The Hydrogen Ion

In order to answer that question, we have to look at the defining feature of an acid. The Hydrogen Ion (H+) is the primary feature of an acid, and the relationship between the Hydrogen ion and acid (or base) molecules is the big concept to understand when it comes to acids.

So why Hydrogen? What makes the Hydrogen Ion so important when it comes to acids and bases? Well, let’s think about what makes up a Hydrogen Ion.

Hydrogen is element 1 on the periodic table. This means that a hydrogen atom contains one proton and one electron. Now, the other major component of atoms is the neutron. Hopefully you know that an element can have a varied number of neutrons, though it might not be stable with the wrong number. These atoms with varying numbers of neutrons are called isotopes.

Hydrogen has three main isotopes, $^1H$, $^2H$, and $^3H$, each containing 0, 1, and 2 neutrons, respectively. Perhaps surprisingly, the most common isotope for Hydrogen is the $^1H$ isotope, which does not contain any neutrons at all. This isotope actually comprises over 99% of all Hydrogen!

An ordinary hydrogen-1 atom has one proton and one electron, with no neutron. Hydrogen-2 and hydrogen-3 isotopes have one and two neutrons respectively. Losing the electron leaves the hydrogen-1 nucleus, a proton; in water it is hydrated rather than an isolated bare proton.

In water, H⁺(aq) denotes the hydrated proton, commonly represented as H₃O⁺. It is not a bare solitary proton freely approaching atomic nuclei. Hydration structures and proton transfer involve the surrounding water molecules.

A Brønsted–Lowry acid donates a proton to a base. For example, HCl+H₂O→H₃O⁺+Cl⁻. Water acts as the proton acceptor in that reaction; atoms and total charge are conserved.

Removing a proton from an acid leaves its conjugate base. The ability of a species to donate or accept a proton depends on its bonding, charge, electronic structure and solvent environment, rather than a neutron count rule.

Use molecular structures and a balanced proton-transfer reaction to identify acid–base partners. Compare acid strengths by the specified equilibrium constants and solvent conditions; nuclear size is not an AP acid-strength model.

Proton transfer can occur among water molecules and other bases in solution. The source bare-sphere drawing is superseded by a molecular hydrated-proton/proton-transfer representation.

Keep nuclear composition and acid–base behavior as distinct models. Hydrogen isotope differences can affect physical chemistry, but “absent neutrons allow acidity” does not explain ordinary aqueous proton transfer.

At some point, you have probably encountered the term pH. You may even know it has something to do with acids. What is pH, though, really?

Due to the great importance of the Hydrogen Ion we just reviewed, chemists like to look at the concentration of Hydrogen in solution quite often. So often, in fact, that they created the term pH to describe just that.

In formal terms, pH is calculated via the following equation:

$pH = -log[H^+]$

The concentrations that the Hydrogen typically exists at in solution are much lower than what you are used to. Normally the lowest concentrations you have been interested in are probably down to 0.001M. If you had an $H^+$ concentration of 0.001M, it would actually be really high!

Typically, the pH scale is given as 0-14. This is the typical range you will be dealing with when learning about acids and bases. Let’s convert that scale to actual $H^+$ concentrations.

$14 = -log[H^+]$
$-14 = log[H^+]$
$[H^+] = 10^{-14}$
$[H^+] = 0.00000000000001M$

As you can see, the Hydrogen Ion concentration can reach extremely low values and still be relevant.

As we said before, when the Hydrogen Ion is around any other molecule or atom, it is unavoidably attracted to it. Therefore, the lone proton does not exist in solution. In water, the Hydrogen Ion bonds itself to water molecules, forming what is known as the Hydronium Ion, $H_3O^+$.

$H_2O+H^+\rightarrow H_3O^+$

Normally, this is assumed, and we still look at pH as being the concentration of $H^+$ in the water, even though it actually is $H_3O^+$.

It is important to note here that the water molecule is accepting a Hydrogen Ion when it forms the Hydronium Ion. As you will find out later, this is one property of a base. Water is a special substance, as it can also release a Hydrogen Ion in the right environment, forming Hydroxide.

$H_2O\leftrightarrow OH^-+H^+$

Water autoionization has K_w=[H₃O⁺][OH⁻] in the ideal concentration convention. K_w depends on temperature. At 25 °C it is approximately 1.0×10⁻¹⁴, which lets us determine the neutral pH at that temperature.

Example Problem 6.1.1
The Equilibrium constant for the deprotonation of water is $K = 10^{-14}$. Use this information to determine the pH of pure water. The equation for the deprotonation of water is:
$H_2O\leftrightarrow H^++OH^-$

We tend to say that a solution with a pH below 7 is acidic, and a solution with a pH above 7 is basic. Most natural systems are within one or two pH units of neutral, but some surprisingly strong acids or bases can be found in nature. One in particular resides in your body! Stomach acid has a pH near 1. This makes it very acidic, and that acidic property is used to digest your food.

On the next column, you can see a cool chart outlining the pH of various common solutions and chemicals.

Sometimes you will see something known as the pOH. This is very similar to pH, as its formula is nearly the same, but with the Hydroxide Ion instead of the Hydrogen Ion.

$pOH = -log[OH^-]$

Thanks to the dissociation reaction of water, we can give a nice relationship between pH and pOH:

pH+pOH=pK_w(T). At 25 °C this becomes approximately pH+pOH=14.

The reason for this relationship relies on the Equilibrium Equation for the water dissociation. Recall that for water, $K = 10^{-14}$. Let’s use that to prove the relationship between pH and pOH.

$[H^+][OH^-]=10^{-14}$
$log{[H^+][OH^-]}=-14$
$log[H^+] + log[OH^-] = -14$
$-log[H^+]-log[OH^-]=14$
$pH+pOH = 14$

The relationship follows from water autoionization. Use pK_w at the stated temperature; the value 14 is the approximate 25 °C value. Other solvents have their own acid–base and autoionization conventions.