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 ↓


The concept of balance


Equilibrium is a fundamental concept in chemistry that describes the state of equilibrium in a chemical reaction. It is a dynamic situation where the rate of the forward reaction is equal to the rate of the reverse reaction. Equilibrium does not mean that the concentrations of reactants and products are equal or that the reaction has stopped; instead, it suggests that their concentrations remain constant over time.

Understanding chemical equilibrium

Chemical equilibrium occurs in a closed system where neither reactants nor products can escape. In such a system, any reaction proceeding in one direction is counteracted by an equal reaction in the opposite direction. It is dynamic because although macroscopic properties (such as concentration, pressure, etc.) remain unchanged, reactions continue to occur at the molecular level.

To represent a reaction in equilibrium, we use a two-headed arrow in a chemical equation:

A + B ⇌ C + D

Here, the forward reaction is A + B → C + D and the reverse reaction is C + D → A + B

Law of mass action

At equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction. The law of mass action provides a mathematical way to understand this balance. It states that at a constant temperature, the concentrations of reactants and products can be expressed in terms of the equilibrium constant K

K = [C]^c [D]^d / [A]^a [B]^b

Here, [A], [B], [C], and [D] represent the molar concentrations of the compounds, and a, b, c, and d are their respective coefficients from the balanced chemical equation.

The idea of balance

Consider a simple system of water in a closed container. When this water evaporates, it turns into vapor. In a closed system, the vapor eventually begins to condense back into liquid water. When the rate of evaporation equals the rate of condensation, equilibrium is achieved.

liquid H2O H2O vapor Evaporation ⇌ Condensation

Characteristics of the balance

Some important characteristics of a system in chemical equilibrium are:

  • Dynamic process: Even though no macroscopic changes can be observed, reactions continue at the molecular level.
  • Closed system required: Equilibrium can only be achieved in a closed system, where no substance can escape.
  • No net change: The concentrations of reactants and products remain constant.
  • Unaffected by initial concentration: The final equilibrium position does not depend on the initial concentration; however, it will affect how long it takes to reach equilibrium.

Le Chatelier's principle

Le Chatelier's principle describes how a system in equilibrium reacts to a disturbance. If a dynamic equilibrium is disturbed by changing conditions (concentration, temperature, pressure), the equilibrium position shifts to counteract the change. This principle helps to predict the direction in which a system will move due to an external stress.

Changes in concentration

If more reactants are added to the system, the equilibrium position shifts toward the products, causing a decrease in the added reactants. Conversely, if more products are added, the system shifts toward the reactants.

Example:

N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

Adding more N2 shifts the equilibrium to the right, producing more NH3.

Changing pressure

Changes in pressure only affect the equilibrium in reactions involving gases. Increased pressure favors the side that has fewer moles of gas, while decreased pressure favors the side that has more moles of gas.

Example:

2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

Increasing the pressure shifts the equilibrium to the right, causing more SO3 to form.

Changing temperatures

An increase in temperature promotes endothermic reactions (absorbs heat), while a decrease promotes exothermic reactions (releases heat).

Example:

N2 (g) + 3H2 (g) ⇌ 2NH3 (g) + Heat

An increase in temperature shifts the equilibrium to the left, reducing NH3 production.

Common misconceptions

Several misconceptions arise regarding chemical equilibrium:

  • Equilibrium means equal concentration: It does not mean equal concentration but constant concentration.
  • The reactions stop at equilibrium: instead, the reactions continue at the same rate on both sides.
  • Adding a catalyst affects the equilibrium: the catalyst speeds up the rate at which equilibrium is reached, but does not change the position of the equilibrium.

Applications of chemical equilibrium

Understanding chemical equilibrium has many practical applications:

  • Industrial synthesis: The Haber process for industrial ammonia synthesis relies on equilibrium principles to maximize production.
  • Biological systems: Many biochemical reactions are equilibrium-based, such as the binding of oxygen to hemoglobin.
  • Environmental science: Understanding equilibrium is important in studying chemical distributions and reactions in atmospheric and environmental contexts.

Conclusion

Chemical equilibrium provides the basis for understanding how reactions proceed and balance in closed systems. It is a dynamic, ongoing process where the concentrations of reactants and products are constant. By applying the principles of chemical equilibrium, chemists can predict how changes in conditions such as concentration, pressure, and temperature will affect the system. This understanding is fundamental to chemical industries, environmental science, and biochemistry.

In summary, equilibrium is not only central to chemistry, but also impacts a variety of fields and real-world applications.


Grade 11 → 7.1


U
username
0%
completed in Grade 11

← Prev (7)
Balance
Next (7.2) →
Equilibrium constant and its applications

Comments 



The concept of balance

https://www.chemistryelite.com/g11/7.1?ceal=en