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 → Thermodynamics ↓


Enthalpy of solution and neutralization


Chemistry involves many processes where energy is either absorbed or released. Two important concepts in thermodynamics related to this are enthalpy of solution and enthalpy of neutralization. These concepts are part of a field called calorimetry, which is the study of measuring heat changes resulting from chemical reactions.

What is enthalpy?

Enthalpy is a measure of the total energy of a thermodynamic system. It includes internal energy, which is the energy needed to create the system, and pressure-volume energy, which is the energy needed to make space for the system by displacing its environment. In simple terms, enthalpy is the heat content of a system.

In chemistry, we often express the change in enthalpy with the symbol ΔH. It represents the heat absorbed or released at constant pressure.

Enthalpy of solution

The enthalpy of solution, also known as the heat of solution, is the change in enthalpy caused by dissolving one mole of solute in a solvent. This process can either absorb heat from the surroundings (endothermic) or release heat to the surroundings (exothermic).

Let us consider a common example—dissolving salt in water. When table salt (NaCl) is dissolved in water, the interaction between the Na+ and Cl- ions in the salt and the water molecules releases energy due to the formation of ion-dipole interactions. However, energy is also required to break the ionic bonds in the salt. The overall change in energy determines whether the process is endothermic or exothermic.

Visualizing the enthalpy of solution

To visualize this process, consider the following simple representation of the interaction:

solute Solvent ΔH

Example calculation

If dissolving 1 mol of salt in water leads to an increase in temperature, this means that the process is exothermic and ΔH is negative. Suppose that dissolving 1 mol of a substance raises the temperature by 5 °C, and the specific heat capacity of the solution is known, then you can calculate ΔH using the formula:

        ΔH = -m × c × ΔT

where m is the mass of the solution, c is the specific heat capacity, and ΔT is the change in temperature.

Neutralization enthalpy

The enthalpy of neutralization is the change in enthalpy when the acid and base combine to form one mole of water. This is usually an exothermic reaction, which means that heat is released. The general reaction looks like this:

        HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

During neutralization, hydrogen ions H+ from the acid react with hydroxide ions OH- the base to form water. Energy is released by the formation of water molecules.

Visualization of enthalpy of neutralization

Here is a simplified visualization of the process:

Acid Base + heat

Example calculation

To calculate the enthalpy change during the neutralization reaction, you need to measure the heat change in the calorimeter using the following formula:

        q = -m × c × ΔT

where q is the heat absorbed or released, m is the mass of the solution, c is the specific heat capacity, and ΔT is the change in temperature from the surroundings. In practice, the enthalpy of neutralization is usually given in kilojoules per mole of water formed (kJ/mol).

Factors affecting enthalpy change

Several factors can affect the measured enthalpy change of a reaction:

  • Concentration: More concentrated solutions can have a greater change in enthalpy due to increased ion interactions.
  • Temperature: Reaction rates and energy changes can vary with temperature.
  • Nature of the solvent: Different solvents have different interactions with the solute, which affect the energy change.

Summary

The concepts of enthalpy of solution and neutralization are important for understanding energy changes in chemical reactions. Enthalpy of solution involves solute-solvent interactions and can be endothermic or exothermic depending on the substances involved. Enthalpy of neutralization helps in understanding the energy released during the reaction of acids and bases, which are primarily exothermic. Both concepts are not only fundamental in chemistry but also have important implications in industrial applications where energy efficiency and heat management are important.


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