one extra proton and each conjugate base has one less proton.
Consider the example of ionization of hydrochloric acid in water. `color{red}(HCl(aq))` acts as an acid by donating a proton to `color(H_2O)` molecule which acts as a base.
`color{red}(undersettext(acid)(HCl(aq)) + undersettext(base)(H_2O(l)) ⇌ undersettext(conjugate acid)(H_3O^(+)(aq)) + undersettext(conjugate base)(Cl^(-) (aq)))`
`color{red}(undersettext(acid)(HCl(aq)) oversettext(loses proton)→ undersettext(conjugate base)(Cl^(-)(aq)))`
`color{red}(undersettext(base)(H_2O(l)) oversettext(adds proton) → undersettext(conjugate acid)(H_3O^(+)(aq)))`
It can be seen in the above equation, that water acts as a base because it accepts the proton. The species `color{red}(H_3O^+)` is produced when water accepts a proton from `color{red}(HCl)`. Therefore, `color{red}(Cl^–)` is a conjugate base of `color{red}(HCl)` and `color{red}(HCl)` is the conjugate acid of base `color{red}(Cl^–)`. Similarly, `color{red}(H_2O)` is a conjugate base of an acid `color{red}(H_3O^+)` and `color{red}(H_3O^+)` is a conjugate acid of base `color{red}(H_2O).`
It is interesting to observe the dual role of water as an acid and a base. In case of reaction with `color{red}(HCl)` water acts as a base while in case of ammonia it acts as an acid by donating a proton.
`color{purple}(✓✓)color{purple} " DEFINITION ALERT"`
According to Brönsted-Lowry theory, acid is a substance that is capable of donating a hydrogen ion `color{red}(H^+)` and bases are substances capable of accepting a hydrogen ion, `color{red}(H^+)`. In short, acids are proton donors and bases are proton acceptors.
Consider the example of dissolution of `color{red}(NH_3)` in `color{red}(H_2O)` represented by the following equation:
`color{red}(undersettext(base)(NH_3(aq)) +undersettext(acid)(H_2O(I)) ⇌ undersettext(conjugate acid)(NH_4^(+)(aq)) + undersettext(conjugate base)(OH^(-) (aq)))`
`color{red}(NH_3(aq) oversettext(adds proton)→ undersettext(conjugate acid)(NH_4^(+)(aq)))`
`color{red}(undersettext(acid)(H_2O(l)) oversettext(loses proton) → undersettext(conjugate base)(OH^(-)(aq)))`
•The basic solution is formed due to the presence of hydroxyl ions. In this reaction, water molecule acts as proton donor and ammonia molecule acts as proton acceptor and are thus, called Lowry-Brönsted acid and base, respectively.
•In the reverse reaction, `color{red}(H^+)` is transferred from `color{red}(NH_4^(+))` to `color{red}(OH^–)`. In this case, `color{red}(NH_4^+)` acts as a Bronsted acid while `color{red}(OH^–)` acted as a Brönsted base.
•The acid-base pair that differs only by one proton is called a conjugate acid-base pair. Therefore, `color{red}(OH^–)` is called the conjugate base of an acid `color{red}(H_2O)` and `color{red}(NH_4^+)` is called conjugate acid of the base `color{red}(NH_3)`.
•If Brönsted acid is a strong acid then its conjugate base is a weak base and viceversa. It may be noted that conjugate acid has one extra proton and each conjugate base has one less proton.
Consider the example of ionization of hydrochloric acid in water. `color{red}(HCl(aq))` acts as an acid by donating a proton to `color(H_2O)` molecule which acts as a base.
`color{red}(undersettext(acid)(HCl(aq)) + undersettext(base)(H_2O(l)) ⇌ undersettext(conjugate acid)(H_3O^(+)(aq)) + undersettext(conjugate base)(Cl^(-) (aq)))`
`color{red}(undersettext(acid)(HCl(aq)) oversettext(loses proton)→ undersettext(conjugate base)(Cl^(-)(aq)))`
`color{red}(undersettext(base)(H_2O(l)) oversettext(adds proton) → undersettext(conjugate acid)(H_3O^(+)(aq)))`
It can be seen in the above equation, that water acts as a base because it accepts the proton. The species `color{red}(H_3O^+)` is produced when water accepts a proton from `color{red}(HCl)`. Therefore, `color{red}(Cl^–)` is a conjugate base of `color{red}(HCl)` and `color{red}(HCl)` is the conjugate acid of base `color{red}(Cl^–)`. Similarly, `color{red}(H_2O)` is a conjugate base of an acid `color{red}(H_3O^+)` and `color{red}(H_3O^+)` is a conjugate acid of base `color{red}(H_2O).`
It is interesting to observe the dual role of water as an acid and a base. In case of reaction with `color{red}(HCl)` water acts as a base while in case of ammonia it acts as an acid by donating a proton.