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 formation and combustion


In the study of thermodynamics, especially in chemistry, it is important to understand the concepts of enthalpy of formation and combustion. These concepts help us understand how energy is absorbed or released during chemical reactions. We will explore these concepts using simple language and many examples to make learning interesting and effective.

What is enthalpy?

Before diving into the enthalpy of formation and combustion, let's first understand what enthalpy is. Enthalpy, represented by the symbol H, is a measure of the energy in a thermodynamic system. It is the sum of the system's internal energy and the product of its pressure and volume. In other words:

H = U + PV

Where:

  • H is the enthalpy
  • U is the internal energy
  • P is the pressure
  • V is the volume

Enthalpy is a state function, which means it is determined by the state of the system, not how it got to that state. This makes changes in enthalpy, such as those that occur in chemical reactions, particularly important.

Understanding enthalpy of formation

Enthalpy of formation is the heat change that occurs when one mole of a compound is formed from its elements in their standard states. The standard state of an element is its form under standard conditions (1 atm pressure and specified temperature, usually 25 °C).

Standard enthalpy of formation (ΔH f °)

The standard enthalpy of formation, ΔH f °, of a pure element in its standard state is zero. This serves as a reference point. For example, the standard enthalpy of formation of oxygen gas, O 2 (g) is zero because it is in its elemental form.

Consider a simple chemical reaction in which hydrogen gas and oxygen gas react to form water:

2H 2 (g) + O 2 (g) → 2H 2 O(l)

To find the enthalpy of water formation, you would write the equation for the formation of one mole of water:

H 2 (g) + 1/2 O 2 (g) → H 2 O(l)

Here, ΔH f ° for water is the temperature change when one mole of water is formed from its elements in their standard states.

Visual example

H2 Gas O2 Gas H2O Liquid

In this example, hydrogen and oxygen gases react to form liquid water, and the thermal change associated with this reaction is the enthalpy of water formation.

Calculating enthalpy change from Hess's law

Sometimes a direct measurement of ΔH f ° is not possible. Instead, Hess's law can be used. According to Hess's law, the total enthalpy change for a chemical reaction is the same no matter what path it takes. This allows us to use known enthalpies to calculate unknown enthalpies by considering hypothetical steps.

For example, to calculate the enthalpy of formation of methane (CH 4), we use:

C(s) + 2H 2 (g) → CH 4 (g)

The enthalpy change for this process can be determined using known values from other processes that form CH 4 (g) from its elements.

Text example

Using Hess's law, suppose we have the following reactions with their enthalpy changes respectively:

1. C(s) + O 2 (g) → CO 2 (g) ΔH = -393.5 kJ/mol
2. H 2 (g) + 1/2 O 2 (g) → H 2 O(l) ΔH = -285.8 kJ/mol
3. CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O(l) ΔH = -890.4 kJ/mol

We can express the formation of CH 4 (g) from C(s) and H 2 (g) using these reactions, and apply Hess's law to find ΔH f ° of methane.

Understanding combustion enthalpy

The enthalpy of combustion is the heat change that occurs when one mole of a substance burns completely in oxygen under standard conditions. This enthalpy change is usually exothermic, releasing energy.

Standard enthalpy of combustion (ΔH c °)

For example, consider the combustion of methane, a common fuel:

CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O(l)

The enthalpy change in this reaction, ΔH c °, represents the standard enthalpy of combustion of methane.

Visual example

CH 4 Gas + O 2 O2 Gas CO2 Gas + 2H 2 O Liquid

In this visual illustration, methane and oxygen react to form carbon dioxide and water. The energy change in this combustion process is the enthalpy of combustion.

Energy release in combustion

Combustion reactions are typically highly exothermic. The energy released during combustion is used in a variety of applications, from heating homes to powering engines. This is why understanding the enthalpy of combustion is important in both chemical thermodynamics and practical energy management.

Calculating enthalpy using combustion data

The enthalpy of formation of compounds can be determined by using the known enthalpy of combustion along with Hess's law. By treating the reactions as a series of simple combustion reactions, chemists can calculate the energy changes required.

Text example

Suppose we know the following enthalpy of combustion:

1. C(s) + O 2 (g) → CO 2 (g) ΔH c ° = -393.5 kJ/mol
2. H 2 (g) + 1/2 O 2 (g) → H 2 O(l) ΔH c ° = -285.8 kJ/mol

Using these values, you can calculate the enthalpy of formation of a compound such as ethanol, C 2 H 5 OH from its combustion products.

Applications of enthalpy concepts

Understanding and calculating enthalpy changes helps us in several ways:

  • Designing energy-efficient industrial processes.
  • Developing sustainable and clean energy sources.
  • Understanding atmospheric phenomena related to energy transformation.
  • Optimization of chemical reactions in laboratories and industries for desired energy production.

Real World Energy Management

The applications of understanding enthalpy are widespread in real-world scenarios, from maximizing fuel efficiency in power plants to advancing refrigeration technology. The concepts ensure that engineers and scientists can design processes that take full advantage of energy transformations in chemical reactions.


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