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 ↓


Electrochemistry


Welcome to the wonderful world of electrochemistry! In this guide, we'll explore the fascinating science of how chemicals produce electricity, and how electricity can bring about chemical changes. Electrochemistry is the study of chemical processes that cause electrons to move. This movement of electrons is called electricity, and electrochemistry is about how chemicals and electrical power can be interchanged.

Basics of electrochemistry

At its core, electrochemistry involves two main types of reactions: oxidation and reduction, often called redox reactions. Let's break these down into simpler concepts:

Oxidation

Oxidation is a reaction in which a substance loses electrons. When a metal such as iron rusts, it is oxidized. Here's a simple example involving magnesium:

Mg → Mg2+ + 2e-

In this reaction, solid magnesium (Mg) loses two electrons (2e-) and becomes magnesium ion (Mg2+).

Reduction

Reduction is the opposite of oxidation. It is a reaction in which a substance gains electrons. Consider this example with copper ions:

Cu2+ + 2e- → Cu

In this case, the copper ion (Cu2+) gains two electrons and becomes solid copper (Cu).

Redox reactions

Redox reactions involve both oxidation and reduction occurring simultaneously. One substance is oxidized, and the other is reduced. Together, they allow the flow of electrons from one substance to another, which can be used to do work, such as generating electricity.

Example of a redox reaction

Zn + Cu2+ → Zn2+ + Cu

This reaction shows that zinc (Zn) is being oxidized and copper ions (Cu2+) are being reduced. Electrons move from the zinc to the copper ions, turning them into copper metal.

Galvanic cells

A galvanic cell or voltaic cell is an electrochemical cell that uses spontaneous redox reactions to generate electricity. It consists of two electrodes placed in an electrolyte solution, connected by a wire and a salt bridge.

Zinc Copper Salt bridge

In our example of a galvanic cell made of zinc and copper:

  • The zinc electrode (anode) loses electrons and undergoes oxidation (Zn → Zn2+ + 2e-).
  • The copper electrode (cathode) gains electrons and undergoes reduction (Cu2+ + 2e- → Cu).
  • The salt bridge allows the ions to flow, maintaining neutrality in the solution.

Electrochemical series

The electrochemical series is a list of elements arranged according to their standard electrode potential. This list enables us to predict the outcomes of redox reactions, determine which elements easily lose or gain electrons, and the voltages produced in electrochemical cells.

Li | Lithium | E° = -3.04 V ... Zn | Zinc | E° = -0.76 V ... Cu | Copper | E° = +0.34 V ... Au | Gold | E° = +1.50 V

In this table:

  • Elements higher up lose electrons and are easily oxidized (e.g., lithium).
  • Elements at the bottom easily accept electrons and get reduced (for example, gold).
  • Elements with more negative standard electrode potentials (e.g., zinc) are better reducing agents, while elements with positive potentials (e.g., copper) are better oxidizing agents.

Applications of electrochemistry

Electrochemistry is not just theoretical; it has practical applications in our daily lives and technology.

Batteries

Batteries are devices that store and release electrical energy using electrochemical reactions. Alkaline batteries are a common example:

Zn (anode) MnO2 (cathode)

In alkaline batteries:

  • Zn acts as the anode, undergoing oxidation.
  • MnO2 acts as cathode and undergoes reduction.

Electrolysis

Electrolysis is a non-spontaneous chemical reaction driven by an external electric current. It is used, for example, in the purification of metals or in electroplating.

Cathode Anode Power source

For example, in the electrolysis of water:

  • Hydrogen gas is produced at the cathode (2H2O + 2e- → H2 + 2OH-).
  • Oxygen gas is produced at the anode (2H2O → O2 + 4H+ + 4e-).

Concluding remarks

Electrochemistry combines concepts of chemistry and electrochemical technology, providing practical understanding in areas ranging from biological systems to industrial processes. Understanding the principles of redox reactions, galvanic cells, and electrolysis opens the door to innovations and advancements in a variety of scientific and engineering fields.


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Electrochemistry

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