Buffer Solutions – Definition, Types, Preparation, Uses
When a few drops of a strong acid or base are applied to water, the hydrogen ion concentration completely changes. In many industrial, chemical, and biological operations, it is necessary to have a solution whose pH value does not change significantly when small volumes of strong acids and strong bases are added. Many fluids, such as blood, have specific pH values, and variations in these values indicate that the body is malfunctioning. Controlling pH is also crucial in a wide range of chemical and biological activities. For the manufacture and usage of many medical and cosmetic products, a specific pH is required. Buffer solutions are what these solutions are called.
What is Buffer Solution?
A buffer solution is a solution that resists changes in hydrogen ion concentration when a modest amount of acid or base is added. In other words, such solutions are known to have reverse acidity and reverse basicity and to keep a reasonably steady pH value. A good example of a natural buffer solution is human blood. Despite eating a wide array of meals, our blood maintains a pH of roughly 7.35.
Buffer Solution is a water-based solvent-based solution made up of a weak acid and its conjugate base, or a weak base and its conjugate acid. They are resistant to changes in pH caused by dilution or the addition of relatively small amounts of acid or alkali. When a small amount of strong acid or strong base is added, the pH of the buffer solution changes very little. As a result, they’re used to maintaining a steady pH. When a tiny amount of strong acid or base is given to it, its pH varies very little, and it is thus used to keep a solution’s pH stable.
A buffer solution is one that can maintain its hydrogen ion concentration (pH) with just slight dilution or the addition of a small amount of acid or base. Fermentation, food preservation, medicine administration, electroplating, printing, enzyme activity, and blood oxygen-carrying capability all require particular hydrogen ion concentrations (pH) in buffer solutions. Buffer solutions are made up of a weak acid and its conjugate base or a weak base and its conjugate acid that can maintain pH.
Types of Buffer Solution
Acidic and basic buffers are the two types of buffer solutions that are extensively classified. These are discussed in greater depth down below.
- Acidic Buffers- Acidic buffer solutions are made up of equimolar amounts of a weak acid and its salt, as well as a strong base. These solutions are used to keep the environment acidic. Acid buffer is made by combining a weak acid and its salt with a strong base to create an acidic pH. The pH of an aqueous solution containing an equal amount of acetic acid and sodium acetate is 4.74. Furthermore, the pH of these liquids is less than seven. These solutions are made up of a weak acid and its salt.
- Basic Buffers- A weak base and its salt are equimolar with a strong acid in a basic buffer solution. These buffer solutions are employed to keep basic conditions. To generate a basic buffer with a basic pH, a weak base and its salt are combined with a strong acid. An aqueous solution of equal parts of ammonium hydroxide and ammonium chloride has a pH of 9.25. These solutions have a pH of greater than seven. They contain a weak base and a weak base salt.
Preparation of Buffer Solution: A buffer solution can be made by controlling the salt acid or salt base ratio if the dissociation constants of the acid (pKa) and the base (pKb) are known. Weak bases and their conjugate acids, or weak acids and their respective conjugate bases, are used to make these solutions. The Handerson-Hasselbalch equation and the preparation of acidic buffer and basic buffer.
Preparation of Acid Buffer
In an acid buffer solution with a strong base (KOH), consider a weak acid (HA) and its salt (KA). The weak acid (HA) ionises, and the equilibrium is as follows:
H2O + HA ⇋ H+ + A–
Acid dissociation constant is,
Ka = ([H+] [A–])/[HA]
The RHS and LHS take negative logs:
-log Ka = -log [H+] – log ([A–]/[HA])
pKa = pH – log ([salt]/[acid])
pH = pKa + log ([salt]/[acid])
pH of acid buffer solution = pKa + ([salt]/[acid])
The Henderson-Hasselbalch equation is sometimes known as the Henderson equation.
Preparation of Base Buffer
Consider a basic buffer solution with strong acid, having salt (BA) and a weak base (B).
As a result, the basic buffer solution will be,
pOH = pKb + log ([salt]/[base])
pOH of a basic buffer solution = pKb + log ([salt]/[acid])
pH of a basic buffer solution = pKa – log ([salt]/[acid])
Maintenance of pH
Consider the case of a buffer solution containing sodium acetate and acetic acid to better understand how buffer solutions maintain a steady pH. It’s worth noting that the sodium acetate is virtually totally ionized in this example, but the acetic acid is just faintly ionized. The following are the equilibrium reactions:
CH3COOH ⇌ H+ + CH3COO–
CH3COONa ⇌ Na+ + CH3COO–
When strong acids are supplied, the H+ ions combine with the CH3COO– ions to form a weakly ionized acetic acid, causing the pH of the surroundings to shift insignificantly. When very alkaline substances are added to this buffer solution, hydroxide ions react with the free acids in the solution to produce water molecules, as seen in the reaction below.
CH3COOH + OH– ⇌ CH3COO– + H2O
As a result, the hydroxide ions combine with the acid to create water, keeping the pH constant.
Uses of Buffer Solutions
- Buffer solutions are referred to by several different names, including pH buffers and hydrogen ion buffers.
- Many organisms employ buffer solutions to maintain an optimum pH for enzyme activity.
- The use of bicarbonate and carbonic acid buffer system to manage the pH of animal blood is an example of a buffer used in pH regulation.
- Enzyme function may be hindered, features may be lost, or the enzymes may even denature if certain buffers are not present. The enzymes’ catalytic function can be completely deactivated by this denaturation process.
Question 1: What is the Henderson-Hasselbalch equation’s limitation?
The Henderson – Hasselbalch equation isn’t applicable to strong acids and strong bases.
Question 2: What is buffering capacity?
The buffer capacity is the number of millimoles of acid or base that must be given to a liter of buffer solution to adjust the pH by one unit.
Question 3: What is a buffer action?
The ability of a buffer solution to withstand changes in pH when a little amount of acid or base is added is referred to as buffer action.
Question 4: Calculate the pH of a buffer solution containing 0.2 moles of NH4Cl per liter and 0.1 moles of NH4OH. NH4OH has a Kb of 1.85×10-5.
According to the Henderson’s Equation,
pOH = pKb + log ([salt]/[acid]) ……… (1)
Now pKb = -log(Kb) = -log(1.85×10-5) = 4.733
Since the buffer solution contains 0.2 moles of NH4Cl per litre, i.e., 0.2 moles of salt and 0.1 moles of NH4OH per litre, i.e. 0.1 moles of base. So, on substituting all the values in equation (1), we get
pOH = 4.7333+ log ((0.2)/(0.1))
= 4.733 + 0.301
pH = 14 – pOH
pH = 14 – 5.034
pH = 8.966
Question 5: What are the benefits of buffer solutions?
When an acidic or basic component is added, a buffer is a solution that can tolerate the pH change. It has the ability to neutralise small amounts of additional acid or base, allowing the pH of the solution to remain stable. This is critical for processes and reactions that require specific and stable pH ranges.
Question 6: What is the Handerson-Hasselbalch Equation’s significance?
The Handerson-Hasselbalch equation can be used to
- Determine the pH of the buffer made from salt and a weak acid/base mixture.
- Determine the pKa value.
- Produce a buffer solution with the required pH.