# Acidity

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The acidity of a solution is determined by the concentration of hydrogen (H+) or hydronium (H3O+) ions dissolved in an aqueous solution, in units of moles per Liter (aka Molarity). More commonly, the level of acidity of a solution is represented by the pH scale.

## pH Scale

The pH scale, usually reported in values on a range 0-14, provides a measure of the acidity or alkalinity of a solution. Solutions with a pH < 7 are considered acidic; solutions with a pH > 7 are considered basic; and solutions with a pH = 7 are considered neutral (neither acidic nor basic). The calculation of pH is given by:

${\displaystyle pH=-log[H^{+}]}$,      or more accurately as ${\displaystyle pH=-log[H_{3}O^{+}]}$     Eq. 1

How is the pH scale range defined? In other words, why are pH values usually between 0-14 and not some other range of values? As seen from Eq. 1 above, the pH scale is on a logarithmic scale, meaning that for every 1 pH unit change, the molar concentration of hydrogen ions changes by a factor of 10. For example, a solution with pH = 4 means that [H+] = 1x10-4 mol/L, while a solution with pH = 5 means that [H+] = 1x10-5 mol/L.

## Autodissociation of Water and the pH Scale

When two water molecules (or even two moles of water) come together, then can undergo a process called autodissociation. This is an equilibrium denoted by the following:

${\displaystyle H_{2}O(l)+H_{2}O(l)\rightleftharpoons H_{3}O^{+}(aq)+OH^{-}(aq)}$      Eq. 2

Here, two water molecules react to transfer a proton (H+) from one water molecule to the other. In essence, one water molecule acts as an acid and donates a proton and the other molecule acts as a base and accepts a proton. The resulting products are the aqueous forms of acid (H3O+) and base (OH-).
The degree to which this reaction goes from reactants (left side) to products (right side) is guided by an equilibrium product or equilibrium constant, denoted as K. Mathematically, this is defined as:

${\displaystyle K={\frac {products}{reactants}}}$,       or specifically ${\displaystyle K=[H_{3}O^{+}][OH^{-}]}$      Eq. 3

In calculating an equilibrium constant, K, only aqueous species and gas pressures are considered. Pure liquids and solids play no role in the calculation of K values.
This special case of equilibrium for the autodissociation of water is defined as Kw is has a value of Kw = 1.00x10-14. In pure water, where [H3O+] and [OH-] can only come from H2O, the concentrations of both [H3O+] and [OH-] can be solved by:

${\displaystyle K_{w}=[H_{3}O^{+}][OH^{-}]=1.00x10^{-14}}$     Eq. 4

Because [H3O+] and [OH-] are created in equal amounts from H2O, we can say that:

${\displaystyle 1.00x10^{-14}=[x][x]=x^{2}}$     Eq. 5

and so the concentrations of both [H3O+] and [OH-], represented by "x", are equal:

${\displaystyle x=1.00x10^{-7}}$     Eq. 6

Therefore, in neutral water, ${\displaystyle pH=-log[1.00x10^{-7}]=7.00}$

## Acidic and Basic Solutions

In solutions containing anything other than pure water, the pH may be affected by dissolved species that upset the equilibrium shown above. An acidic species would be any compound that increased the [H3O+] concentration, and a basic species would be any compound that increased the [OH-] concentration (or alternately, decreased the [H3O+] concentration).

For the purposes of this project, acids will be considered mainly as proton donors. That is, when mixed in water an acid will dissociate into H+ and some anion. For example, hydrochloric acid dissociates fully in water (known as a strong acid) as:

${\displaystyle HCl(aq)\rightarrow H^{+}(aq)+Cl^{-}(aq)}$

Other common strong acids include nitric acid (HNO3) and sulfuric acid (H2SO4). Other acidic compounds, such as organic acids found in plants and animals, are considered weak acids because they do not fully dissociate. For example, acetic acid, the main ingredient in vinegar does not fully dissociate:

${\displaystyle CH_{3}COOH(aq)\rightleftharpoons CH_{3}COO^{-}(aq)+H^{+}(aq)}$

Strong bases are compounds that fully dissociate into [OH-] ions and some associated cation. For example, sodium hydroxide dissociates fully in water as:

${\displaystyle NaOH(s)\rightarrow Na^{+}(aq)+OH^{-}(aq)}$

Which increases the [OH-] concentration in water (or, decreases the relative [H3O+] concentration). In the case of weak bases, such as ammonia and organic bases found in plants and animals, they provide increased [OH-] concentration as shown below by reacting with water:

${\displaystyle NH_{3}(aq)+H_{2}O(l)\rightleftharpoons NH_{4}^{+}(aq)+OH^{-}(aq)}$

...coming soon

## Notes and References

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