Welcome to LearnChemistry by Inam Jazbi.

In today’s lesson, we break down one of the most important topics of physical chemistry—Buffers or Buffer Solutions. Students often get confused between acidic buffers, basic buffers, their formulas, and how they actually resist pH change.

So in this blog, I have explained everything in the simplest and most scoring way, exactly how examiners expect in MDCAT, FSC, Cambridge O Levels, and board exams.
You’ll learn:
✔ What are Buffers?
✔ Types of Buffers
✔ How Acidic & Basic Buffers work
✔ Examples used in the human body
✔ Formulas that students mostly forget
✔ Smart tricks to solve buffer numericals

If you really want chemistry to feel easy and logical, this is the perfect guide for you—by Inam Jazbi (LearnChemistryby Inam Jazbi).

Buffers or Buffer Solutions

Definition of Buffers and Buffer Action

It is a solution that resists changes in its pH as a result of dilution or small addition of acids or bases.

A buffer solution is one whose pH is not changed significantly on dilution or even by the addition of the small amount of acid or base at constant temperature. A buffer solution or merely buffer is a solution or mixture of a weak acid and its salt with strong base i.e. its conjugate base (CH₃COOH + CH₃COONa) or a weak base and its salt strong acid i.e. its conjugate acid (NH₄OH + NH₄Cl) that has the ability to resist a change in its pH upon the addition of small amounts of either an acid or a base or on dilution. Buffer is also defined as the solution of reserve acidity or alkalinity which resists change of pH upon the addition of a small amount of acid or alkali.

In other words

a buffer solution contains a weak acid-base conjugate pair (i.e. a weak acid and its conjugate base or a weak base and its conjugate acid) in approximately equal concentrations that maintains a fairly constant H⁺ ions concentrations or pH despite small additions of acid or base.

Thus a buffer solution contains two species, an acidic (anionic) specie to neutralize OH^⁻ ions of added base and a basic specie to neutralize H⁺ ions of added acid. (It is necessary that the acid and the base components of the buffer must not consume each other through a neutralization reaction). Buffer is just the application of common ion effect.

For example, the pH of pure water is 7.0. When an acid or base gets added to water, it turns into an acidic or a basic solution. However, water + buffer solutions have very little change in their pH upon the addition of acid or base.

Reason of Buffer Action

An equilibrium is established between the weak acid and conjugate base of a buffer solution.

The weak acid and conjugate base are in excess, meaning that the position of the equilibrium established between them will not be sensitive to changes in their concentrations, but it will be very sensitive to changes in the concentration of H⁺_(aq) ions.

When H⁺_(aq) ion concentration is increased or decreased, the position of equilibrium moves to oppose the change.

Buffer Action

The resistance offered by a buffer solution to change its pH on adding small amount of acid or base is called Buffer Action.

Buffer Range

It is the range of pH over which a buffer solution remains effective on addition of strong acid or base.

Buffer Capacity

A buffer solution cannot keep the pH constant if enough acid or base is added into it. The quantity of acid or base added to buffer solution before changing its pH is called buffer capacity. Buffer capacity is a measure of how little the pH changes with the addition of a given amount of acid or base. Buffer capacity is maximum when acid to salt ratio or base to salt ratio is equal to 1.

The buffering capacity or buffering ability i.e. the effectiveness of the buffer solution is the amount of acid or base the buffer can neutralize without a significant changes in pH. The buffer capacity of a solution is the capability of a buffer to resist the changes of pH i.e. it measures that how much extra acid or base, the solution can absorb (neutralize) before the buffer is essentially destroyed or consumed.

The buffer capacity depends upon the concentrations of the components of the buffer mixture. (i.e. amount of acid and base from which the buffer is made). The larger the amount, the greater the buffering capacity.

Why does temperature decrease buffer capacity?

Temperature is directly proportional to the dissociation of ions. So, the higher the temperature, the more the dissociation of ions. Hence, the numbers of undissociated ions become less resulting in decreased buffer capacity.

Types of Buffers

1. Acidic buffers with acidic pH values (4-7), contains weak acid and its conjugate base

2. Basic buffer with basic pH values (7-11), contains weak base and its conjugate acid

3. Simple Buffer; contains acid salt and normal salt of a polybasic acid e.g. Na₂HPO₄ + Na₃PO₄

Examples of Buffer Mixtures or Solution

A buffer mixture or solution contains two species, one capable of reacting with H+ ions and the other with OH¯ ions. A buffer solution or mixture is prepared by mixing a weak acid and its salt with strong base (acidic buffer solution) or a weak base and its salt with strong acid (basic buffer solution).

(i) A weak acid and a salt of it with strong base e.g. CH₃COOH + CH₃COONa (& Boric acid + Borax) OR

(ii) A weak base and a salt of it with strong acid e.g. NH₄OH + NH₄Cl.

Characteristics of Buffer

(1). A buffer solution has a specific, constant pH value.

