What Is A Buffer & How Does It Work?
A buffer is an aqueous solution that can resist significant changes in pH levels upon the addition of a small amount of acid or alkali. Each buffer is characterized by a set capacity, which is defined as the quantity of strong acid or base that must be added to change the pH of one liter of the solution by one pH unit. In other words, buffer capacity is the amount of acid or base that can be added before the pH begins to change significantly.
A buffer range is the specific pH range in which a buffer effectively neutralizes the added acid or base while maintaining a nearly constant pH. The capacity and range of a particular buffer ensure that the added small amount of acid/base is neutralized and the chemical reaction keeps going without giving a wrong outcome for the experiment/process. Simply put, a buffer is a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid.
Types of Buffers
Acidic Buffer
Acid buffer solutions have a pH of less than 7. It is generally made from a weak acid and one of its salts (often called conjugate*). Commonly used acidic buffer solutions are a mixture of ethanoic acid and sodium ethanoate in solution, which have a pH of 4.76 when mixed in equal molar concentrations. You can change the pH of the buffer solution by changing the ratio of acid to a salt or by choosing a different acid and one of its salts.
Alkaline Buffer
Alkaline buffer solutions have a pH greater than 7 and are made from a weak base and one of its salts. A very commonly used example of an alkaline buffer solution is a mixture of ammonia and ammonium chloride solution. If these were mixed in equal molar proportions, the solution would have a pH of 9.25.
What is a buffer composed of?
For a buffer to effectively maintain a specific pH range, it should be composed of a weak conjugate acid-base pair. This means that it can either be a weak base with its conjugate acid or a weak acid with its conjugate base, depending on the desired pH when preparing the buffer.
- Acetic acid (a weak organic acid with the formula CH3COOH) and a salt containing its conjugate base, the acetate anion (CH3COO-), like sodium acetate (CH3COONa), can function as buffers when combined in solution.
- Pyridine (a weak base with the formula C5H5N) and a salt containing its conjugate acid, the pyridinium cation (C5H5NH+), like Pyridinium Chloride, can also serve as buffers.
- Ammonia (a weak base with the formula NH3) and a salt containing its conjugate acid, the ammonium cation, like Ammonium Hydroxide (NH4OH), can be another example of a buffer system.
How Buffers Work?
Buffers work by neutralizing any added acid (H+ ions) or base (OH- ions) to maintain a moderate pH, making them a weaker acid or base. Let'sLet's take an example of a buffer made up of the weak base ammonia, NH3, and its conjugate acid, NH4+. When HCl (strong acid) is added to this buffer system, the extra H+ ions added to the system are consumed by the NH3 to form NH4+. Now, because all the extra H+ ions are locked up and have formed a weaker acid, NH4+, thus the pH of the system does not change significantly. Similarly, when NaOH (strong base) is added to this buffer system, the ammonium ion donates a proton to the base to become ammonia and water, thus neutralizing the base without any significant pH change.
Now there is a term we call ''Breaking of the buffer solution,'' which arises when the entire base and its conjugate acid (in the above case, NH3 and NH4+) are consumed to neutralize the added acid or base. The further addition of an acid or base to the buffer will change its pH quickly. Thus the breaking of the buffer is its capacity, or in other words, it is the amount of acid or base a buffer can absorb before breaking its capacity. It is to be noted that a solution with a weak base has a higher buffer capacity for the addition of a strong acid, and a solution with weak acid has a greater buffer capacity for the addition of a strong base.
A conjugate acid is a species formed by the gain of a proton by a base, while in reverse, a conjugate base is a species formed by the removal of a proton from an acid.
We can understand this by the following equation-
Acid + Base ⇌ Conjugate Base + Conjugate Acid
H2O (l) + NH3 (g) → OH− (aq) + NH4+ (aq)
In this equation, H2O is the conjugate acid, and its corresponding conjugate base is OH−, while NH3 and NH4+ represent the base and its conjugate acid pair. Here it is to be noted that the stronger the acid or base, the weaker the conjugate, and the weaker the acid or base, the stronger the conjugate.
Importance and Use of Buffers
Many solutions are acidic in nature, which can affect the outcome of chemical reactions. Buffer solutions help maintain a specific pH in a reaction medium, which is crucial for the reaction to occur or occur at a suitable rate. In particular, biochemical reactions are very sensitive to pH. The reason being many biological molecules contain groups of atoms that can be charged or neutral based on the pH of the environment. These charged or neutral groups significantly impact the biological activity of the molecule. The fluid within the cell and the fluids around the cells have a characteristic and almost constant pH in all multicellular organisms. This pH is preserved in several ways, and one of the most important is through buffer systems.
