Grade 10
  1. Matter and its properties  1.1. States of matter  1.2. Physical and chemical properties of matter  1.3. Classification of substances (elements, compounds, mixtures)  1.4. Changes in Matter (Physical and Chemical Changes)  1.5. Law of conservation of mass and law of definite proportions
  2. Atomic Structure  2.1. Historical development of the atomic model  2.2. Subatomic particles (protons, neutrons, electrons)  2.3. Atomic number, mass number and isotopes  2.4. Electronic Configuration and Energy Levels  2.5. Valency, ion formation and oxidation state
  3. Periodic table  3.1. History and Development of the Periodic Table  3.2. Modern Periodic Law and Periodic Trends  3.3. Groups and Periods in the Periodic Table  3.4. Trends in the Periodic Table  3.5. Metals, non-metals, metalloids and their properties  3.6. Noble gases and their chemical inertness
  4. Chemical bond  4.1. Formation of Chemical Bonds (Octet Rule)  4.2. Ionic Bonding and Properties of Ionic Compounds  4.3. Covalent Bond and Properties of Covalent Compounds  4.4. Metallic bond and its properties  4.5. Molecular geometry and VSEPR theory  4.6. Polarity and dipole moment of molecules  4.7. Intermolecular forces
  5. Chemical Reactions and Equations  5.1. Types of Chemical Reactions  5.2. Writing and balancing chemical equations  5.3. Law of conservation of mass in chemical reactions  5.4. Factors Affecting Chemical Reactions  5.5. Catalysts and their role in chemical reactions  5.6. Exothermic and Endothermic Reactions
  6. Stoichiometry and Chemical Calculations  6.1. Mole concept and Avogadro number  6.2. Molar Mass and Gram-Mole Conversions  6.3. Empirical and molecular formula  6.4. Stoichiometric calculations and chemical yield  6.5. Limiting and Excess Reactants
  7. Acids, Bases and Salts  7.1. Properties and Definitions of Acids and Bases (Arrhenius, Bronsted-Lowry, Lewis)  7.2. pH Scale, pOH, and Indicators  7.3. Neutralization reactions and salt formation  7.4. Strength of Acids and Bases (Strong vs. Weak Acids and Bases)  7.5. Common Acids and Bases in Daily Life  7.6. Salts, their solubility and applications
  8. Metals and Nonmetals  8.1. Physical and chemical properties of metals  8.2. Physical and Chemical Properties of Non-Metals  8.3. Reactivity Series of Metals and Displacement Reactions  8.4. Extraction of metals from ores  8.5. Corrosion, its causes and prevention  8.6. Alloys, their composition and uses
  9. Carbon and its compounds  9.1. Allotropes of Carbon  9.2. Hydrocarbons and their classification (alkanes, alkenes, alkynes)  9.3. Functional groups in organic compounds  9.4. Nomenclature of organic compounds (IUPAC system)  9.5. Isomerism in organic chemistry (structural and geometrical)  9.6. Important organic reactions (combustion, addition, substitution, polymerization)  9.7. Applications of organic compounds in industry and daily life
  10. Environmental Chemistry  10.1. Air Pollution (Sources, Effects and Control)  10.2. Water Pollution (Causes, Effects and Remedies)  10.3. Greenhouse effect and global warming  10.4. Ozone layer depletion and its effects  10.5. Green Chemistry and Its Applications  10.6. sustainable chemistry practice
  11. Thermochemistry  11.1. Heat and Energy Changes in Chemical Reactions  11.2. Exothermic and Endothermic Reactions  11.3. Enthalpy, enthalpy change, and heat of reaction  11.4. Specific heat capacity and calorimetry  11.5. Energy Changes in Combustion Reactions
  12. Electrochemistry  12.1. Electrolytes and non-electrolytes  12.2. Electrolysis and its applications (electroplating, extraction of metals)  12.3. Redox reactions (oxidation and reduction)  12.4. Galvanic cell and electrochemical series  12.5. Corrosion and electrochemical prevention methods
  13. Chemical kinetics and equilibrium  13.1. Rate of chemical reactions and its measurement  13.2. Factors affecting the reaction rate (catalyst, temperature, concentration)  13.3. Reversible and irreversible reactions  13.4. Dynamic equilibrium and equilibrium constant  13.5. Le Chatelier's principle and its applications
  14. Gases and Gas Laws  14.1. Properties of gases and kinetic molecular theory  14.2. Boyle's Law (Pressure-Volume Relation)  14.3. Charles's Law  14.4. Gay-Lussac's Law  14.5. Combined Gas Law and Ideal Gas Law  14.6. Real Gases vs Ideal Gases
  15. Nuclear Chemistry  15.1. Introduction to Radioactivity and Nuclear Reactions  15.2. Types of radioactive decay (alpha, beta, gamma)  15.3. Half-life and applications of radioactive isotopes  15.4. Nuclear fission and fusion reactions  15.5. Nuclear power generation and its environmental impact