(2). Its pH value is very slightly changed upon the small additions of strong acid or strong base. (The buffering capacity of any buffer is exhausted if enough strong acid or strong base is added that will use up original amount of weak base or weak acid).

(3). Its pH does not alter either on keeping for a long time or on dilution. Thus the pH of a buffer solution does not depend on the volume of the solution because a change in solution volume changes the concentrations of the acid base components of buffer by the same amount and the pH remain unchanged. So a buffer solution can be diluted without a change in pH.

(4). The pH of a buffer solution depends only on pKₐ and on the relative molar amounts of weak acid (or base) and conjugate base (or conjugate acid).

Acid-Base Buffer Systems in the body’s fluids

Acid-base buffer systems consist of chemicals that combine with excess acids or bases. Buffer system chemicals can combine with strong acids, which release more hydrogen ions, to convert them into weak acids, which release fewer hydrogen ions. The three most important acid-base buffer systems in the body’s fluids are as follows:

Bicarbonate buffer system:

This system is present in both intracellular and extracellular fluids, using the bicarbonate ion as a weak base and carbonic acid as a weak acid. It is sometimes called the carbonic acid bicarbonate buffer system. Carbonic acid is formed when hydrogen ions are excessive and dissociates when conditions are basic or alkaline. The body maintains a readily available bicarbonate reserve.

Phosphate buffer system:

This system also operates in both intracellular and extracellular fluids and is very important in controlling hydrogen ion concentrations in the fluid of the nephrons and in urine. It consists of monohydrogen phosphate and dihydrogen phosphate.

Protein buffer system:

Consists of plasma proteins and certain cell proteins. When the solution pH falls, amino groups accept hydrogen ions; when it rises, carboxyl groups release hydrogen ions. For red blood cells, which are densely packed with hemoglobin molecules that buffer hydrogen ions, a chloride shift occurs. This involves dissociation of carbonic acid and bicarbonate ions diffusing into the plasma in exchange for chloride ions. In the lungs this occurs in reverse and is known as the hemoglobin buffer system.

Applications of Buffers

1. Maintenance of pH of Blood

2. Maintenance of pH in Laboratory Reactions

3. Maintenance of pH in Various Industries

4. Maintaining the pH of culture media for the growth of bacteria in Pathological Laboratories

5. For Good Yield of Crops

6. For Preservation of foods and Fruits

7. Use in and Analytical chemistry and Other Sciences

some real-life applications of buffer solutions

⇒ Food preservation chemicals

⇒ Photographic materials

⇒ Leathers

⇒ Dyes

⇒ Calibration of pH meters

⇒ Maintaining the pH of culture media for the growth of bacteria

⇒ The pH of soil, etc.

buffer solutions in daily life

Bicarbonates and phosphate buffer solutions are mostly used in daily life. Although, if we talk about living beings, hemoglobin and proteins are buffers too.

Applications of Buffers (Details)

1. Buffers solutions play important roles in maintaining the pH in biological and physiological processes. The maintenance of blood pH is necessary for life. Buffers help to maintain pH of blood at remarkably constant level 7.3 to 7.4 due to bicarbonate and carbonic acid buffer.

Organisms have built-in buffers to protect them against changes in pH. Human blood has a pH near 7.4 that is maintained by combination of carbonate (H₂CO₃/NaHCO₃), phosphate(Na₂HPO₄ /NaH₂PO₄) and protein buffer systems.

2. In laboratory many inorganic and organic reactions are performed in buffered solution to minimize any adverse effect caused by acids or bases that might be consumed or produced during reaction.

3. They are widely used in many industrial processes to maintain pH during fermentation, dye processes, manufacturing of pharmaceuticals, tanning of leather, manufacture of sugar, paper, drugs etc.

4. They are used in preparing culture media in biological and pathological laboratories to maintain a constant pH for suitable growth of bacteria and viruses.

5. They are used in agriculture to maintain the pH of soil for suitable crop yield.

6. They are used in foods industry to maintain the pH of various food items for preservation of their flavor, appearance and micro-biological stability.

7. Buffers are important in many areas of analytical chemistry (and allied sciences like molecular biology, microbiology, cell biology, soil sciences, nutrition and clinical analyses).

Mechanism of Buffer Action (of Acetic acid-acetate buffer i.e. acidic buffer)

The buffering action (i.e. the resistance offered by a buffer solution to change its pH on the addition of an acid or a base) can be explained on the basis of Le-Chatelier’s Principle and common ion effect.

To explain buffer action, consider the ionization of an acidic buffer solution consisting of weak acetic acid, CH₃COOH (weak electrolyte) and sodium acetate, CH₃COONa (Salt of weak acid with strong base; a strong electrolyte).

(i) Maintenance of pH of Buffer

Both acetic acid and sodium acetate provide CH₃COO^⁻ ions, however CH₃COO^⁻ ion comes from strong basic salt CH₃COONa are high in concentration.

