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 ↓


Galvanic cell and electrochemical series


Galvanic cells, also known as voltaic cells, are devices used in electrochemistry to convert chemical energy into electrical energy. This process occurs through the redox reactions that take place in these cells. The concept of a galvanic cell is fundamental to understanding how batteries work and how electricity can be generated from chemical reactions.

What is a galvanic cell?

A galvanic cell consists of two different metals connected by a salt bridge or porous membrane. Each metal is immersed in an electrolyte solution. The two metals or electrodes have different tendencies to lose or gain electrons. This difference in tendency causes electrons to flow from one electrode to the other, creating an electric current.

ZnSO4 solution CuSO4 solution E- flow

This diagram shows a basic galvanic cell. Zinc (the anode) is on the left, and copper (the cathode) is on the right, containing solutions of ZnSO4 and CuSO4, respectively. Electrons flow from the zinc electrode to the copper electrode.

Components of a galvanic cell

Anode and cathode

The anode is the electrode where oxidation occurs. In a galvanic cell, it is the negative electrode. Electrons are generated at the anode and travel through an external circuit to the cathode. For example, in a zinc-copper galvanic cell, zinc is the anode:

Zn(s) → Zn2+ (aq) + 2e-

The cathode is the electrode where reduction takes place. It is the positive electrode in a galvanic cell. Electrons travel from the external circuit to the cathode. In a zinc-copper cell, copper is the cathode:

Cu2+ (aq) + 2e- → Cu(s)

Salt bridge

The salt bridge is a vital component that helps a galvanic cell function properly. It allows ions to move between the two half-cells so that charge balance can be maintained. It is usually made of a salt solution such as KCl or KNO3 that does not react with the chemicals in the cell.

Electrochemical series

The electrochemical series, also known as reactivity series, is a list of elements arranged by standard electrode potential. These potentials indicate the ability of an element to be oxidized or reduced. This series helps us predict the direction of redox reactions and also tells which electrode will be the anode or cathode in a galvanic cell.

Explanation of standard electrode potential

The standard electrode potential (E0) is measured in volts (V) and indicates the ability of a substance to gain electrons (reduction potential) or lose electrons (oxidation potential) under standard conditions. Elements with higher reduction potentials are more likely to gain electrons.

Hydrogen is given a standard reduction potential of 0.00 V, which serves as a reference point:

H2 → 2H+ + 2e- E0 = 0.00 V

Use of electrochemical series

Consider two metals, metal A and metal B, with their standard electrode potentials:

  • Metal A: E0 = -0.76 V (example: zinc)
  • Metal B: E0 = +0.34 V (example: copper)

Since copper has a high standard reduction potential, it acts as the cathode, and zinc acts as the anode. The electron flow is from zinc to copper.

Calculating cell potential

The overall cell potential can be calculated using the following equation:

Ecell = Ecathode - Eanode

For Zn-Cu cell:

  • Zinc (anode): Eanode = -0.76 V
  • Copper (cathode): Ecathode = +0.34 V

Thus, the cell potential is:

Ecell = 0.34 V - (-0.76 V) = 1.10 V

Applications of galvanic cells

Galvanic cells are used in a variety of applications to power devices. Here are some common uses:

  • Batteries: Galvanic cells are the basis of batteries that power electronic devices, vehicles, etc.
  • Corrosion prevention: Sacrificial anodes are used in galvanic cells to prevent corrosion of metal structures.
  • Electroplating: The process of electroplating uses galvanic cells to deposit metals on surfaces.

Example of a simple battery

A common example of a simple battery is the Daniell cell, with a zinc anode and a copper cathode:

  1. Anode reaction: Zn(s) → Zn2+ (aq) + 2e-
  2. Cathode reaction: Cu2+ (aq) + 2e- → Cu(s)

This basic concept can be extended to create more complex batteries, such as alkaline batteries, lithium-ion batteries, etc., depending on the materials and chemicals used.

Conclusion

Understanding the galvanic cell and the electrochemical series is essential to understanding how chemical reactions can produce electricity. These concepts are foundational in the development of batteries, which are vital to modern technology. By exploring the various components and reactions within a galvanic cell, we gain insight into the electrochemical processes that power many aspects of our daily lives.


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Redox reactions (oxidation and reduction)
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Galvanic cell and electrochemical series

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