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Buffers

• Acetate buffer: acetic acid/sodium acetate (CH₃COOH)/CH₃COONa), pH ~ 4.7
• Ammonia buffer: ammonia/ammonium chloride (NH₃/NH₄Cl), pH ~ 9.3
• Hydrogen carbonate (bicarbonate) buffer: carbonic acid/sodium hydrogen carbonate (H₂CO₃/NaHCO₃), pH ~ 6.4
• Hydrogen carbonate (bicarbonate) buffer: sodium hydrogen carbonate/carbonate (NaHCO₃ / Na₂CO₃), pH ~ 10.3
• Phosphate buffer: sodium dihydrogen phosphate/disodium hydrogen phosphate (NaH₂PO₄/Na₂HPO₄), pH ~ 7.2

Buffers can also be made by starting with just weak acid and then neutralizing a portion of that weak acid with a strong base (achieving the mixture of weak acid and its salt by neutralization).

pH of a buffer
A buffer can resist pH change because the two components (conjugate acid and conjugate base) are both present in appreciable amounts at equilibrium and are able to neutralize small amounts of other acids and bases (in the form of H⁺ and OH^-) when they are added to the solution.

Acidic buffers
Acidic buffers are buffers that maintain a solution at pH below 7 and typically consist of a weak acid and one of its salts to provide the anion which acts as a base (its conjugate base). The function of the salt component is to act as a source of ions (a supply of extra A^- by adding a soluble salt of HA which fully ionizes such as NaA).

They work because the dissociation of a weak acid is an equilibrium reaction,

HA ⇌ H⁺ + A^-.

The conjugate base A^- removes the added acid: when an acid is added the [H⁺] increases and these H⁺ react with the conjugate base (A^-) to produce undissociated HA. The equilibrium shifts to the left (in accordance with Le Chatelier’s principle), removing most of the H⁺ ions. Since HA is a weak acid, the A^- will tend to combine with any available H⁺ to form HA, thus the hydrogen ion (H⁺) concentration (i.e., the pH) will change only slightly.

The weak acid removes the added alkali: when an alkali is added the [OH^-] increases. The small amount of H⁺ present reacts with the OH^- to form H₂O. The weak acid dissociates, causing the equilibrium to shift to the right to restore most of the H⁺ ions, so the pH tends to remain almost the same.

Basic buffers
Basic buffers are buffers that regulate the pH around values ​​greater than 7 and typically consist of a weak base and a salt of its conjugated acid. Similarly, to acidic buffers, in basic buffers the weak base removes the added H⁺ and the conjugated acid removes added OH^-.

Neutral buffers
They maintain a solution at pH value of around 7 and consist of polyprotic acid salts. In some cases, both the one that plays the role of the weak acid, and its conjugate base are salts from the partial neutralization of an acid that has several protons, such as phosphoric acid.

The pH of a buffer solution depends on the dissociation constant, K_a of the acid (or K_b of the base) as well as on the ratio of acid (base) and its conjugate base (acid) concentrations. This dependence is described by the well-known Henderson-Hasselbalch equation which is often used in chemistry and biochemistry to perform the necessary calculations in the preparation of buffer solutions for use in the laboratory, or in other applications. Using the Henderson-Hasselbalch equation, the exact amount of acid and conjugate base needed to obtain a buffer of a given pH can be determined.

Henderson-Hasselbalch equation (buffer equation)
For a weak acid HA and its conjugate base A⁻:

pH = pK_a + log [HA]/[A^-]

The Henderson-Hasselbalch equation is derived from the acid ionization (dissociation) constant:

K_a = [H⁺][A^-]/[HA]

Taking the negative log on both sides of the equation gives:

-log K_a = -log [H⁺][A^-]/[HA]

This corresponds to:

-log K_a = -log [H⁺] - log [A^-]/[HA]

Relatively:

-log K_a = -log [H⁺] + log [HA]/[A^-]

By definition:

-log K_a = pK_a

-log [H⁺] = pH

Thus:

pK_a= pH + log [HA]/[A^-]

This equation is then arranged to:

pH = pK_a + log [HA]/[A^-]

Buffers are characterized by the pH range over which they can maintain a more or less constant pH and by their buffer capacity, the amount of strong acid or base that can be absorbed before the pH changes significantly.

