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 → Structure of the atom ↓


Electronic configuration of elements


The concept of electronic configuration in chemistry is fundamental to understanding how elements interact, form bonds, and participate in chemical reactions. The electronic configuration of an element describes the distribution of its electrons among different atomic orbitals. In this comprehensive lesson, we will delve deep into the details of electronic configuration, providing both textual descriptions and visual representations for ease of understanding.

What is electronic configuration?

Electronic configuration refers to the specific arrangement of electrons in the orbitals of an atom or molecule. Electrons are distributed in orbitals around the nucleus in a way that is determined by several rules, including the Pauli exclusion principle, Hund's rule, and the Aufbau principle. Understanding electronic configuration helps predict the chemical properties of elements, their location in the periodic table, and their reactivity.

Principles governing electronic configuration

Aufbau principle

The Aufbau principle, derived from the German word "aufbauen" meaning "to raise", states that electrons fill atomic orbitals starting from the lowest available energy level before filling higher levels. This is similar to filling a container from the bottom before moving upwards. The order of filling is based on the increasing energy levels of the orbitals.

Pauli exclusion principle

The Pauli exclusion principle, formulated by Wolfgang Pauli, stipulates that no two electrons in an atom can have the same set of four quantum numbers. In simple terms, an atomic orbital can hold a maximum of two electrons, and the spins of these electrons must be opposite.

Hund's law

According to Hund's rule, each orbital in a subshell gets one electron before any orbital gets a second electron. This reduces electron-electron repulsion and makes the electron configuration more stable.

Structure of electronic configuration

Electronic configurations are often represented using a notation that shows the energy level, the type of orbital, and the number of electrons in those orbitals. For example:

1s2 2s2 2p6

Let's break down this notation:

  • 1s2 - This represents two electrons in the 1s orbital.
  • 2s2 - This represents two electrons in the 2s orbital.
  • 2p6 - This represents six electrons in the 2p orbitals.

Visualization of electron shells

The distribution of electrons in different shells can be seen through the following illustration:

Example: Helium (He) has the configuration 1s2. Neon (Ne) has the configuration 1s2 2s2 2p6.

Writing the electronic configuration

Step-by-step guide

  1. Identify the element's atomic number. This tells you the total number of electrons.
  2. Use the Aufbau principle to place electrons in orbitals in order of increasing energy level.
  3. Apply Hund's rule to distribute electrons in the orbitals of the same subshell.
  4. Remember the Pauli exclusion principle, make sure there are no more than two electrons per orbital.

Example

1. Hydrogen (H)

Atomic Number: 1

Configuration: 1s1

2. Carbon (C)

Atomic Number: 6

Configuration: 1s2 2s2 2p2

Here, we distribute the 2p electrons according to Hund's rule, placing one electron in each p orbital before pairing.

3. Sodium (Na)

Atomic Number: 11

Configuration: 1s2 2s2 2p6 3s1

Orbital diagram

Orbital diagrams visually represent the arrangement of electrons in an atom's electron orbitals. They use boxes to represent orbitals and arrows to indicate electrons and their spins. The following describes how to draw an orbital diagram:

1. Understanding orbitals

According to the Pauli exclusion principle, each orbital can hold a maximum of two electrons with opposite spins. Different types of orbitals include:

  • s-orbitals: spherical, can hold 2 electrons.
  • p-orbitals: Dumbbell shaped, each p subshell has 3 orbitals, can hold a total of 6 electrons.
  • d-Orbitals: More complex shape, can hold 5 orbitals per d subshell, 10 electrons in total.
  • f-Orbitals: Complex shape, can hold 7 orbitals per f subshell, total of 14 electrons.

2. Steps to drawing an orbital diagram

  1. Write the electronic configuration of the element.
  2. Draw boxes or lines for each orbital and fill them with electrons, indicated by arrows.
  3. Follow the Hund rule by filling each orbital with one electron before doubling it.
  4. Make sure the electron spins are indicated by up (↑) and down (↓) arrows in the same orbital.

3. Example of an orbital diagram

Example: Oxygen (O) - atomic number 8

Configuration: 1s2 2s2 2p4

Orbital diagram:

1s ↑↓   
2s ↑↓   
2P ↑ ↑ ↑↓

Explanation: The 1s and 2s orbitals are filled first, with two electrons each. In the 2p orbitals, we follow Hund's rule, placing one electron in each p orbital before pairing in the third box.

Example: Chlorine (Cl) - atomic number 17

Configuration: 1s2 2s2 2p6 3s2 3p5

Orbital diagram:

1s ↑↓
2s ↑↓
2P ↑↓ ↑↓ ↑↓ ↑↓
3s ↑↓
3p ↑↓ ↑↓ ↑

Explanation: The orbital diagram follows the same logic: initially filling the low-energy orbitals and then distributing the electrons in the 3p orbitals according to Hund's rule.

Conclusion

Understanding electronic configurations is vital to predicting and explaining the chemical behavior of elements. This involves applying principles such as the Aufbau principle, the Pauli exclusion principle, and Hund's rule. By looking at electron arrangements through numerical notation and orbital diagrams, students can gain a comprehensive understanding of atomic structures and the fundamentals of chemistry.


Grade 11 → 2.8


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Electronic configuration of elements

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