Acids and Bases

Three definitions of Acids and Bases

In general:

Acids are chemicals that increase H+ ion concentration in water.

ex. HCl(s)+ H2O(ℓ) → H+(aq) + Cl−(aq)

Bases are chemicals that increase OH- ion concentration in water.

ex. NaOH(s) + H2O(ℓ) → Na+(aq)+ OH−(aq)

Arrehnius Definition

Svante Arrhenius, a Swedish chemist, proposed this definition of acids and bases:

Acids can generally be recognized as ionic compounds with H in the chemical formula because they ionize to produce H+ ions:

Other examples:

H2SO4 → 2H+ + SO42− (diprotic, producing 2 hydrogen ions)
HNO2 → H+ + NO2−

Bases can generally be recognized by ending with OH in the chemical formula because they ionize to produce OH− ions:

Other examples:

LiOH → Li+ + OH−
Ca(OH)2 → Ca2+ + 2OH−

Brø nsted-Lowry

Not all acids and bases follow Arrhenius's definition.

Brønsted-Lowry defines acids as hydrogen donors and bases as hydrogen acceptors.

Ammonia is basic, but does not contain a OH− ion. How does it change the concentration of OH−? It reacts with a H+ in water to leave an OH− remaining.

Conjugate acids-base pairs

Acid-base reactions are reversible. Hence the double arrows in the reaction above. When the base accepts the proton (H+) it forms an acid, called a conjugate acid.

The remaining ion leftover from the acid is called the conjugate base.

Acids are always paired with its conjugate base (the base that it produces).

Bases are always paired with its conjugate acid (the acid it produces).

Strange case of water

Water can act as either an acid or a base. Sometimes it accepts H+ and acts as a base or sometimes it donates H+ and acts as an acid. Therefore water is said to be amphoteric.

Water can even act as its own acid and base, which is called the autoionization of water.

Lewis Acids and Bases

Lewis acids accept electron pairs.

Lewis bases donate electron pairs.

This is the most general and flexible definition of acids and bases.

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