Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Classification of elements and periodicity in properties ↓


Periodic trends in properties


The periodic table is a powerful tool in chemistry. It is arranged in such a way that the elements are listed in order of increasing atomic number, creating periodic patterns or trends. These trends provide valuable information about the properties of elements, helping scientists predict behavior in chemical reactions. Let's look at these trends in detail.

1. Atomic radius

The atomic radius is one of the most fundamental chemical properties. It refers to the size of the atom, more precisely, the distance from the nucleus to the outermost shell of electrons.

radius

Trends over a period

As you move from left to right across a period, the atomic number increases. This means that more protons are added to the nucleus, and electrons are added to the same shell. The increased nuclear charge pulls the electrons closer to the nucleus, resulting in a smaller atomic radius.

Example: Going from left to right in the 3rd period, atomic radius of Na (sodium) is larger than that of Cl (chlorine).

Downward trend in the group

The atomic radius increases as you move down a group. This is because each successive element has an additional electron shell, which more than accounts for the increased nuclear charge, increasing the size of the atom.

Example: Potassium (K) has a much larger radius than Lithium (Li).

2. Ionization energy

Ionization energy is the energy required to remove an electron from an isolated gaseous atom. It provides information about the strength of the electron bond of the atom.

Trends over a period

Ionization energy generally increases across a period. With a greater number of protons in the nucleus, the electrons are bound more tightly, making it harder to remove one.

Example: In the 2nd period, ionization energy of Na is less than that of Ar due to increase in nuclear charge across the period.

Downward trend in the group

Ionization energy decreases as we go down the group. The outer electrons are farther from the nucleus and experience a shielding effect from the inner-shell electrons, making them easier to remove.

Example: Comparing Li with Cs, caesium has a lower ionization energy because its outer electron is farther from the nucleus.

3. Electron affinity

Electron affinity measures how easily an atom can gain an electron. This property represents the energy change when an electron is added to a neutral atom.

Trends over a period

Generally, electron affinity becomes more negative (more energy is released) as you proceed across a period from left to right. Atoms more eagerly accept electrons to achieve a full valence shell.

Example: Chlorine (Cl) has a more negative electron affinity than sodium (Na).

Downward trend in the group

Electron affinity becomes less negative as we go down the group. The additional electron shells reduce the positive attraction of the nucleus on the incoming electrons.

Example: Electron affinity of F is greater than I (iodine).

4. Electronegativity

Electronegativity indicates the tendency of an atom to attract shared electrons in a chemical bond. It is an important factor in determining the nature of the bond between atoms.

Trends over a period

Electronegativity increases as you move across a period. Atoms on the right side of the table have more protons, creating a greater attraction for electrons.

Example: Oxygen (O) is more electronegative than carbon (C).

Downward trend in a group

Electronegativity decreases as we go down the group. Larger atomic radius and additional electron shells weaken the pull on shared electrons.

Example: In group 17, fluorine (F) is more electronegative than iodine (I).

Conclusion

Periodic trends in atomic radius, ionization energy, electron affinity and electronegativities reveal the systematic and predictable nature of the behaviour of elements. Understanding these trends helps scientists make informed predictions about chemical reactivity and the nature of bonds. Such patterns are fundamental to guiding intensive explorations, experiments and theoretical developments in chemistry.


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Modern Periodic Law and Periodic Table
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Periodic trends in properties

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