Buffer range
Each conjugate acid-base pair has a characteristic pH range where it works as an effective buffer. Useful range is when pH = pK_a ± 1. This represents the point in the titration that is halfway to the equivalence point. This region is the most effective for resisting large changes in pH when either acid or base is added. Based on the Henderson-Hasselbalch equation, it can be seen that the pH of the buffer solution is equal to the pK_a value of the acid when the concentrations of the conjugate acid and conjugate base are approximately equal (within about a factor of 10 of each other), because log(1) = 0. In other words, the ratio of base-to-acid in the buffer, must be greater than 0.1 but less than 10. As soon as this ratio becomes outside this interval, then the buffer range has been exceeded, and we no longer expect the mixture to neutralize external acids/bases effectively.

Buffer capacity
Buffer capacity may be defined as the quantity of a strong acid or strong base that must be added to one liter of a solution to change it by one pH unit.

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. Although the useful pH range of a buffer depends largely on the chemical properties of the weak acid and weak base used to prepare the buffer (i.e., on K_a or K_b), its buffer capacity depends solely on the concentrations of the species in the buffer solution. Buffers with a high capacity contain large concentrations of buffering components (acid/conjugate base) that react with added H⁺ or OH^- ions to maintain pH. Therefore, the concentrations of the acid and its conjugate base in a buffer will determine how much additional acid or base can be added to the solution before its buffering ability is exhausted. The higher the concentrations of acid and conjugate base, the larger the buffer capacity.

The buffer capacity equation is as follows:

β = Δn/ΔpH

Δn is the number of moles of acid or base added per liter of buffer solution
ΔpH is the change in pH

Applications
Buffer solutions are very useful in a wide range of applications where a stable pH is needed. They play an important role in maintaining pH in many chemical and biochemical processes.
Buffer solutions are widely used in industry. Industrial processes requiring buffer solutions include fermentation (e.g. in breweries or bakeries), control of dyeing processes in textile industries, manufacture of pharmaceuticals, manufacture of personal hygiene products and cosmetics, in the food industry to maintain the physical, chemical and microbiological stability of foods, etc. They are also used in chemical analysis and calibration of pH meters.

Important buffers

Some reagents used for buffer preparation

Buffers for calibration of pH meters or pH electrodes
The measurement of pH is essential in practically every industry and laboratory involving any chemical or biochemical process or analysis. A pH meter is an instrument that measures the hydrogen-ion activity in aqueous solutions, indicating their acidity or alkalinity expressed as pH. The pH meter measures the difference in electrical potential (voltage) between a pH electrode and a reference electrode and displays the result converted to the corresponding pH value. One of the electrodes is sensitive to hydrogen ions (usually a glass electrode) and the other is the reference electrode (e.g. a silver chloride electrode). They are often combined as one compact electrode together with a temperature probe. The electrodes, or probes, are immersed into the solution to be analized.

A pH meter must be calibrated against a known standard to ensure accuracy each time it is used or before each series of measurement. The more accurate the results need to be, the more often it must be calibrated. Standards are buffer solutions with a known pH (e.g. 4.00, 7.00, 10.00). Calibration is required using at least two standard buffer solutions that span the range of pH values to be measured. Temperature can affect pH readings. To obtain an accurate reading, the pH meter must be calibrated at the same temperature as the samples to be analized.