Acetic acid being a weak acid (or electrolyte) undergoes very little dissociation and feebly dissociates to give few acetate ions and H⁺ ions. Sodium acetate being a strong electrolyte dissociates fully to give common acetate ions (which may hydrolyze to give acetic acid). The increased concentration of acetate ions (common ion) suppresses the ionization of acetic acid through the common ion effect of acetate ions by shifting acid equilibrium to left thereby reducing the H⁺ ions concentration forming undissociated acetic acid. (The presence of acetic acid suppresses the hydrolysis of acetate ions). Thus there is no appreciable change in the pH of buffer solution.

CH₃COOH_(aq) ⇌ CH₃COO⁻ _(aq) + H⁺_(aq)

CH₃COONa_(aq) ⇌ CH₃COO⁻_(aq) + Na⁺_(aq) Acetate ion (Common ion)

If one goes on adding sodium acetate in acetic acid solution, then the added concentrations of CH₃COO¯ decrease the dissociation of acetic acid and the pH of solution increases. It is obvious that greater the concentration of acetic acid as compared to CH₃COONa, lesser is the pH of buffer solution. The following table shows how the pH value of a buffer mixture of two compounds is maintained.

(ii) Effect of addition of a small amount of strong acid

When a small amount of strong acid (such as HCl) is added to buffer solution, the additional H⁺ ions (given by acid) will be consumed by the base component (conjugate base) i.e. acetate ions to give unionized acetic acid and hence the pH of solution will remain at original level.

HCl_(aq) ⇌ H⁺_(aq) + Cl¯_(aq)

H⁺_(aq) + CH₃COO¯_(aq) ⇌ CH₃COOH_(aq)

(iii) Effect of addition of a small amount of strong base

When small amount of strong base (e.g. NaOH) is introduced in buffer solution, the additional OH^⁻ ions given by base will be neutralized by acid component (weak acid) of the buffer to form water and acetate ions shifting the equilibrium to the right to produce more H⁺ ions till practically all the excess OH^⁻ ions are neutralized and the original buffer pH is restored and hence the pH of solution practically remains constant.

NaOH_(aq) ⇌ Na⁺_(aq) + OH¯_(aq)

OH¯_(aq)+ CH₃COOH_(aq) ⇌ H₂O_(l) + CH₃COO¯_(aq)

H⁺_(aq) + OH¯_(aq) ® H₂O_(l)

Conclusion

Thus, the buffer solution resists any change in its pH value by the addition of small amount of acid or base.

✔ Correct Answer: C

⭐ 1. Types of Buffers (Colourful Table)

🟩 Type	🟦 Components	🟧 Example
Acidic Buffer	Weak acid + salt with strong base	CH₃COOH + CH₃COONa
Basic Buffer	Weak base + salt with strong acid	NH₄OH + NH₄Cl

⭐ 2. Henderson–Hasselbalch Formulas

🟦 For Acidic Buffer

pH = pKa + log (Salt / Acid)

🟩 For Basic Buffer

pOH = pKb + log (Salt / Base)

⭐ 3. Common Buffer Pairs (Beautiful Quick Table)

🟪 Buffer Type	🟨 Pair	🟧 Use
Acidic Buffer	CH₃COOH/CH₃COONa	Lab buffer
Basic Buffer	NH₄OH/NH₄Cl	Titrations
Blood Buffer	H₂CO₃/HCO₃⁻	Maintains pH 7.4
Phosphate Buffer	H₂PO₄⁻/HPO₄²⁻	Cells + kidneys

⭐ 4. Glycine Buffer

🟩 Acid Form	🟦 Base Form	🟧 Type
NH₃⁺–CH₂–COOH (Glycine HCl)	NH₂–CH₂–COOH (Glycine)	Acidic Buffer

Conclusion

Buffers are everywhere — inside our blood, medicines, biological systems, and laboratory solutions. Understanding buffer solutions not only makes chemistry easier, but also helps you solve numerical problems confidently.

I hope this colourful guide by Inam Jazbi | LearnChemistry helped you master this chapter in the simplest way.

Buffer Solutions: Types, Formulas, Examples, and Easy Memory Tricks

MDCAT-STYLE MCQs

🌟 MCQ 1 (Concept)

🌟 MCQ 2 (Definition)

🌟 MCQ 3 (Acidic Buffer)

🌟 MCQ 4 (Basic Buffer)

🌟 MCQ 5 (Function)

🌟 MCQ 6 (Exam-Favorite)

🌟 MCQ 7

🌟 MCQ 8

🌟 MCQ 9

🌟 MCQ 10

🌟 MCQ 11

🌟 MCQ 12

🌟 MCQ 13

🌟 MCQ 14

🌟 MCQ 15

🌟 MCQ 16

🌟 MCQ 17

🌟 MCQ 18

🌟 MCQ 19

🌟 MCQ 20