Buffer Solution Examples
Buffers can be found in various biological and chemical contexts, serving as crucial components in maintaining stable pH levels. Some examples of buffer solutions include:
- Blood: The bicarbonate buffer system is an essential component of blood, working to maintain a constant pH level, which is crucial for various physiological processes.
- Tris buffer: Tris (hydroxymethyl) aminomethane is a commonly used buffer in molecular biology, biochemistry, and electrophoresis, as it has a relatively stable pH in the range of biological interest (pH 7.0 to 9.0).
- Phosphate buffer: This buffer system is widely used in biological and biochemical research because of its buffering capacity within the pH range of 6.0 to 8.0. It is particularly useful in cell culture applications, as it helps maintain stable pH levels in culture media.
- Citrate buffer: Citrate buffers are commonly used in food, pharmaceutical, and chemical industries due to their buffering capacity over a wide pH range (2.5 to 7.5). They are particularly useful in enzyme-based applications, as they help maintain optimal pH levels for enzyme activity.
Universal Buffers
Most buffers work over a relatively narrow pH range. Citric acid is an exception to the rule because it has three pKa values. This allows for a larger pH range to be available for a buffer when a compound has multiple pKa values.
Moreover, combining buffer solutions is possible as long as their pKa values are close, differing by 2 or less. The pH can be adjusted to achieve the desired range by using a strong base or acid. One example of a combined buffer is McIvaine's buffer, which involves mixing Na2PO4 and citric acid in varying ratios to achieve an effective buffer range from pH 3.0 to 8.0. Another example is a mixture of citric acid, boric acid, monopotassium phosphate, and barbituric diethyl acid, which can cover a broad pH range from 2.6 to 12.
Meanwhile, by understanding the unique characteristics and applications of each buffer solution, researchers and scientists can ensure that they select the most suitable buffer system for their specific needs, thereby maintaining stable pH levels and promoting the desired outcomes in their experiments and processes.
Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation is a formula used to calculate the pH of a buffer solution. It relates the pH of the buffer solution to the dissociation constant (pKa) of the weak acid or base used in the buffer and the ratio of the concentrations of the conjugate base and acid. It is commonly used in chemistry and biology to design and prepare buffer solutions for various applications.
The equation is:
pH = pKa + log([conjugate base]/[weak acid])
Acidic and basic buffer solutions can be created using the equation.
Acid buffer solution: a weak acid (HA) + its salt (KA) + a strong base (KOH).
The ionization of weak acid HA can be written as:
HA + H2O ⇋ H+ + A−
and the acid dissociation constant is:
Ka = [H+] [A–]/HA
The Henderson-Hasselbalch equation, pH of acid buffer = pKa + ([salt]/[acid]), can be derived from taking the negative log of RHS and LHS.
Similarly, base buffer solution: a weak base (B) + it's salt (BA) + a strong acid.
This equation is used to calculate the pH of the buffer, and pKa value, and prepare a buffer solution of a needed pH.
Limitations of the Henderson-Hasselbalch equation: Although very useful, it cannot be used for strong acids and strong bases.
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FAQs
What is a ph buffer?
A pH buffer, also known as a buffer solution, is a chemical system designed to maintain a stable pH level even when small amounts of acids or bases are added. Buffer solutions are crucial in various biological, chemical, and industrial processes, as they help to prevent significant fluctuations in pH that could potentially disrupt the intended reactions or negatively impact the system.
A pH buffer consists of a weak base and its conjugate acid or a weak acid and its conjugate base. The balance between these two components allows the buffer to neutralize the added acid or base without significantly changing the overall pH of the system.
What Is a Buffer in Chemistry?
Buffers play a vital role in maintaining stable pH levels in chemical systems. They achieve this by neutralizing small amounts of added acid or base, thereby preventing significant fluctuations in pH. This pH maintenance is crucial because many chemical reactions are sensitive to changes in pH, which can affect the reaction rate and product formation or even cause irreversible damage to molecules involved in the reaction.
What Is a Buffer in Biology?
Buffers are essential for maintaining optimal pH conditions for cellular processes and enzyme activity. Most enzymes exhibit maximal activity within a narrow pH range, and any deviation from this range can lead to a decrease in efficiency or even denaturation of the enzyme. Buffers help to maintain a stable pH in cellular environments and biological fluids, such as blood and extracellular fluids.
For instance, the human body has several buffer systems in place to regulate blood pH, including the bicarbonate buffer system. The bicarbonate buffer system maintains blood pH within a narrow range (7.35 to 7.45) by neutralizing excess acids or bases, ensuring the proper functioning of various physiological processes.
Is water a buffer solution?
In practical terms, however, water is an extremely poor buffer because it can't evolve more H+ than OH– because both evolve from the same reactant at a 1:1 stoichiometric ratio. Keep in mind that a buffer has to be able to compensate for small changes in pH. This is something water alone cannot do.