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 ↓


Unusual Properties of the First Elements


The periodic table is a remarkable tool used for the classification and organization of elements in chemistry. Elements are arranged based on their atomic numbers and grouped with other elements with similar properties. An interesting phenomenon in chemistry is the unique, often unexpected, properties of the first element in any group in the periodic table. These unusual properties are primarily due to their small atomic size, high electronegativities, and the absence of d-orbitals. In this lesson, we will delve deeper into the reasons for these anomalies and explore several examples to better understand them.

Understanding unusual properties

The main reasons for the unusual behaviour of the first elements in each group are:

  • Small atomic size: The atomic size of the first element is usually much smaller than that of the other members of the group.
  • High electronegativities: Being smaller, these elements hold on to their valence electrons more tightly, which leads to their high electronegativities.
  • Absence of d-orbitals: Elements from the second period onwards may have d-orbitals for bond formation but elements in the first group do not.

Examples of abnormal behavior

1. Hydrogen

Hydrogen is the first element in the periodic table and is a classic example of unusual behavior. It has unique properties that make it quite different from the elements in group 1 (alkali metals) and group 17 (halogens). For example:

H 2 (molecular hydrogen) is a nonmetal and forms covalent bonds.

Unlike the alkali metals, which form ionic compounds, hydrogen forms covalent compounds. For example, in H 2 O (water), hydrogen shares electrons with oxygen.

2. Lithium and Beryllium

Lithium (Li) and beryllium (Be) are the first members of groups 1 and 2, respectively, with properties that are unusual compared to their own groups.

Lithium: Although it is in the alkali metal group, lithium forms more covalent compounds. For example, LiCl (lithium chloride) shows more covalent properties than the other alkali chlorides.

Beryllium: Known for its diagonal bond with aluminum, beryllium compounds are largely covalent. BeCl 2 (beryllium chloride) has a polymeric structure, unlike the ionic chlorides of other group 2 elements.

3. Boron

Boron is the first element of group 13 and shows some unique behaviour.

Boric Acid: Acts as a monobasic acid unlike other members of H 3 BO 3 13 group which have minimal or no basic properties.

The Lewis acid character of boron compounds, such as BF 3 (boron trifluoride), is important because of its inability to expand its octet.

Visualizing unusual behaviors

Atomic size decreases in a period

This diagram shows how atomic size decreases as we move across a period from left to right. The first element in each group has a smaller atomic radius, which affects its chemical behavior.

Comparison of properties with subsequent group elements

4. Carbon

Carbon, the first element of group 14, exhibits unique properties that are quite different from other elements in the group.

Tetravalent: Carbon always forms four covalent bonds, as seen in methane (CH 4).

Chainability: Carbon has the unique ability to form long chains of carbon atoms, which is not evident in other group 14 elements such as silicon.

5. Nitrogen

Nitrogen, the first element of group 15, has several unique characteristics:

Diatomic Molecules: Nitrogen exists as a diatomic molecule (N 2), while the other group 15 elements form larger structures.

Triple Bond: The presence of triple bond in N 2 makes it very stable, which is not found in other members of the group.

Role of bonding and structure

When considering the unusual properties of the first elements, analysis of bonding and structure helps to provide clarity:

Hybridization: First elements often display different hybridization patterns than their group counterparts. For example, in group 14, carbon displays sp 3 hybridization, forming strong sigma bonds.

C H Sigma Bonds

Periodic trends and their impact

Periodic trends such as ionization energy, electronegativities, and atomic radii strongly influence the anomalies found in the first elements:

High ionization energy: The first elements usually have high ionization energy. For example, helium and neon have high ionization energy due to their full shells, making them less reactive.

Electronegativity: The first elements are often the most electronegative in their group. For example, oxygen is more electronegative than its group counterpart sulfur.

Exploring further examples

6. Oxygen

Oxygen is the first element of group 16 and shows special behaviour.

High electronegativities: Oxygen is highly electronegative, due to which it has the ability to form hydrogen bonds, which is lacking in sulphur.

Allotropes: Oxygen forms ozone (O 3) allotrope, whereas sulphur has many allotropes like S8.

7. Fluorine

Fluorine represents the halogen in group 17:

Most electronegative element: This high electronegativity allows fluorine to form strong ionic bonds, such as sodium fluoride (NaF).

Reactivity: As the most reactive halogen, fluorine can form compounds with noble gases such as xenon.

Conclusion

The unusual properties of the first elements in the periodic table arise from their unique position and electronic configuration. These properties are important for understanding the chemistry and behavior of these elements. The distinctive features are mainly due to their small size, high electronegativities, and the absence of d-orbital participation. By exploring examples such as hydrogen, lithium, carbon, and fluorine, we gain a greater appreciation of the diversity and complexity found in the periodic table.


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Unusual Properties of the First Elements

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