Corrosive Substances

Additional scientific description

Both acids and bases can cause major public harm. They are both principally corrosive because they catalyse hydrolysis reactions and other nucleophilic substitution reactions that result in decomposition of biological molecules. (Brown, 2018; Burrows et al., 2021; McMurry, 2004; Wade and Simek, 2022) Acids are chemical species that are capable of donating a proton (Bronsted Acid) or capable of forming a covalent bond with an electron pair (Lewis Acid). (IUPAC, 2014) The dissociation process of strong acids almost goes to completion, meaning almost all of the acid molecules break apart into ions, producing high concentrations of hydronium ions that can initiate undesirable chemical reactions. This characteristic is in contrast to weak acids, which only partially dissociate in water. (Wade and Simek, 2022) Examples of strong acids include hydrochloric acid, nitric acid (ICSC 0183, 2016), perchloric acid (ICSC 1006, 2000) and sulfuric acid (ICSC 0362, 2016).

Bases are chemical species that have an available pair of electrons capable of forming a covalent bond with a proton (Brønsted base) or with the vacant orbital of some other species (Lewis base). (IUPAC, 2014) The dissociation process of strong bases goes to completion, producing high concentrations of hydroxide ions. They have a high affinity for accepting protons which fully neutralises them. The strength of the base is measured by its ability to donate hydroxide ions in solution. Lewis theory defines a strong base as an electron pair donor, making them effective at forming bonds with electron-deficient species. (Atkins & Jones, 2010; Burrows et al.,2021; McMurry et al., 2004; Wade and Simek, 2022) Examples of strong bases include caesium hydroxide (ICSC 1592, 2006), calcium hydroxide (ICSC 0408, 1997), lithium hydroxide (ICSC 0913, 2009), potassium hydroxide (ICSC 0357, 2010) and sodium hydroxide (ICSC 0360, 2010).

Lewis acids are electron-pair acceptors that are able to react with a Lewis acid base to form a Lewis adduct, by sharing the electron pair donated by the Lewis base. (IUPAC, 2014)

𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 𝐴𝐴𝐴𝐴𝐿𝐿𝐴𝐴 + 𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 𝐵𝐵𝐵𝐵𝐿𝐿𝐿𝐿 → 𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿𝐿 𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴𝐴

Acids and bases can form conjugate acid-base pairs, where the acid formed on protonation of a base is called the conjugate acid, and the base is the conjugate base. The conjugate acid always carries one unit of positive charge more than the base, but the absolute charges of the species are immaterial to the definition. (IUPAC, 2014h) Strong acid-base reactions are typically exothermic - the new bond formed between the proton and the base is stronger than the bond that was broken to release the proton, therefore the released energy raises the temperature of the surroundings. The strength of an acid is inversely proportional to the strength of its conjugate base - i.e. the conjugate base of a strong acid must be a weak base; and the conjugate base of a weak acid must be a strong base. In the reaction of an acid with a base, the equilibrium generally favours the weather acid and base. (Wade and Simek, 2022)

Metrics and numeric limits

Metrics are used to quantify the strength of acids, including their pH, pOH, pK_a, and pK_b; acidity/ionisation constants (K_a, K_b and K_w); degree of dissociations; and concentrations. pH (equation 1) is a measure of how acidic or alkaline a substance is, ranging (for aqueous solutions) from 0 (strongly acidic) to 14 (strongly basic).

Equation 1: 𝑝𝑝𝑝𝑝 + 𝑝𝑝𝑝𝑝𝑝𝑝 = 14

Strong acids have high concentrations of hydronium ions, typically giving low pH values close to zero when fully dissociated in water. The pK_a (equation 2) is a measure of the concentration of hydrogen ions when there are equal concentrations of both the ionised and unionised form of a substance. Strong acids also tend to have pK_a values below zero due to their equilibrium lying to the far right, resulting in negligible concentrations of undissociated acid. Whereas strong bases have high concentrations of hydroxide ions, typically giving high pH values close to 14 when fully dissociated in water. The pOH is a measure of the concentration of hydroxide ions in a solution. Strong bases also tend to have pOH values close to zero.

Equation 2: 𝑝𝑝𝐾𝐾_𝑎𝑎+ 𝑝𝑝𝐾𝐾_𝑏𝑏 = 𝑝𝑝𝐾𝐾_𝑤𝑤

The acidity constant (K_a equation 3) is another equilibrium constant for splitting off a proton from a charged or uncharged acid. This metric is usually very high for a strong acid, indicating complete dissociation. (IUPAC, 2014) Due to this dissociation, the concentrations are also typically very high.

Equation 3: 𝐾𝐾_𝑎𝑎𝐾𝐾_𝑏𝑏 = 𝐾𝐾_𝑤𝑤

The ionisation constant is another equilibrium constant for quantifying the extent to which a base ionises in solution. This metric is usually very large for a strong base, indicating complete ionisation. The degree of dissociation is how much the strong base completely dissociates into ions when dissolved in water, producing hydroxide ions. Acids can be neutralised by strong bases, producing water and salt during the vigorous reaction. This is due to the high concentration of hydroxide ions, giving an overall high concentration for a strong base.