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 bond


Chemical bonding is the process in which atoms or molecules combine to form more complex structures through interactions between electrons. Chemical bonds are important because they hold the building blocks of matter together, contributing to the properties and behavior of substances.

Understanding atoms

An atom consists of a nucleus, which contains protons and neutrons, and electrons that revolve around the nucleus. The number of protons in the nucleus of an atom determines the identity of the element and is known as the atomic number. Electrons orbit the nucleus in different energy levels or shells, and their arrangement is important for bonding.

Electron configuration

Electrons are arranged in shells around the atom's nucleus. The first shell can have up to 2 electrons, the second shell can have up to 8, and so on. The electrons in the outermost shell are called valence electrons, and they are important in chemical bonding.

For example, in a helium atom with 2 electrons:
Shell 1: 2 electrons (complete)
    
In a neon atom with 10 electrons:
Shell 1: 2 electrons
Shell 2: 8 electrons (complete)

Why do atoms bond together?

Atoms bond to achieve more stable electron configurations. Atoms are generally more stable when their outer electron shells are full, similar to the noble gases that are naturally stable. When atoms react, they lose, gain, or share electrons to fill their outer shells, forming different types of chemical bonds.

Types of chemical bonds

Ionic bond

Ionic bonds form when electrons transfer from one atom to another, forming ions. An ion is an atom that has a net positive or negative charge due to the loss or gain of electrons. This usually occurs between metals and non-metals.

Example of ionic bond:

Sodium (Na) has 1 electron in its outer shell, chlorine (Cl) has 7 electrons.
Na (1 valence electron) --> Na⁺ + e⁻
Cl (7 valence electrons) + e⁻ --> Cl⁻
Na⁺ + Cl⁻ --> NaCl
no⁺ Cl⁻

The formation of oppositely charged ions results in an attractive force, forming an ionic bond that holds the ions together in a compound such as sodium chloride (NaCl).

Covalent bonds

Covalent bonds are formed when two atoms share one or more electron pairs. This type of bond usually occurs between non-metal atoms.

Example of covalent bond:

Hydrogen (H) atoms have 1 electron each.
By sharing their electrons, they can fill their outer shell, forming covalent bonds:

H : H (H₂ single covalent bond)

When the atoms share electrons equally, the bond is a nonpolar covalent bond. If the atoms do not share electrons equally, the bond is called a polar covalent bond.

Metal bonding

Metallic bonds are formed between metal atoms. In this type of bond, electrons are not shared between individual atoms. Instead, electrons move around freely in a 'sea' between the lattice of atoms. This electron mobility is responsible for many of the properties of metals, such as conductivity and malleability.

E⁻ E⁻ E⁻

Examples of bonds in everyday life

Chemical bonds exist everywhere around us. The air we breathe, the food we eat, and the objects we use are all products of atoms bonded in different ways. Understanding how these bonds work enriches our understanding of the physical world.

Common compounds

  • Water (H2O): Two hydrogen atoms and one oxygen atom form a covalent bond.
  • Carbon dioxide (CO2): One carbon atom forms double covalent bonds with two oxygen atoms.
  • Sodium chloride (NaCl): An ionic bond between sodium and chloride ions.

Role of electronegativity in bonding

Electronegativity is a measure of an atom's ability to attract and hold electrons. In general, nonmetals have higher electronegativities than metals. The difference in electronegativities between atoms affects the type of bond that forms:

  • If the electronegativities difference is large (usually > 1.7), an ionic bond is likely to form.
  • If the difference is small or zero, a covalent bond is formed.
  • With a moderate difference (between 0.4 and 1.7), a polar covalent bond is likely, where the electrons are shared unequally.
Example:
Na and Cl: high difference, ionic bond.
O and H: medium difference, polar covalent bond.
F and F: no difference, non-polar covalent bond.

Molecules and compounds

A molecule is a group of two or more atoms held together by covalent bonds. If the atoms are different, the molecule is a compound. Compounds can have completely different properties than the individual elements that make them up.

Example:
O₂ - Molecule composed of two oxygen atoms.
H₂O - Molecules and compounds composed of hydrogen and oxygen.

Macroscopic properties of substances

The type of chemical bond affects the macroscopic properties of substances:

  • Ionic compounds: Usually solid at room temperature, have high melting and boiling points, conduct electricity when melted or dissolved in water.
  • Covalent compounds: These may be gases, liquids or solids, have low melting and boiling points, they generally do not conduct electricity.
  • Metallic substances: good conductors of electricity and heat, malleable, ductile, lustrous.

Conclusion

Chemical bonding is fundamental to chemistry, enabling the formation of compounds and influencing the properties of substances. Whether through ionic, covalent or metallic bonding, the interactions between atoms create the diverse world of substances seen in nature. Understanding these bonds helps us understand the behavior and transformations of matter.


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Noble gases and their chemical inertness
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Chemical bond

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