# Why isn't water acidic?

The definition of an acid is a compound with a hydrogen cation and a non-metal anion. Water is a hydroxide bonded to a hydrogen atom. This hydroxide has a net negative charge, too, since the negative oxygen's charge is stronger than the positive hydrogen's. Water has a pH of 7, though, meaning that it's neutral. Why isn't water acidic

If you look at what happens in the equilibrium dissociation of water, both a hydronium ion (Arrhenius acid) and a hydroxide ion (Arrhenius base) are formed. Since these two ions are formed in equal quantities ($\ce{10^{-7} mol/L}$ each in pure water), the water will be neither acidic nor basic.

$$\ce{H2O_{(l)} + H2O_{(l)} <=> H3O+_{(aq)} + OH^{-}_{(aq)}}$$ (from wikipedia)

Wikipedia can explain more about acids and their definitions if you are interested: http://en.wikipedia.org/wiki/Acid.

• In addition to this, pure water is defined as neutral, as it is a most common solvent in nature. – ssavec Dec 10 '14 at 6:59
• Just keep in mind also that water's pH is often a bit less than 7 because of the carbone dioxide in the air dissolving in the water forming carbonic acid. – Babounet Dec 10 '14 at 8:35
• Also the pH is not related to neutrality. Neutrality is when the number of hydroxide and hydrogen ions are equal. pH is just a measure of hydrogen ion concentration, and this is temperature dependent. At higher temperatures, water's pH decreases because it dissociates more but it remains neutral because each dissociation produces one hydrogen ion and one hydroxide. – bon Apr 30 '16 at 15:46
• pH is a measure of hydrogen ions, just as pOH is a measure of hydroxide ions. a pH or an pOH of 7 DOES mean that that the number of hydrogen and hydroxide ions are in balance - or "neutral" in that the solution is not acidic or basic. – Stephen Inoue Apr 16 '18 at 22:39

Water has “the neutral pH” because pH = 7 (a meme only useful in aqueous chemistry) is used to estimate acidity/basicity of a substance with respect to water namely.

In Brønsted–Lowry approach, water is an acid and a base simultaneously: H2O sometimes dissociates to OH (the conjugate base of H2O) and H+, i. e. we see deprotonation. The latter immediately reacts with surrounding H2O to form aqueous varieties of hydrogen cation (conjugate acid of H2O, simplistically called “hydronium”), i. e. we see protonation. But there is an equilibrium, namely OH and (H2O)nH+ react back to form the covalent water. This results in certain stable (under constant conditions) concentration of the hydrogen cation, that is used to define pH.

When we dissolve in water some substance that (preferentially) protonates it, an excess (H2O)nH+ forms and pH became lower than of pure H2O. Thus a Brønsted–Lowry acid (wrt H2O) becomes an Arrhenius acid. When we dissolve in water some substance that (preferentially) deprotonates it, (H2O)nH+ is consumed, an excess OH forms, and pH became higher than of pure H2O. Thus a Brønsted–Lowry base (wrt H2O) becomes an Arrhenius base (or alkali).

But the paragraph above make sense only in the aqueous framework. When we mix H2O with NH3, a huge amount of ions forms:

• OH, the conjugate base of H2O;
• NH4+, the conjugate acid of NH3 (in other words, an ammonia-based form of H+).

In Arrhenius picture, we see an alkali: there is almost no (H2O)nH+ and a lot of OH, whereas NH4+ isn’t “hydrogen cation”. But in NH3-centric picture we see indeed a strong acid! There is a lot of NH3’s conjugate acid, whereas NH3’s conjugate base (NH2) is absent. In “symmetric” Brønsted–Lowry picture we see that H2O acted as acid (it deprotonated) and NH3 acted as base (it was protonated).

Hence, acidic and basic solutions are all relative. Water is acidic, with respect to substances that are bases wrt H2O.

In your example, the hydrogen (+1) bonded to the hydroxide (-1) will balance out to a net charge of 0. Because there is no overall charge, the pH will be neutral.

