Hydrochloric acid in solution donates a proton to the solvent water and behaves as a Bronsted acid:

HCl + H₂O ----> H₃O⁺ + CL^-

H₃O⁺ is referred to as the hydronium ion. For convenience, will refer to the hydronium ion as a hydrogen ion [H⁺] . The basic properties of ammonia are clearly accounted for by this theory. Ammonia accepts a proton from the water producing OH^- :

H₂O + NH₃ ----> OH⁺ + NH₄⁺

The terms acid or base strength and concentration are easily confused. Strength is a measure of the degree of dissociation of an acid or a base in solution, independent of its concentration. Concentration refers to the amount of acid or base per amount of solution (volume).

The strength of acids and bases in water depends on the extent that they react with the water solvent. Acids and bases are considered strong when they react nearly 100%, that is that they nearly all dissociate. A weak reaction with water is much less than 100% complete.

Important strong acids include the following:

This reaction has little or no tendency to proceed in the reverse direction. All of the molecules are converted to ions.

All strong bases are metal hydroxides. Strong bases completely dissociate, or ionize, in aqueous solutions to produce hydroxide ions and metal cations. Of the common metal hydroxides, only NaOH and KOH are soluble in water and are the only readably usable strong bases:

Weak acids and bases dissolve in water principally in the molecular form. Only a small percentage of the molecules dissociate to form the hydronium ion or the hydroxide ion.

The most important weak acids are the following:

The most fundamental difference between strong acids and bases is their equilibrium situation.

A strong acid such as HCl does not exists to any measurable degree in equilibrium with it's ions, H⁺ and CL^- . On the other hand, a weak acid such as acetic acid establishes a dynamic equilibrium (LaChatelier's principle) with it's ions , H⁺ and C₂H₃O₂ ^- . This equilibrium may be represented as the ratio of concentration (bracket like this [ ] indicate concentrations that are molar or M) of products and reactants.

Ratio of the concentration:

of products [H⁺] [C₂H₃O₂ ^-]

to reactants [HC₂H₃O₂]

At a fixed temperature and in a specified solvent this ratio is a constant value, Ka:

Ka = [H⁺] [C₂H₃O₂ ^-] / [HC₂H₃O₂]

where Ka is the equilibrium constant or acid dissociation constant for the given acid. The magnitude or size of this constant shows the strength of the acid. The larger the Ka the (large numerator small denominator), the more dissociated (stronger) the acid will be. The situation for bases is the same.

There are many table listed in books and scientific literature for these constants and are useful in performing calculations related to acid base solution composition. Calculations of this type are essential to many areas of chemistry.

Although pure water is virtually 100% molecular (non dissociated), a small number of water molecules do ionize. This process occurs by the transfer of a proton of one hydrogen ion to another water molecule. Again we will not be concerned with the resulting hydronium ion (H₃O⁺) but only the hydrogen ion (H⁺). This equilibrium is shown below:

H₂O ----> H⁺ + HO^-

Kw = [H⁺] [HO^-]

Kw is a constant because its value does not depends on the nature or concentration of the solute, as long as the temperature does not change. Kw is a temperature-dependent quantity.

The pH scale correlates the hydrogen ion concentration with a number, the pH, serves as a useful indicator of the degree of acidity or basicity of a solution. To help develop the pH concept consider:

- Addition of an acid (proton donor) to water increases the [H⁺] and decreases the [HO^-]

- Addition of an base (proton acceptor) to water decreases the [H⁺] and increases the [HO^-]

- [H⁺] = [HO^-] equal amounts of acid and base are present

The pH of a solution may be calculated if the concentration of either H⁺ or HO^- is known. Conversely, measurement of pH allows the calculation of H⁺ or HO^- concentration.

In biochemistry the maintenance of pH is vital. Alteration of our blood pH by as little as 1 pH unit would lead to a coma, shutting down of most internal organs and resulting in a death that even young Dr Carter from ER couldn't fix. A buffer solution contains components that enable the solution to results large changes in pH when acids or bases are added.

Consider the case of acetic acid we talked about earlier. A buffer solution may be prepared by prepared from the addition of an acid and base. An equilibrium is established in solution between the weak acid and conjugate base (it's salt ion).

A buffer solution functions in accordance with LaChatelier's principle, which states that an equilibrium system, when stressed, will shift its equilibrium to alleviate that stress. Meaning that if you were to add one of the species in the above equation the amount of the remaining species would shift in concentration to reestablish the equilibrium. OR if you add H⁺ by adding an acid, then the concentration of acetic acid and water will increase to re-adjust for the displacement from equilibrium. The magnitude of shift depends on the Ka for the chemicals in question.

ex. Addition of base (OH^-) to a buffer system causes the following changes:

- OH- from the base reacts with H+ producing water (reducing the amount of H+ present)

- Acetic acid dissociates to replenish the H+ consumed by the base, maintaining the pH close to its initial level

or: Addition of acid (H+) to a buffer system causes the following changes:

- H+ from the acid increases the overall H+ concentration

- The system reacts to the stress to from more acetic acid (molecular acid), by combining acetate (conjugate base) with water

In other words Adding OH^- shifts the equilibrium to the right and adding H⁺ shifts the equilibrium to the left. Note that buffering against base is a function of the concentration of the weak acid and buffering against acid is a function of the concentration of the conjugate base.

A buffer by definition, resists changes in the pH of a solution. A buffer must contain the chemical species for "neutralizing" added amounts of acid or base. Generally a good buffer is a solution of a weak base and its conjugate base, or a weak base and it's conjugate acid (e.g.. ammonia and ammonia chloride). A buffer is selected on the basis of its pKa value and its chemical nature.

The Henderson-Hasselbach equation (given below) is the relationship between pH, pKa and the ratio of the concentration of the salt and acid forms of the buffer.

pH = pKa + log [A-]/[HA]

As shown by this equation, when the concentration of the conjugate base and the un-dissociated acid are equal, [A-] = [HA], the pH of the solution is equal to the pK of the buffer

pH = pKa + log 1

log 1 = 0

pH = pKa

When [A-] = 10 x [HA], then:

log [A-]/[HA] = log 10 = 1

and

pH = pKa + 1

When [A-] = 1/10 [HA], then;

log [A-]/[HA] = log 1/10 = -1

and

pH = pKa - 1

Thus, buffers are most effective in the range of one pH unit of the pKa. Once beyond that, there is very little buffering capacity left of the reacting species to continue to resist pH changes. The molar strength of the two components of a buffer should be chosen to give adequate buffering capacity. The total ionic strength of the buffer may also be an important consideration - for example, in enzyme studies. Buffer strength is expressed expressed in terms of the total concentration of conjugate base or conjugate acid. Using a buffer called Tris as an example.

The pKa for Tris is 8.3; therefore this substance can effectively serve as a buffer between pH 7.8 and 9.3 . The molecular weight is 121.1g. To make 0.25 liter of a 50 mM Tris buffer with pH 7.8;

50 mM = 0.05 M =(g/mw)/liters = (g/121.1)/0.25 = 1.51 g of Tris solid

However, the pH of the buffer would have to be adjusted to 7.8 by adding acid: HCl would be the best choice because the Cl- made in adjusting the pH is not a problem with most biological systems.

The chemical nature of a buffer is especially important in biological investigations. Solubility, stability, defined interactions with other ions in the system, light absorption and potential inhibition or stimulation of the function of the biological system are all criteria to be considered in buffer selection.