We offer a wide range of pH buffer solutions under the PanReac AppliChem brand, traceable to NIST SRMs, for the calibration of pH meters. All these solutions are very stable and can be stored for long periods of time (up to 6 years). In most cases an accuracy of ± 0.02 pH units at a temperature of 20 °C is guaranteed. For each buffer solution you will find the pH variation with temperature printed on the label or on its TDS.

Ready-to-use buffer solutions (colorless)

Ready-to-use buffer Solutions (color coded)
The colored buffer solutions are dyed in different colors so that they can be easily identified when working.

Knowing the pH value of substances or solutions is important or critical in many situations, including laboratory chemical analysis. pH meters are used for soil measurements in agriculture, water quality for municipal water supplies, swimming pools, environmental monitoring; brewing of wine and beer; manufacturing, life sciences and clinical applications such as blood chemistry; and in many chemical and pharmaceutical industries ranging from laboratory experimentation to quality. They are used to analyze the exact pH value of food and chemical products to ensure safety and quality, or they can be used to evaluate acidity/alkalinity of drugs in pharmaceutical and biotechnology industries.

Sectors

• Analytical Laboratories, Quality Control, Chemical and Biochemical labs
• Research, Development and Innovation (R&D)
• Biopharmaceutical and Pharmaceutical development and manufacturing
• Life Sciences
• Cosmetics and Healthcare
• Food and Beverages
• Inorganic synthesis
• Educational or academic institutions
• Paper industry
• Textile industry
• Soap and detergent industry
• Metallurgy
• Agriculture
• Pool maintenance

Important buffers in living systems
Buffering is important in living systems as a means of maintaining a stable, relatively constant internal environment, also known as homeostasis. Small molecules such as bicarbonate and phosphate provide buffering capacity as do other substances, such as hemoglobin and other proteins.

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 to the environment, e.g., by exhaling CO₂.

Hydrogen carbonate (bicarbonate) buffer
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. The maintenance of blood pH is regulated via the bicarbonate buffer. This system consists of carbonic acid and bicarbonate ions. When the blood pH drops into the acidic range, this buffer acts to form carbon dioxide gas. The lungs expel this gas out of the body during the process of respiration. Under alkaline conditions, this buffer returns pH to neutral by causing excretion of bicarbonate ions through the urine. This maintains the blood pH between 7.35 and 7.45 and balances 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.

Phosphate buffer
The phosphate buffer system acts similarly to the bicarbonate buffer, but has much stronger action. The internal environment of all cells contains this buffer comprising hydrogen phosphate ions and dihydrogen phosphate ions. Under conditions when excess hydrogen enters the cell, it reacts with the hydrogen phosphate ions, which accept them. Under alkaline conditions, the dihydrogen phosphate ions accept the excess hydroxide ions entering the cell.

Protein buffer
Proteins consist of amino acids linked together by peptide bonds. Amino acids possess an amino group and a carboxylic acid group. At physiological pH, the carboxylic acid exists as the carboxylate ion (COO^-) with a negative charge and the amino group exists as the NH₃⁺ ion. When the pH becomes acidic, the carboxyl group takes up excess hydrogen ions to return back to the carboxylic acid form. If the blood pH becomes alkaline, there is a release of a proton from the NH₃⁺ ion, which takes the NH₂ form.

Hemoglobin buffer
The respiratory pigment present in blood, hemoglobin, also has buffering action within tissues. It can bind with either protons or oxygen at a given point of time. The binding of one releases the other. In hemoglobin, the binding of protons occurs in the globin portion whereas oxygen binding occurs at the iron of the heme portion. During exercise, protons are generated in excess. Hemoglobin helps in the buffering action by binding these protons, and simultaneously releasing molecular oxygen.

Life Science buffers
Many biochemical processes are significantly impaired by even small changes in pH value. Therefore, it is often necessary to stabilize the pH by adding a suitable buffer to the medium, without affecting the system under investigation. Tris, Glycine, HEPES, PBS buffers are some examples of our life science buffers. See also more information on our Life Science buffers section.