Strong and weak acids are key concepts in chemistry. Strong acids completely dissociate into their ions in water, while weak acids incompletely dissociate. There are only a few strong acids, but many weak acids.
Strong acids completely dissociate in water into their ions and produce one of more protons or hydrogen cations per molecules. Inorganic or mineral acids tend to be strong acids. There are only 7 common strong acids. Here are their names and formulas:
- HCl – hydrochloric acid
- HNO3 – nitric acid
- H2SO4 – sulfuric acid (note: HSO4– is a weak acid)
- HBr – hydrobromic acid
- HI – hydroiodic acid
- HClO4 – perchloric acid
- HClO3 – chloric acid
Strong Acid Dissociation
A strong acid in water ionizes completely, so when the dissociation reaction is written as a chemical reaction, the reaction arrow points right:
- HCl → H+(aq) + Cl–(aq)
- HNO3 → H+(aq) + NO3(aq)–
- H2SO4 → 2H+(aq) + SO42-(aq)
While there are only a few strong acids, there are many weak acids. Weak acids incompletely dissociate in water to yield an equilibrium state that contains the weak acid and its ions. As an example, hydrofluoric acid (HF) is considered a weak acid because some HF remains in an aqueous solution, in addition to H+ and F– ions. Here is a partial list of common weak acids, ordered from strongest to weakest:
- HO2C2O2H – oxalic acid
- H2SO3 – sulfurous acid
- HSO4 – – hydrogen sulfate ion
- H3PO4 – phosphoric acid
- HNO2 – nitrous acid
- HF – hydrofluoric acid
- HCO2H – methanoic acid
- C6H5COOH – benzoic acid
- CH3COOH – acetic acid
- HCOOH – formic acid
Weak Acid Dissociation
Weak acids incompletely dissociate, forming an equilibrium state containing the weak acid and its ions. So, the reaction arrow points both ways. An example is the dissociation of ethanoic acid, which forms the hydronium cation and ethanoate anion:
CH3COOH + H2O ⇆ H3O+ + CH3COO–
Acid Strength (Strong vs. Weak Acids)
Acid strength is a measure of how readily the acid loses a proton or hydrogen cation. One mole of a strong acid HA dissociates in water to yield one mole of H+ and one mole of the acid’s conjugate base A−. In contrast, one mole of a weak acid yields less than one mole each of hydrogen cation and conjugate base, while some of the original acid remains. The two factors that determine how easily deprotonation occurs are the size of the atom and the polarity of the H-A bond.
In general, you can identify strong and weak acids based on the equilibrium constant Ka or pKa:
- Strong acids have high Ka values.
- Strong acids have low pKa values.
- Weak acids have small Ka values.
- Weak acids have large pKa values.
Concentrated vs. Dilute
The terms strong and weak are not the same as concentrated and dilute. A concentrated acid contains very little water. A dilute acid contains a large percentage of water. A dilute solution of sulfuric acid is still a strong acid solution and can cause a chemical burn. On the other hand, 12 M acetic acid is a concentrated weak acid (and still dangerous). If you dilute acetic acid enough, you get the concentration found in vinegar, which is safe to drink.
Strong vs. Corrosive
Most acids are highly corrosive. They can oxidize other substances and produce chemical burns. However, the strength of an acid is not a predictor of its corrosivity! The carborane superacids are not corrosive and can be safely handled. Meanwhile, hydrofluoric acid (a weak acid) is so corrosive it passes through skin and attacks bones.
Types of Acids
The three major acid classifications are Brønsted–Lowry acids, Arrhenius acids, and Lewis acids:
- Brønsted–Lowry acids: Brønsted–Lowry acids donate protons. In aqueous solution, the proton donor forms the hydronium cation (H3O+). However, Brønsted–Lowry acid-base theory also allows for acids in solvents besides water.
- Arrhenius acids: Arrhenius acids are hydrogen donors. Arrhenius acids dissociate in water and donate a hydrogen cation (H+) to form the hydronium cation (H3O+) . These acids are also characterized by turning litmus red, having a sour taste, and reacting with metals and bases to form salts.
- Lewis acids: Lewis acids are electron pair acceptors. Under this definition of an acid, the species either immediately accepts electron pairs or else it donates a hydrogen cation or proton and then accepts an electron pair. Technically, a Lewis acid must form a covalent bond with an electron pair. By this definition, Lewis acids often are not Arrhenius acids or Brønsted–Lowry acids. For example, HCl is not a Lewis acid.
All three acid definitions have their place in predicting chemical reactions and explaining behavior. Common acids are Brønsted–Lowry or Arrhenius acids. Lewis acids (e.g., BF3) are specifically identified as “Lewis acids.”
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- Petrucci R.H., Harwood, R.S.; Herring, F.G. (2002). General Chemistry (8th ed.) Prentice-Hall. ISBN 0-13-014329-4.