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 → Periodic table ↓


Trends in the Periodic Table


The periodic table is an essential tool in chemistry that provides us with a lot of information about the elements. As Class 10 Chemistry students, understanding the trends in the periodic table will enhance your understanding of the chemical behaviour of the elements. The three major trends to explore include atomic radius, ionisation energy, and electronegativities. Let's dive deeper into each of these properties.

Atomic radius

The atomic radius is the distance from the center of an atom's nucleus to the outermost electron shell. This concept is important to understand because the size of an atom affects its chemical properties and reaction tendencies.

During a period

As you move from left to right across a period on the periodic table, the atomic radius decreases. This may seem surprising, so let's find out why this happens:

  • The number of protons in the nucleus increases across a period. This increased positive charge pulls the electrons closer to the nucleus.
  • Although the number of electrons also increases, they are added to the same shell and not to a new shell, so the added electrons do not increase the size of the atom appreciably.

Example: Consider elements in the second period, such as lithium (Li) and fluorine (F). Lithium has an atomic number of 3, while fluorine has 9 As you move from lithium to fluorine, the atomic radius decreases because of the increased nuclear charge.

                  ,
Li(Lithium) ----> F(Fluorine)
  <-------- decrease in atomic radius ------->

Group down

Conversely, as you move down a group in the periodic table, the atomic radius increases. This is because:

  • Going down a group means adding a new electron shell, which increases the distance between the outermost electron and the nucleus, thus increasing the size of the atom.

Example: In group 1, consider hydrogen (H) and cesium (Cs). Although hydrogen has an atomic number of 1 and cesium has an atomic number of 55, cesium is much larger because of its extra electron shell.

         ,
H(Hydrogen) -------------------------> Cs(Caesium)
  <-------- Increase in atomic radius --------->

Ionization energy

Ionization energy refers to the energy required to remove an electron from a neutral atom in the gaseous state. It is an important concept because it affects an element's ability to participate in chemical reactions.

During a period

Ionization energy generally increases as you move from left to right across a period. This is because:

  • As you move across a period, the atom has more protons, which means more positive charge, which attracts electrons more strongly.
  • The greater attraction between the nucleus and the valence electron means that more energy will be required to remove it.

Example: In the second period, lithium (Li) has a lower ionization energy than neon (Ne) because the nucleus of neon holds on to the electrons more tightly due to its higher charge.

                    ,
Li(Lithium) ------> Ne(Neon)
 <--- increase in ionization energy ---->

Group down

As you go down the group, ionization energy decreases. This is because:

  • As mentioned, going down a group, more electron shells are added, which increases the distance of the valence electrons from the nucleus.
  • The increased distance weakens the attraction between the nucleus and the outermost electrons, making them easier to remove.

Example: In group 1, caesium (Cs) has a lower ionization energy than lithium (Li) due to greater shielding and distance from the nucleus.

       ,
Li(Lithium) ----------------------> Cs(Caesium)
 <--- decrease in ionization energy ---->

Electronegativity

Electronegativity is the measure of an atom's ability to attract electrons and form bonds with them. It plays an important role in determining the type of bond that forms between atoms.

During a period

Electronegativity increases as you move from left to right in a period. This is because:

  • At each pass of an element, more protons in the nucleus means more pull on the shared electrons.
  • This increased ability to attract electrons results in higher electronegativities.

Example: In the second period, lithium (Li) is less electronegative than fluorine (F), because fluorine has a greater tendency to attract electrons towards itself.

                   ,
Li(Lithium) -----> F(Fluorine)
 <---- increase in electronegativities ---->

Group down

Generally, a decrease in electronegativities is observed as we go down the group. This is because:

  • Larger atomic size means that the valence electrons are farther from the nucleus and the effective pull of the nucleus on these electrons is reduced.

Example: In group 17, fluorine (F) is more electronegative than iodine (I), because fluorine's electrons are closer to the nucleus, enabling it to attract additional electrons more effectively.

         ,
F(fluorine) ------------------> I(iodine)
 <--- decrease in electronegativities ---->

Conclusion

Understanding the trends in atomic radius, ionization energy, and electronegativities in the periodic table helps us predict and explain the chemical behavior of elements. This knowledge is foundational to understanding more complex topics in chemistry. Continue to explore and relate these trends as you move forward in your studies. By doing so, you will develop a more profound understanding and appreciation of the physical world around you.


Grade 10 → 3.4


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Trends in the Periodic Table

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