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 ↓


Hess's law of constant heat summation


Hess's law of constant heat addition, commonly called Hess's law, is an important principle studied in chemistry, particularly in the field of thermochemistry. It deals with the heat involved in chemical reactions and allows chemists to determine the enthalpy changes of reactions even when they cannot be measured directly. The law is named after Russian chemist Germain Hess, who formulated it in 1840.

Understanding Hess's law

To begin understanding Hess's law, it's important to know that in chemistry, a chemical reaction often involves the breaking and forming of bonds, resulting in a change in energy. This energy change is often measured as heat or enthalpy change, represented as ΔH. Hess's law states that the total enthalpy change of a chemical reaction is the same no matter how it is carried out, whether directly in one step or indirectly in multiple steps.

Hess's law can be expressed as follows:

ΔH_total = ΔH_1 + ΔH_2 + ... + ΔH_n

Here, ΔH_total is the total enthalpy change for the reaction, and ΔH_1, ΔH_2, ..., ΔH_n are the enthalpy changes for each step in the reaction pathway.

Thermodynamics and enthalpy

Thermodynamics is the branch of physics that deals with heat and temperature, and their relation to energy and work. In the context of chemical reactions, enthalpy (H) is a measure of the total energy of a thermodynamic system, including internal energy and the energy needed to displace the environment to make way for the system. The change in enthalpy (ΔH) is what we are primarily interested in, since it indicates the heat absorbed or released under conditions of constant pressure.

Importance of Hess's law

Hess's law is important for several reasons:

  • Predicting enthalpy change: This allows us to predict the total enthalpy change for a chemical reaction, even if it is difficult to measure directly.
  • Path Independence: This law states that the enthalpy change for a reaction is path-independent and depends only on the initial and final states.
  • Thermochemical calculations: It helps significantly in thermochemical calculations where direct measurement is impractical.

Visual example of Hess's law

Let's look at an example of the formation of water (H₂O) from hydrogen (H₂) and oxygen (O₂).

Step 1: H₂(g) → 2H(g) ΔH₁
Step 2: O₂(g) → 2O(g) ΔH₂
Step 3: 2H(g) + O(g) → H₂O(g) ΔH₃

Overall reaction: H₂(g) + 1/2 O₂(g) → H₂O(g) ΔH_total
    
        
            
            
            
            H₂(g) + 1/2 O₂(g)
            ΔH₁, ΔH₂
            
            2H(g) + O(g)
            ΔH₃
            
            H₂O(g)

The overall enthalpy change for water formation can be calculated by adding the enthalpy changes of these steps:

ΔH_total = ΔH₁ + ΔH₂ + ΔH₃

Text example using Hess's law

Suppose we want to determine the enthalpy change when graphite is burned to form carbon dioxide:

Step 1 (known): C(graphite) + O₂(g) → CO₂(g) ΔH₁ = -393.5 kJ/mol
Step 2 (hypothetical): CO₂(g) → C(graphite) + O₂(g) ΔH₂ = +393.5 kJ/mol
Overall reaction: C(graphite) + O₂(g) → CO₂(g) ΔH_total

Using Hess's law, the enthalpy change of the direct reaction would be the same as in Step 1, but using a hypothetical inverse step shows how Hess's law applies:

ΔH_total = ΔH₁ + ΔH₂
ΔH_total = -393.5 + 393.5 = 0 (for imaginary inverse)
ΔH_total = -393.5 kJ/mol (for the actual reaction, Step 1)

This shows how Hess's law ensures consistency in measuring the enthalpy change, no matter which route is taken.

Application of Hess's law in chemistry

Hess's law finds applications in various fields of chemistry:

  • Determination of enthalpy of formation: It helps in calculating the enthalpy of formation of compounds using data obtained from known reactions.
  • Understanding reaction pathways: It is used to analyze reaction pathways and intermediate steps in complex reactions.
  • Design of chemical processes: Industrial chemists use it to design energy-efficient processes.

Conclusion

Hess's law of constant heat summation is a fundamental principle in thermochemistry that underlies the conservation of energy. By showing that the enthalpy change is independent of the path taken by a reaction, it allows chemists to calculate the heat of reaction in a systematic and predictable way. This law is not simply a theoretical construct; it is a widely used tool in research and industry to solve practical problems involving chemical energy transformations.

Through an understanding of this law, students and chemists can appreciate the beauty and usefulness of thermodynamics concepts in predicting and measuring energy changes associated with chemical reactions.


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Hess's law of constant heat summation

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