Acid Base Buffer Solution-Definition, Types, Formula, Examples, and FAQs

Buffer Solution is a special aqueous solution that resists the change in its pH when some quantity of acid and Base is added.

Many fluids, such as blood, have specific pH values of 7.14, and variations in these values indicate that the body is malfunctioning.  

The change in pH of Buffer Solutions on adding a small quantity of acid or bases is very minimal and hence they are used to make solutions that resist the change in pH.

Let us learn about Buffer solution, its types, and others in this article.

Buffer Solution Definition

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 Solutions

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. An 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.

Buffer Solution is a special aqueous solution that resists the change in its pH when some quantity of acid and Base is added.

Basic or Alkaline 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.

Buffer Solution is a special aqueous solution that resists the change in its pH when some quantity of acid and Base is added.

Also Read: Acids, Bases, and Salts

Mechanism of Buffering Action

To understand how a buffer works, think about the example of a buffer solution created by combining sodium acetate and acetic acid. Acetate acid, as implied by its name, is an acid with the chemical formula CH3COOH, whereas sodium acetate dissociates in solution to produce the conjugate base CH3COO acetate ions. The reaction equation is:

CH3COOH (aq) + OH(aq)    ⇢    CH3COO (aq) + H2O (aq)

Now, the acetate ion can be neutralized by adding a strong acid to this solution as, 

CH3COO (aq) + H+ (aq)    ⇢    CH3COOH (aq)

Thus, the original buffer reaction equilibrium changes, and hence, the pH remains constant. 

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.

Henderson-Hasselbalch Equation

Henderson-Hasselbalch equation is the equation that provides a relationship between the pH of acids and their pKa.

The Henderson-Hasselbalch equation is,

pH = pKa + log10 ([A] / [HA])

where 
[A] denotes the molar concentration of the Conjugate Base (of the Acid)
[HA] denotes the molar concentration of the Weak Acid

Importance of Henderson Equation

The significance of the Henderson Equation is,

  • It helps to determine the pH of the buffer obtained from a mixture of salt and weak acid or base.
  • To determine the pKa value.
  • To maintain the buffer solution of the required pH.

Limitations of Henderson-Hasselbalch Equation

The Henderson-Hasselbalch Equation is not applicable in many cases are:

  • This equation does not explain the water self-dissociation tendency. 
  • The assumptions stated by this equation are not applicable when strong acids or bases are associated.

A buffer solution can easily be prepared using Henderson-Hasselbalch Equation

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) ionizes, and the equilibrium is as follows:

H2O + HA ⇋ H+ + A

The acid dissociation constant is,

Ka = ([H+] [A])/[HA]

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])

Preparation of Base Buffer

Consider a basic buffer solution with strong acid, 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])

Buffering Capacity

Buffering Capacity is the property of solutions which is represented by the symbol β. It is defined as the ratio of the moles of acid or base required to change the pH of the solution by 1 and the pH change and the volume of buffers.

Β = millimoles /(ΔpH)

Uses of Buffer Solutions

Various uses of Buffer solution are,

  • 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.

pH Maintenance

How the pH of any solution is kept constant can easily be explained with the help of the example given below,

Take formic acid and sodium formate ions in an aqueous solution they behave as buffer solution and their equilibrium reaction is,

  • HCOOH ⇌ H+ + HCOO
  • HCOONa ⇌ Na+ + HCOO

If the strong acids are added to the above buffer solution, the hydrogen ion (H+ ions) combine with the HCOO ions to give a weakly ionized formic acid, and thus, the pH of the solution is kept constant or slightly changed.

If the strong base is added to the above buffer solution, the hydroxide ion (OH ions) combine with the H+ ions available in the solution to give water molecules, and thus, the pH of the solution is kept constant. The reaction occurring in the process is given below.

HCOOH + OH ⇌ HCOO + H2O

Thus, the pH of the buffer solution is kept constant.

Read More,

Solved Examples on Buffer Solution

Example 1: What is the pH of a buffered solution of 1.5 M NH3 and 2.5 M NH4Cl when 0.5 M HCl is added to the solution?

Solution:

We know that,

pKb of ammonia is 4.75

pKa = 14 – pKb. 

       =  14 – 4.75 = 9.25

Now, on adding 0.5 M HCl

0.5 M H+ ions are available in the aqueous solution which reacts with 0.5 M NH3 to form 0.5 M NH4Cl

Now the remaining concentration of ammonia is 1 M and that of NH4Cl is 3 M.

