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 ↓


Arrhenius concept


The Arrhenius concept of acids and bases is one of the fundamental theories in chemistry that describes how acids and bases behave in aqueous solution. This theory was proposed by Swedish chemist Svante Arrhenius in 1887. His work laid the groundwork for understanding the nature of acids and bases. According to Arrhenius, there are special definitions for the structure of acids and bases, based on how they interact with water.

Understanding the Arrhenius definition of acids and bases

In the Arrhenius concept, an acid is defined as a substance that increases the concentration of hydrogen ions (H +) in an aqueous solution. When an acid dissolves in water, it donates H + ions to the solution.

A common example of an Arrhenius acid is hydrochloric acid (HCl). When HCl is dissolved in water, it dissociates to form hydrogen ions:

HCl (aq) → H + (aq) + Cl - (aq)

In contrast, an Arrhenius base is defined as a substance that increases the concentration of hydroxide ions (OH -) in an aqueous solution. When a base dissolves in water, it donates OH - ions to the solution.

A typical example of an Arrhenius base is sodium hydroxide (NaOH). When NaOH is dissolved in water, it decomposes as follows:

NaOH (aq) → Na + (aq) + OH - (aq)

The role of water

In the Arrhenius concept, water plays an important role as a solvent, providing a medium for acids and bases to dissociate and form ions. The presence of water allows acids to dissociate into H + ions and bases into OH- ions, which characterizes their acidic or basic nature.

Limitations of the Arrhenius concept

While the Arrhenius concept provides a basic understanding of acid and base behavior, it comes with some limitations. One major limitation is that it only applies to aqueous solutions. This means that acids and bases that do not dissociate or form ions in water cannot be effectively described by this theory.

Another limitation is that this concept does not take into account acid-base reactions that do not involve the formation of water. For example, the reaction between ammonia (NH 3) and hydrogen chloride gas (HCl) results in the formation of ammonium chloride (NH 4Cl), but this reaction occurs without water as a solvent:

NH 3 (g) + HCl (g) → NH 4Cl (s)

Concept of Arrhenius

To better understand the Arrhenius concept, we can imagine the dissociation of an acid such as hydrochloric acid in water:

HCl H + CL - Separation

Example: Comparison of different acids and bases

Now, let's look at some examples and understand how different acids and bases behave according to the Arrhenius definition:

Acid

  • Sulfuric acid (H 2SO 4): When sulfuric acid is dissolved in water, it dissociates to release two hydrogen ions: 
    H 2SO 4 (aq) → 2H + (aq) + SO 4 2- (aq)
  • Nitric acid (HNO 3): When nitric acid dissolves in water, it dissociates to yield hydrogen ions and nitrate ions: 
    HNO 3 (aq) → H + (aq) + NO 3 - (aq)
  • Acetic acid (CH 3COOH): Acetic acid, when dissolved in water, partially dissociates to release hydrogen ions: 
    CH 3COOH (aq) ⇌ H + (aq) + CH 3COO - (aq)

Bases

  • Potassium hydroxide (KOH): Potassium hydroxide dissociates in water to release hydroxide ions: 
    KOH (aq) → K + (aq) + OH - (aq)
  • Calcium hydroxide (Ca(OH) 2): Calcium hydroxide dissociates in water to form two hydroxide ions: 
    Ca(OH) 2 (aq) → Ca 2+ (aq) + 2OH - (aq)
  • Ammonium hydroxide (NH 4OH): Ammonium hydroxide is a weak base that dissociates slightly in water: 
    NH 4OH (aq) ⇌ NH 4 + (aq) + OH - (aq)

Conclusion

The Arrhenius concept provides a fundamental understanding of how acids and bases form by emphasizing their behavior in water. While this theory is limited to aqueous solutions and does not explain all acid-base reactions, it is an important stepping stone to the more comprehensive theories used by chemists today. This concept provides an accessible introduction to the dynamic world of acid-base chemistry, setting the stage for further discovery and study.


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Arrhenius concept

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