Buffers and pH

Water is a key ingredient of life. One of water molecule’s properties that makes it so indispensible is its ability to dissociate into proton and hydroxide ions:

 This dissociation happens only in a very small fraction of water molecules, but it controls many chemical processes, especially in biology. Of a particular importance is concentration of protons H⁺ in solution because protons are universally involved in many acid-base reactions and equilibria.

pH

Because concentration of protons in solutions may vary by 14 orders magnitude, it is practical to express its value in a negative logarithm form or pH:

Acidic solutions have higher concentration of protons and pH value less than 7, alkaline solutions have pH value above 7. Life occurs mostly around neutral conditions with pH 6-8;

pK

Many functional groups can bind and dissociate protons. This protonation/deprotonation equilibrium is controlled by concentration of protons in the solution :

For each protonatable group there is such pH value where half of molecules are in protonated from and half is in deprotonated form. Such pH value is called a pK of a group and is its important characteristic that expresses affinity of a group toward protons. 

Buffers

Because of its role in stability and reactivity of biological samples, pH of the medium has to be carefully controlled. In the laboratory this is achieved using pH buffering systems – equilibrium solutions with low chemical reactivity that allow adjusting and maintaining pH within given range. Buffering systems, or buffers, always consist of an acid, base, and protons.

One of the ions in the buffering system, either acid or base, has to be weak, i.e. does not dissociate readily. The counter-ion is always strong. Phosphoric acid is a common example of a weak acid-based buffer, while ammonium is a weak base. There is a great variety of inorganic and organic buffers that cover entire pH range.

To use buffers effectively, it is important to understand how they work. Buffers work best for controlling concentration of protons when half of its molecules have lost protons (deprotonated) while another half remains protonated, or around its pK value.  Whenever protons are added to or withdrawn from the solution, protonation equilibrium of salt shifts and buffer responds by binding or releasing protons thus establishing a new equilibrium close to the original. At pH equal pK value of an acid buffer has greatest latitude in absorbing jump in concentration of protons, also called buffer capacity.

At pH values of one pH unit apart from pK, protonation equilibrium of the buffer shifts to 5 or 95% in base and acid. Buffering capacity is significantly reduced at these conditions. Thus, practical use of a buffer is limited to the range of one pH unit away from pK value of the acid. Many weak acids have several protonation groups, making then useable in several pH ranges. It is also possible to prepare cocktails of buffers so that buffer capacity of one buffer picks up when another is exhausted.

Practical consideration in selecting and preparing buffers:

• Select desired pH range.

• Select  buffer that:

• has pK is close to desired pH, best within 0.5 pH units;
• has no spectral interference – most organic buffers have pronounced UV and IR absorption;
• has sufficient solubility;
• does not interact chemically with other components.
• etc .

• Prepare stock solution of acidic or alkaline form of buffer at ×2-4 final concentration.

• Using pH meter and magnetic stirrer adjust pH of stock to almost desired pH (0.1 -0.2 pH units) using 2-5M solution of HCl or NaOH.

• Dilute buffer to desired concentration. Typical values range from 20 to 200 mM, with 50-100 mM being the most common. 

• Check actual pH and make final pH adjustment.