Grade 10 → Chemical kinetics and equilibrium ↓


Reversible and irreversible reactions


In the world of chemistry, reactions tell us how substances transform into new substances. To better understand how these reactions work, it is important to know the difference between reversible and irreversible reactions. Both types of reactions play an important role in chemical kinetics and equilibrium.

Reversible reactions

Reversible reactions are chemical reactions where reactants form products, which can later be converted back into reactants under certain conditions. This means that the reaction can go both forward and backward. Let's write a simple equation to represent a reversible reaction:

A + B ⇌ C + D

The double arrow (⇌) indicates that the reaction can proceed in both directions: left to right and right to left. How far a reversible reaction proceeds in each direction depends on factors such as temperature, pressure, and concentration.

Examples of reversible reactions

A classic example of a reversible reaction is the Haber process:

N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g)

In this reaction, nitrogen gas (N 2) reacts with hydrogen gas (H 2) to form ammonia (NH 3). However, under certain conditions, ammonia can break back down into nitrogen and hydrogen.

N 2 + 3H 2 2NH 3

Another common example is the dissociation of acetic acid in water:

CH 3 COOH ⇌ CH 3 COO - + H +

Acetic acid can dissociate into acetate ions and hydrogen ions, but these products can also recombine to form acetic acid.

Irreversible reactions

Irreversible reactions are chemical reactions where reactants turn into products and cannot easily turn back into reactants. These reactions occur in only one direction, usually releasing energy in the form of heat or light. Here's an example:

A + B → C + D

The single arrow (→) indicates that the reaction proceeds in only one direction, from reactants to products.

Examples of irreversible reactions

Combustion reactions are typical irreversible reactions. Consider the combustion of methane:

CH 4 + 2O 2 → CO 2 + 2H 2 O

Methane reacts with oxygen to form carbon dioxide and water. Once the reaction is complete, the products do not turn back into reactants under normal conditions.

CH 4 + 2O 2 CO 2 + 2H 2 O

Another example of an irreversible reaction is the formation of a salt from its constituent elements:

2Na + Cl 2 → 2NaCl

Sodium metal reacts with chlorine gas to form sodium chloride (table salt), and under normal conditions the products do not change back to sodium and chlorine.

Chemical equilibrium

In reversible reactions, chemical equilibrium is the state at which the rates of the forward and backward reactions are equal. When a system reaches equilibrium, the concentrations of reactants and products remain constant over time, although they are not necessarily equal.

Consider the equilibrium of the Haber process:

N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g)

At equilibrium, the rate at which nitrogen and hydrogen form ammonia is equal to the rate at which ammonia dissociates back into nitrogen and hydrogen.

Factors affecting chemical equilibrium

Many factors can affect the chemical equilibrium, including:

  • Concentration: Changing the concentration of reactants or products shifts the equilibrium. An increase in reactants usually shifts the equilibrium to the right, favoring product formation, and vice versa.
  • Pressure: For gases, changes in pressure affect the equilibrium, especially if the reaction involves a change in the number of moles of gas. An increase in pressure generally shifts the equilibrium to the side with fewer gas molecules.
  • Temperature: An increase in temperature generally favours the endothermic direction of the reaction, while a decrease in temperature favours the endothermic direction of the reaction.

Le Chatelier's principle

La Chatelier's principle states that if a dynamic equilibrium system is disturbed by a change in concentration, pressure, or temperature, the system will adjust to counteract the disturbance and reestablish equilibrium.

Let's see how this principle applies to a reversible reaction example:

2NO 2 (g) ⇌ N 2 O 4 (g)

If the concentration of NO 2 is increased, the equilibrium will shift to the right causing more N 2 O 4 to be produced.

Conclusion

Understanding reversible and irreversible reactions is essential for students studying chemistry, especially when learning about chemical kinetics and equilibrium. These concepts help explain how reactions work and how different conditions affect a reaction. Remember, reversible reactions can happen both ways, while irreversible reactions only happen one way, which largely affects how chemical processes are managed and applied in real-world scenarios.


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Reversible and irreversible reactions

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