Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Balance → Acid and Base Theory ↓


Bronsted-Lowry concept


The Bronsted-Lowry concept of acids and bases is an important model for understanding chemical reactions and equilibrium in chemistry. Developed by Danish chemist Johannes Nicolaus Bronsted and English chemist Thomas Martin Lowry in 1923, this theory focuses on the movement of protons (H + ions) between substances.

Basic definitions

According to the Bronsted-Lowry concept:

  • Acid: A substance that donates protons (H + ions) to another substance.
  • Base: A substance that accepts protons (H + ions) from another substance.

The ability to donate or accept a proton defines the acidic or basic nature of a compound in the Bronsted-Lowry theory. This concept broadens the Arrhenius theory and provides a more general approach that applies to a wide range of chemical reactions, including reactions in non-aqueous solutions.

Acid-base balance

An important aspect of the Bronsted-Lowry theory is the concept of equilibrium in acid-base reactions. An acid-base reaction can be represented as:

HA + B ⇌ A- + HB+

In this reaction:

  • HA is the acid that donates a proton to base B, forming the conjugate base A-.
  • B accepts the proton and forms the conjugate acid HB+.
  • This reaction is reversible, whereby a state of equilibrium is achieved.

Conjugate acid-base pair

Every acid has a conjugate base, and every base has a conjugate acid. This is due to the reversible nature of proton transfer reactions. The conjugate base is the species that remains after the acid donates its proton, and the conjugate acid is the species that forms when the base accepts a proton.

For example, in the reaction:

NH4+ + OH- ⇌ NH3 + H2O

Here, NH4+ is the acid and OH- is the base. After the reaction:

  • NH3 (ammonia) is the conjugate base of NH4+.
  • H2O (water) is the conjugate acid of the base OH-.

Visualization of acid-base reactions

Let's represent the acid-base reaction graphically:

HA B A- HB+ Equilibrium

In this illustration, the solid blue line shows the forward reaction where HA donates a proton to B to form A- and HB+. The dashed line shows the reverse reaction, illustrating the dynamic nature of the equilibrium.

Water: A special case

In the Bronsted-Lowry framework, water is an amphoteric substance. Amphoteric substances can act as either acids or bases, depending on the reaction conditions. Here are some examples:

HCl + H2O ⇌ Cl- + H3O+

In the first example, water acts as a base, accepting a proton from hydrochloric acid (HCl) to form a hydronium ion (H3O+).

NH3 + H2O ⇌ NH4+ + OH-

In the second example, water acts as an acid, donating a proton to ammonia (NH3) to form a hydroxide ion (OH-).

This ability to behave as both an acid and a base makes water a versatile medium for acid–base reactions in aqueous solution.

Strength of acids and bases

The strength of acids and bases is determined by their tendency to donate or accept protons:

  • Strong acid: Dissociates completely in the reaction, readily donating a proton. For example, hydrochloric acid (HCl).
  • Weak acid: Partially dissociates, not all molecules donate protons. For example, acetic acid (CH3COOH).
  • Strong base: Completely accepts protons in the reaction. For example, sodium hydroxide (NaOH).
  • Weak base: Partially accepts protons, such as ammonia (NH3).

The strength of the acid and base affects the position of equilibrium in an acid-base reaction.

Examples of Bronsted-Lowry reactions

Example 1: Ammonia and water

NH3 + H2O ⇌ NH4+ + OH-

In this reaction:

  • NH3 is a base and H2O initially acts as an acid.
  • The conjugate acid is NH4+ and the conjugate base is OH-.

Example 2: Acetic acid and water

CH3COOH + H2O ⇌ CH3COO- + H3O+

In this reaction:

  • CH3COOH (acetic acid) is an acid and H2O is initially a base.
  • The conjugate base is CH3COO- and the conjugate acid is H3O+.

Role of solvents

Bronsted-Lowry reactions can occur in a variety of solvents, not just water. Solvents can affect the equilibrium by stabilizing or destabilizing acids and bases and their conjugates. The ability of the solvent to donate or accept protons affects the reaction kinetics, providing an additional level of complexity and flexibility in predicting reaction outcomes.

Advantages of the Bronsted-Lowry concept

The Bronsted-Lowry concept has several advantages over other theories:

  • More general: It applies to a wider range of reactions than the Arrhenius theory, which is limited to aqueous reactions.
  • Focusing on proton transfer: By focusing on proton transfer, it provides a clear understanding of the mechanisms involved in acid-base reactions.
  • Conjugate pair: The concept of conjugate acid-base pairs provides insight into the reversibility and equilibrium of reactions.

Conclusion

The Bronsted-Lowry theory is one of the fundamental concepts in chemistry for understanding acids, bases, and their reactions. By emphasizing proton transfer and equilibrium, it provides a more comprehensive understanding of chemical behavior in diverse environments beyond aqueous solutions, which aligns with broader chemical principles and practices.


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Bronsted-Lowry concept

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