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 → Atomic Structure ↓


Valency, ion formation and oxidation state


The concepts of valency, ion formation, and oxidation states are important in understanding how elements interact with each other to form different compounds. These concepts are important for explaining the behavior and characteristics of elements, especially when they engage in chemical reactions. In this lesson, we will delve deeper into these concepts using simple language, visual examples, and concrete textual examples.

Valency

Valency is the ability of an atom to combine with other atoms. It tells how many electrons an atom will gain, lose, or share when forming a chemical bond. Valency is affected by the number of electrons in the atom's outermost shell, known as the valence shell.

To understand valency, consider the following examples:

        Summarized valency:
        - Hydrogen (H): 1
        - Oxygen (O): 2
        - Nitrogen (N): 3
        - Carbon (C): 4

To understand this better, let's imagine an oxygen atom:

Hey

Oxygen has six electrons in its outer shell, but the oxygen atom needs eight electrons to be stable. Therefore, the valency of oxygen is 2 because it needs two more electrons to complete its octet.

Ion formation

When atoms gain or lose electrons, they form ions. An ion is an atom or group of atoms that has an electrical charge. Atoms become ions to achieve a more stable electron arrangement.

Types of ions

  • Cation: Positively charged ions formed by losing electrons. Example, Na^+
  • Anion: Negatively charged ion formed by gaining electrons. Example, Cl^-

Consider the sodium atom (Na), which has atomic number 11 and the electron arrangement 2, 8, 1. When sodium loses an electron, it forms the sodium ion:

        na → na⁺ + e⁻

Here is an illustration of sodium ion formation:

No no⁺

Similarly, the chlorine atom will gain one electron and become a chloride ion:

        Cl + e⁻ → Cl⁻

The visualization of chloride ion formation is as follows:

Chlorine Cl⁻

Oxidation states

Oxidation states or oxidation numbers help us understand the degree of oxidation or reduction of an atom in a compound. They indicate the hypothetical charge that an atom would have if all bonds between atoms of different elements were 100% ionic.

Rules for determining the oxidation state

  1. The oxidation state of a free element (uncombined element) is zero.
  2. The oxidation state of a monatomic ion is equal to the charge of the ion.
  3. The oxidation state of oxygen is normally -2, but in peroxides it is -1.
  4. The oxidation state of hydrogen is +1, but when bonded with metals, it is -1.
  5. In a neutral compound, the sum of the oxidation states of all atoms is zero. In a polyatomic ion, the sum is equal to the charge on the ion.

Consider the molecule H₂O (water). Oxygen usually has an oxidation state of -2, and since it contains two hydrogen atoms, each of which has an oxidation state of +1, the total charge is:

        2(H) + 1(O) = 0
        2(+1) + (-2) = 0

Here's a simplified illustration of the oxidation states in water:

H2O +1 -2 +1

Let us consider another example - NaCl (sodium chloride). Here, the oxidation state of sodium is +1, while the oxidation state of chlorine is -1, resulting in:

        Na^+ + Cl^- = 0
        +1 + (-1) = 0

The oxidation state representation for NaCl is as follows:

sodium chloride +1 -1

Applications and significance

The concepts of valency, ion formation, and oxidation states are fundamental in the study of chemistry because they allow us to understand the formation and transformation of molecules in chemical reactions. Knowing how atoms can combine and transform helps chemists develop new materials, medicines, and solve environmental problems.

These concepts form the basis for more advanced topics of chemistry, such as molecular geometry, chemical bonding theory, and electrochemical reactions.

Consider the role of these concepts in determining the process of rusting of iron, which is represented by the formation of iron oxide (Fe₂O₃):

        4Fe + 3O₂ → 2Fe₂O₃

Here, we use the concept of oxidation state to identify that iron is oxidized from 0 in Fe to +3 in Fe₂O₃.

Through these insights, students can appreciate the complex dance of atoms and energy that creates all chemical phenomena. As they progress in their studies, these fundamental concepts will continue to serve as important tools for understanding and manipulating the behavior of matter.


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Valency, ion formation and oxidation state

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