Using the Henderson-Hasselbalch equation,

pKa – log ([salt]/[acid]) = 9.25 – log (3/1) 

                                     = 9.25 – 0.477

                                     = 8.773

Example 2: How many moles of sodium formate and formic acid are required to prepare 1.00 L of a 0.25 mol/L buffer solution with pH 4.00? (given pKa of Sodium formate is 3.86)

Solution:

pH = pKa + log([A][HA])

Given,  pKa of Sodium formate is 3.86

6.00 = 3.86 + log([A][HA])

log ([A][HA]) = 4 – 3.86

log ([A][HA]) = 0.14

[A][HA] = 100.14

              = 1.38

[A⁻] = 1.38[HA]

Given, 

[A⁻] + [HA] = 0.25 mol/L

1.38[HA] + [HA] = 0.25 mol/L

2.38[HA] = 0.25 mol/L

[HA] = 0.25 mol/L / 2.38

[HA] = 0.105 mol / L

[A⁻] = (0.250 – 0.105) mol/L 

       = 0.145 mol/L

Thus, 0.105 mol of sodium formate and 0.145 mol of formic acid is required to form the required buffer solution.

FAQs on Buffer Solution

Question 1: What is a buffer solution?

Answer:

A solution which resists the small change in the pH by adding some acid and base is called a Buffer Solution.

Question 2: What is buffering capacity?

Answer:

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?

Answer:

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: What are the advantages of buffer solutions?

Answer:

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 5: What is the pH range of an Acidic Solution?

Answer:

The Acidic solution has a pH range of 0-6.

Question 6: What is the pH range of the Basic Solution?

Answer:

The Basic solution has a pH range of 8-14.

Question 7: What is the pH of the Neutral Solution?

Answer:

The Neutral solution has a pH of 7.

Er. Neeraj K.Anand is a freelance mentor and writer who specializes in Engineering & Science subjects. Neeraj Anand received a B.Tech degree in Electronics and Communication Engineering from N.I.T Warangal & M.Tech Post Graduation from IETE, New Delhi. He has over 30 years of teaching experience and serves as the Head of Department of ANAND CLASSES. He concentrated all his energy and experiences in academics and subsequently grew up as one of the best mentors in the country for students aspiring for success in competitive examinations. In parallel, he started a Technical Publication "ANAND TECHNICAL PUBLISHERS" in 2002 and Educational Newspaper "NATIONAL EDUCATION NEWS" in 2014 at Jalandhar. Now he is a Director of leading publication "ANAND TECHNICAL PUBLISHERS", "ANAND CLASSES" and "NATIONAL EDUCATION NEWS". He has published more than hundred books in the field of Physics, Mathematics, Computers and Information Technology. Besides this he has written many books to help students prepare for IIT-JEE and AIPMT entrance exams. He is an executive member of the IEEE (Institute of Electrical & Electronics Engineers. USA) and honorary member of many Indian scientific societies such as Institution of Electronics & Telecommunication Engineers, Aeronautical Society of India, Bioinformatics Institute of India, Institution of Engineers. He has got award from American Biographical Institute Board of International Research in the year 2005.

CBSE Class 11 Chemistry Syllabus

CBSE Class 11 Chemistry Syllabus is a vast which needs a clear understanding of the concepts and topics. Knowing CBSE Class 11 Chemistry syllabus helps students to understand the course structure of Chemistry.

Unit-wise CBSE Class 11 Syllabus for Chemistry

Below is a list of detailed information on each unit for Class 11 Students.

UNIT I – Some Basic Concepts of Chemistry

General Introduction: Importance and scope of Chemistry.

Nature of matter, laws of chemical combination, Dalton’s atomic theory: concept of elements,
atoms and molecules.

Atomic and molecular masses, mole concept and molar mass, percentage composition, empirical and molecular formula, chemical reactions, stoichiometry and calculations based on stoichiometry.

UNIT II – Structure of Atom

Discovery of Electron, Proton and Neutron, atomic number, isotopes and isobars. Thomson’s model and its limitations. Rutherford’s model and its limitations, Bohr’s model and its limitations, concept of shells and subshells, dual nature of matter and light, de Broglie’s relationship, Heisenberg uncertainty principle, concept of orbitals, quantum numbers, shapes of s, p and d orbitals, rules for filling electrons in orbitals – Aufbau principle, Pauli’s exclusion principle and Hund’s rule, electronic configuration of atoms, stability of half-filled and completely filled orbitals.

UNIT III – Classification of Elements and Periodicity in Properties

Significance of classification, brief history of the development of periodic table, modern periodic law and the present form of periodic table, periodic trends in properties of elements -atomic radii, ionic radii, inert gas radii, Ionization enthalpy, electron gain enthalpy, electronegativity, valency. Nomenclature of elements with atomic number greater than 100.

UNIT IV – Chemical Bonding and Molecular Structure

Valence electrons, ionic bond, covalent bond, bond parameters, Lewis structure, polar character of covalent bond, covalent character of ionic bond, valence bond theory, resonance, geometry of covalent molecules, VSEPR theory, concept of hybridization, involving s, p and d orbitals and shapes of some simple molecules, molecular orbital theory of homonuclear diatomic molecules(qualitative idea only), Hydrogen bond.