However, pure, neutral water can behave as a Lewis Base (lone pair donor) due to the lone pairs present on oxygen. For example, water behaves as a Lewis Base and Al3+ as a Lewis Acid (electron acceptor) to form a hexaaquaaluminum(III) complex. That is: Al3+ with 6 water molecules surrounding it in a cubic fashion.

There seems to be a lot of confusion here - both in the question and in some of the answers. Let's try and clear a few things up.

The definition of an acid is a compound with a negatively charged hydroxide group

This is definitely wrong. There are many different definitions of acids and bases. The three most commonly used ones are:

1. Arrhenius theory: This is one of the oldest theories of acids and bases. It says that an acid is a species which dissociates in aqueous solution to give hydrogen ions in solution. A base is a species which gives hydroxide ions.
2. Brønsted-Lowry theory: This is really a generalization of the Arrhenius theory beyond aqueous solutions. It says that an acid is a proton donor (whether in aqueous solution or not) and a base is a proton acceptor.
3. Lewis theory: This is a much more general theory of acids and bases which has strong links to molecular orbital theory, although it was actually developed prior to this. It states that an acid is an electron pair acceptor and a base is an electron pair donor.

The most relevant of these theories to our discussion are the first two. Both of them state that an acid must have labile protons which it can easily lose. In both theories a hydroxide ion would usually be considered a base, not an acid.

Water has a pH of 7, though, meaning that it's neutral

There are two concepts here; neutrality and pH. These are often confused as meaning the same thing but they are actually completely independent. An aqueous solution is neutral when it contains equal numbers of hydrogen and hydroxide ions. pH is a measure of the hydrogen ion concentration in aqueous solution and is defined as: $$\mathrm{pH} = -\log_{10}(a_{\ce{H+}})$$

where $a_{\ce{H+}}$ is the hydrogen ion activity. For dilute solutions this is approximately numerically equal to the hydrogen ion concentration.

If we consider just pure water then there will always be a small amount of hydrogen and hydroxide ions in solution, due to the self-ionization of water. $$\ce{H2O + H2O <=> H3O+ + OH-}$$

As you can see from the equation, the hydrogen (or actually $\ce{H3O+}$) ion concentration will always be the same as the hydroxide ion concentration and so pure water will always be neutral. However, as you heat water, the dissociation equilibrium shifts towards more dissociation so the concentration of hydrogen ions goes up and the pH goes down.

Why isn't water acidic

Using the Brønsted-Lowry definition, we can describe any reaction involving proton transfer in terms of an acid and a base. Therefore, calling something an acid or base is only strictly relevant in the context of a specific reaction. $$\ce{H2O + HCl -> H3O+ + Cl-}$$ $$\ce{H2O + NaNH2 -> NaOH + NH3}$$

In the first reaction, water accepts a proton from hydrogen chloride, so it is acting as a base. In the second reactions, water donates a proton to the amide ion, so it is acting as an acid. So you can see that acidity is all relative.

A way of quantifying this relationship is to measure the $\mathrm{p}K_\mathrm{a}$ of a species. This is a measurement of the equilibrium position of the following reaction. $$\ce{HA + Solvent <=> A- + SolventH+}$$ It is defined in a similar way to pH: $$\mathrm{p}K_\mathrm{a} = -\log_{10}K_\mathrm{a}$$ $$K_\mathrm{a} = \frac{a_{\ce{A-}}a_{\ce{SolventH+}}}{a_{\ce{HA}}}$$

Tables of $p\mathrm{K_a}$ values can be found in places such as here.

Water shows its acidic nature when strong bases are used. Group I elements reacts with water (like any other acid) to liberate hydrogen gas. However, you have to remember that water is a polar compound. Oxygen has enough electronegativity to pull the bonded pair of electrons toward itself but hydrogen being a partial member of group 17, has the power (I mean electron withholding property) to resist the oxygen's effort to pull all electron's toward itself. Hence water is unable to show its acidic nature. Moreover, water has low ionisation potential to liberate hydrogen ions. Even if it dissociate, equal number of $\ce{OH-}$ and $\ce{H+}$ are produced to cancel each other's effect and remain neutral.