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 ↓


Aufbau principle


The Aufbau principle, derived from the German word "aufbauen" meaning "to build up", is a fundamental concept in chemistry that describes the order in which electrons occupy the orbitals of an atom. The Aufbau principle helps us understand and predict the electronic structure of atoms, which enables us to describe the chemical behavior and properties of elements.

Basic concepts of the Aufbau principle

The Aufbau principle is based on the idea of building up the electron configuration of an atom from the lowest energy level to the highest. According to this principle, electrons are added to the lowest energy orbitals available. As we move from lower atomic numbers to higher atomic numbers, electrons fill the subshells in a predictable order. Let's start with the essential rules that guide the Aufbau process:

  • Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers. This means that an orbital can hold a maximum of two electrons with opposite spins.
  • Hund's rule: Electrons will remain alone in degenerate orbitals before forming pairs. This reduces electron repulsion and keeps the atom stable.
  • Order of filling subshells: Electrons fill orbitals in order of increasing energy, often represented by an Aufbau diagram.

Sequence of filling the sub-shell

The order in which electrons fill the subshells can be determined using the n + l rule (also called Madelung rule). The rule states:

  • The sub-shells are filled in increasing order of n + l value.
  • If two sub-shells have the same n + l values, then the sub-shell with the lower n value will be filled first.

Here, n is the principal quantum number and l is the azimuthal quantum number. Consider the following chart which shows the order in which the orbitals are filled:

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p
1s 2s 2P 3s 3P 4S 3D

Use of the Aufbau principle

Let's apply the Aufbau principle to write the electron configuration of the elements. We will consider simple examples using hydrogen, helium, lithium and others.

Hydrogen:

Hydrogen has one electron.

1s 1

The single electron is located in the 1s orbital, which is the lowest energy.

Helium:

Helium has two electrons.

1 to 2

Both electrons occupy the 1s orbital and fill it completely.

Lithium:

Lithium has three electrons.

1s 2 2s 1

The first two electrons fill the 1s orbital, and the third electron enters the 2s orbital.

Continue this process for more complex atoms, such as carbon, nitrogen, or oxygen:

Carbon:

Carbon has six electrons.

1s 2 2s 2 2p 2

Electrons fill the orbitals in order of increasing energy. The 2p orbital begins to fill after the 2s orbital.

Oxygen:

Oxygen has eight electrons.

1s 2 2s 2 2p 4

Even though there are four electrons in the 2p orbital, the 2p subshell is still not completely filled.

Exceptions to the Aufbau principle

While the Aufbau principle provides a useful framework for constructing electron configurations, it is important to note some exceptions, particularly those relating to the transition metals and heavier elements. For example:

Copper (Cu):

Copper has 29 electrons.

[Ar] 3d 10 4s 1

Instead of 4s 2 3d 9 configuration, copper completely fills the 3d sub-shell before the 4s.

Chromium (Cr):

Chromium has 24 electrons.

[Ar] 3d 5 4s 1

The semi-filled 3D sub-shell provides additional stability, resulting in this unusual configuration.

Such anomalies arise due to the subtle balance between exchange energy, electron-electron repulsion, and other quantum effects. While the Aufbau principle serves as a reliable guideline, these exceptions are important for understanding the chemistry of some elements.

Importance of the Aufbau principle

The Aufbau principle is essential in chemistry for several reasons:

  • Predicting chemical properties: Understanding electron configuration helps chemists predict the chemical behavior of elements and their possible reactions.
  • Formation of chemical bonds: The electrons in the outermost shell, or valence electrons, are involved in chemical bonding. The Aufbau principle helps to identify these electrons.
  • The position of the elements in the periodic table and their properties can often be explained through their electron configuration.

Visualizing the Aufbau principle

Consider an analogy to understand the Aufbau principle. Think of electrons as filling seats in a theater, where certain seats are preferred because of better view or comfort (lower energy). Electrons naturally fill the best available seats first and then move to less favorable seats. This is similar to the way they fill orbitals, following the rules of the principle.

1s 2s 2P 3s 3P

Conclusion

The Aufbau principle is the cornerstone of electronic structure theory in chemistry. Although it simplifies the complex interactions within the atom, it provides a practical approach to arranging electrons in atoms. With exceptions noted in specific elements, the theory provides insight into electron configurations, helping us gain a deeper understanding of chemical properties, bonding behavior, and trends observed in the periodic table.


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Aufbau principle

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