UNIT V – Chemical Thermodynamics

Concepts of System and types of systems, surroundings, work, heat, energy, extensive and intensive properties, state functions. First law of thermodynamics – internal energy and enthalpy, measurement of U and H, Hess’s law of constant heat summation, enthalpy of bond dissociation, combustion, formation, atomization, sublimation, phase transition, ionization, solution and dilution. Second law of Thermodynamics (brief introduction)
Introduction of entropy as a state function, Gibb’s energy change for spontaneous and nonspontaneous processes.
Third law of thermodynamics (brief introduction).

UNIT VI – Equilibrium

Equilibrium in physical and chemical processes, dynamic nature of equilibrium, law of mass action, equilibrium constant, factors affecting equilibrium – Le Chatelier’s principle, ionic equilibrium- ionization of acids and bases, strong and weak electrolytes, degree of ionization,
ionization of poly basic acids, acid strength, concept of pH, hydrolysis of salts (elementary idea), buffer solution, Henderson Equation, solubility product, common ion effect (with illustrative examples).

UNIT VII – Redox Reactions

Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, in terms of loss and gain of electrons and change in oxidation number, applications of redox reactions.

UNIT VIII – Organic Chemistry: Some basic Principles and Techniques

General introduction, classification and IUPAC nomenclature of organic compounds. Electronic displacements in a covalent bond: inductive effect, electromeric effect, resonance and hyper conjugation. Homolytic and heterolytic fission of a covalent bond: free radicals, carbocations, carbanions, electrophiles and nucleophiles, types of organic reactions.

UNIT IX – Hydrocarbons

Classification of Hydrocarbons
Aliphatic Hydrocarbons:
Alkanes – Nomenclature, isomerism, conformation (ethane only), physical properties, chemical reactions.
Alkenes – Nomenclature, structure of double bond (ethene), geometrical isomerism, physical properties, methods of preparation, chemical reactions: addition of hydrogen, halogen, water, hydrogen halides (Markovnikov’s addition and peroxide effect), ozonolysis, oxidation, mechanism of electrophilic addition.
Alkynes – Nomenclature, structure of triple bond (ethyne), physical properties, methods of preparation, chemical reactions: acidic character of alkynes, addition reaction of – hydrogen, halogens, hydrogen halides and water.

Aromatic Hydrocarbons:

Introduction, IUPAC nomenclature, benzene: resonance, aromaticity, chemical properties: mechanism of electrophilic substitution. Nitration, sulphonation, halogenation, Friedel Craft’s alkylation and acylation, directive influence of functional group in monosubstituted benzene. Carcinogenicity and toxicity.

To know the CBSE Syllabus for all the classes from 1 to 12, visit the Syllabus page of CBSE. Meanwhile, to get the Practical Syllabus of Class 11 Chemistry, read on to find out more about the syllabus and related information in this page.

CBSE Class 11 Chemistry Practical Syllabus with Marking Scheme

In Chemistry subject, practical also plays a vital role in improving their academic scores in the subject. The overall weightage of Chemistry practical mentioned in the CBSE Class 11 Chemistry syllabus is 30 marks. So, students must try their best to score well in practicals along with theory. It will help in increasing their overall academic score.

CBSE Class 11 Chemistry Practical Syllabus

The experiments will be conducted under the supervision of subject teacher. CBSE Chemistry Practicals is for 30 marks. This contribute to the overall practical marks for the subject.

The table below consists of evaluation scheme of practical exams.

Evaluation SchemeMarks
Volumetric Analysis08
Salt Analysis08
Content Based Experiment06
Project Work04
Class record and viva04
Total30

CBSE Syllabus for Class 11 Chemistry Practical

Micro-chemical methods are available for several of the practical experiments. Wherever possible such techniques should be used.

A. Basic Laboratory Techniques
1. Cutting glass tube and glass rod
2. Bending a glass tube
3. Drawing out a glass jet
4. Boring a cork

B. Characterization and Purification of Chemical Substances
1. Determination of melting point of an organic compound.
2. Determination of boiling point of an organic compound.
3. Crystallization of impure sample of any one of the following: Alum, Copper Sulphate, Benzoic Acid.

C. Experiments based on pH

1. Any one of the following experiments:

  • Determination of pH of some solutions obtained from fruit juices, solution of known and varied concentrations of acids, bases and salts using pH paper or universal indicator.
  • Comparing the pH of solutions of strong and weak acids of same concentration.
  • Study the pH change in the titration of a strong base using universal indicator.

