Buffers

179 products available

Previous 1 2 3 4 5 6 7 8 9

Buffer

A buffer is a mixture of substances whose pH value (concentration of oxonium ions) changes much less when an acid or a base is added than it would be the case in an unbuffered system. The effect of the buffer is based on the conversion of the oxonium ions (H3O+) or hydroxide ions (OH-) supplied by the acid or base to weak acids or bases, which themselves have little tendency to form H3O+ or OH- ions.
Buffers are the aqueous buffer solutions specifically prepared in chemistry. More complex buffer systems are found in bodily fluids such as blood, or in groundwaters that interact with humus, for example.

Chemical principles

Buffer solutions contain a mixture of a weak acid and its conjugate or corresponding base (or the respective salt) or a weak base and its conjugate or corresponding acid. Ampholytes (bifunctional molecules) can also serve as buffers. The factor determining the pH is the ratio or protolysis equilibrium of the buffer pair.

The following applies to the acid-base equilibrium of an acid HA:
KS = c(H3O+) ⋅ c(A-) / c(HA)

By transforming, we obtain:
c(H3O+) = KS ⋅ c(HA) / c(A-)

Forming the negative decadic logarithm from this, we obtain:
- lg c(H3O+) = - lg Ks - lg c(HA) / c(A-)

This corresponds to:
pH = pKS - lg c(HA) / c(A-)

Relatively:
pH = pKS + lg c(A-) / c(HA)

Henderson-Hasselbalch equation (buffer equation)

This equation - which applies under the approximation that the activities of the substances correspond to their concentrations in solution - can be used to determine the concentration ratio of acid and base for a given pH value, given a known pKs value. The higher the concentrations, the lower the effect of additions of acids or bases. The amount of strong base (or acid) that can be absorbed by a buffer solution without significantly changing the pH is expressed by the buffer capacity.
Examples of buffer solutions are the acetic acid/acetate buffer or the ammonium buffer made from ammonium ions and ammonia.

Types of buffer systems

The carbonate buffer (a mixture of carbonic acid and hydrogen carbonates) regulates the CO2 concentration between the atmosphere, oceans and the biosphere. It is also the main part of the blood buffer. This maintains the pH of the blood between pH 7.35 and 7.45 and balances the fluctuations caused by metabolism. When the pH is below 7.35, it is called acidosis, and above 7.45, alkalosis. Death occurs at pH values below 6.8 or above 8.0.
When considering a buffer system, a distinction must be made between closed and open buffer systems. In a closed buffer system (e.g. acetic acid/acetate buffer), the protons (H+) or hydroxide ions (OH-) produced during a chemical reaction are trapped by the buffer substance. They react to form the corresponding or conjugated acid or base of the buffer and thus remain in the solution. In an open buffer system (e.g. the hydrogen carbonate/CO2 buffer system in the lungs), the system is in exchange with the environment. It is able to maintain the appropriate pH by releasing a component into the environment, e.g., by exhaling CO2.

Importance of buffer systems

Buffers have an important role in technical chemistry, such as electroplating or analog photography, as well as in analytics.
Buffer systems also play an important role in soil science; see Buffer area (soil science).
Significance in the life sciences: Buffers are essential for many animals and not least for the human organism. For example, human blood plasma and many enzymes depend on a constant pH value. Without buffers, even the smallest amounts of acid - e.g. lactic acid from energy metabolism - would be enough to paralyze the organism because various proteins would denature and thus become unusable.