2. Study the pH change by common-ion in case of weak acids and weak bases.

D. Chemical Equilibrium
One of the following experiments:

1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing the concentration of either of the ions.
2. Study the shift in equilibrium between [Co(H2O)6] 2+ and chloride ions by changing the concentration of either of the ions.

E. Quantitative Estimation
i. Using a mechanical balance/electronic balance.
ii. Preparation of standard solution of Oxalic acid.
iii. Determination of strength of a given solution of Sodium hydroxide by titrating it against standard solution of Oxalic acid.
iv. Preparation of standard solution of Sodium carbonate.
v. Determination of strength of a given solution of hydrochloric acid by titrating it against standard Sodium Carbonatesolution.

F. Qualitative Analysis
1) Determination of one anion and one cation in a given salt
Cations‐ Pb2+, Cu2+, As3+, Al3+, Fe3+, Mn2+, Ni2+, Zn2+, Co2+, Ca2+, Sr2+, Ba2+, Mg2+, NH4 +
Anions – (CO3)2‐ , S2‐, NO2 , SO32‐, SO2‐ , NO , Cl , Br, I‐, PO43‐ , C2O2‐ ,CH3COO
(Note: Insoluble salts excluded)

2) Detection of ‐ Nitrogen, Sulphur, Chlorine in organic compounds.

G) PROJECTS
Scientific investigations involving laboratory testing and collecting information from other sources.

A few suggested projects are as follows:

  • Checking the bacterial contamination in drinking water by testing sulphide ion
  • Study of the methods of purification of water.
  • Testing the hardness, presence of Iron, Fluoride, Chloride, etc., depending upon the regional
    variation in drinking water and study of causes of presence of these ions above permissible
    limit (if any).
  • Investigation of the foaming capacity of different washing soaps and the effect of addition of
    Sodium carbonate on it.
  • Study the acidity of different samples of tea leaves.
  • Determination of the rate of evaporation of different liquids Study the effect of acids and
    bases on the tensile strength of fibres.
  • Study of acidity of fruit and vegetable juices.

Note: Any other investigatory project, which involves about 10 periods of work, can be chosen with the approval of the teacher.

Practical Examination for Visually Impaired Students of Class 11

Below is a list of practicals for the visually impaired students.

A. List of apparatus for identification for assessment in practicals (All experiments)
Beaker, tripod stand, wire gauze, glass rod, funnel, filter paper, Bunsen burner, test tube, test tube stand,
dropper, test tube holder, ignition tube, china dish, tongs, standard flask, pipette, burette, conical flask, clamp
stand, dropper, wash bottle
• Odour detection in qualitative analysis
• Procedure/Setup of the apparatus

B. List of Experiments A. Characterization and Purification of Chemical Substances
1. Crystallization of an impure sample of any one of the following: copper sulphate, benzoic acid
B. Experiments based on pH
1. Determination of pH of some solutions obtained from fruit juices, solutions of known and varied
concentrations of acids, bases and salts using pH paper
2. Comparing the pH of solutions of strong and weak acids of same concentration.

C. Chemical Equilibrium
1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing
the concentration of eitherions.
2. Study the shift in equilibrium between [Co(H2O)6]2+ and chloride ions by changing the
concentration of either of the ions.

D. Quantitative estimation
1. Preparation of standard solution of oxalic acid.
2. Determination of molarity of a given solution of sodium hydroxide by titrating it against standard
solution of oxalic acid.

E. Qualitative Analysis
1. Determination of one anion and one cation in a given salt
2. Cations – NH+4
Anions – (CO3)2-, S2-, (SO3)2-, Cl-, CH3COO-
(Note: insoluble salts excluded)
3. Detection of Nitrogen in the given organic compound.
4. Detection of Halogen in the given organic compound.

Note: The above practicals may be carried out in an experiential manner rather than recording observations.

We hope students must have found this information on CBSE Syllabus useful for their studying Chemistry. Learn Maths & Science in interactive and fun loving ways with ANAND CLASSES (A School Of Competitions) App/Tablet.

Frequently Asked Questions on CBSE Class 11 Chemistry Syllabus

Q1

How many units are in the CBSE Class 11 Chemistry Syllabus?

There are 9 units in the CBSE Class 11 Chemistry Syllabus. Students can access various study materials for the chapters mentioned in this article for free at ANAND CLASSES (A School Of Competitions).

Q2

What is the total marks for practicals examination as per the CBSE Class 11 Chemistry Syllabus?

The total marks for the practicals as per the CBSE Class 11 Chemistry Syllabus is 30. It includes volumetric analysis, content-based experiment, salt analysis, class record, project work and viva.

Q3

Which chapter carries more weightage as per the CBSE Syllabus for Class 11 Chemistry?

The organic chemistry chapter carries more weightage as per the CBSE Syllabus for Class 11 Chemistry.