Electronic Configuration and the Periodic Table - SS1 Chemistry Lesson Note
Electronic configuration and the periodic table are interconnected concepts that form the basis for understanding the properties and trends of elements. The arrangement of electrons within atoms determines their chemical behaviour, and this information is organised systematically in the periodic table. Let us delve into the intricacies of electronic configuration and its relationship with the periodic table:
Electron Energy Levels and Shells:
Electrons in atoms occupy specific energy levels or shells around the nucleus. The energy levels are labelled as 1, 2, 3, and so on, starting from the innermost shell. Each shell can hold a limited number of electrons:
- The first shell can hold a maximum of 2 electrons.
- The second shell can hold a maximum of 8 electrons.
- The third shell can hold a maximum of 18 electrons.
- Subsequent shells can accommodate even larger numbers.
Subshells and Orbitals:
Within each energy level, electrons occupy subshells of orbitals. There are four types of orbitals: s, p, d, and f orbitals. Each orbital has a specific shape and orientation:
- The s orbital is spherical and can hold a maximum of 2 electrons.
- The p orbital is dumbbell-shaped and can hold a maximum of 6 electrons (in three orbitals: px, py, and pz).
- The d orbital is cloverleaf-shaped and can hold a maximum of 10 electrons (in five orbitals).
- The f orbital is complex and can hold a maximum of 14 electrons (in seven orbitals).
Pauli Exclusion Principle and Hund's Rule:
The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers. This means that each orbital can accommodate a maximum of two electrons, with opposite spins (up and down).
Hund's Rule states that electrons prefer to occupy different orbitals within a subshell before pairing up. This results in unpaired electrons in different orbitals, which contributes to the atom's magnetic properties.
Notation of Electronic Configuration:
The electronic configuration of an atom is represented using a shorthand notation. It follows the pattern of specifying the energy level and the number of electrons in each subshell. For example, the electronic configuration of oxygen (atomic number 8) is 1s² 2s² 2p⁴, indicating that it has two electrons in the 1s orbital, two electrons in the 2s orbital, and four electrons in the 2p orbitals.
Periodic Table and Electronic Configuration:
The periodic table organises elements based on their atomic number (number of protons) and groups them according to their similar electronic configurations and chemical properties.
- Periods: Horizontal rows in the periodic table represent periods. Each period corresponds to a new energy level being filled with electrons.
- Groups: Vertical columns in the periodic table represent groups or families. Elements within the same group have similar outer electron configurations, which contribute to their similar chemical properties.
Periodic Trends:
Electronic configuration plays a vital role in determining periodic trends, such as atomic size, ionisation energy, electron affinity, and electronegativity.
- Atomic Size: Atomic size generally increases from top to bottom within a group and decreases from left to right across a period. This is due to the increasing number of energy levels and effective nuclear charge, respectively.
- Ionization Energy: Ionization energy is the energy required to remove an electron from an atom. It generally decreases from top to bottom within a group and increases from left to right across a period. This trend is influenced by the distance of the outermost electrons from the nucleus and the effective nuclear charge.
- Electron Affinity: Electron affinity is the energy change that occurs when an atom gains an electron. It generally increases from left to right across a period and decreases from top to bottom within a group.
- Electronegativity: Electronegativity is the tendency of an atom to attract electrons towards itself in a chemical bond. It generally increases from left to right across a period and decreases from top to bottom within a group.
Anomalies in Electronic Configuration:
Some elements exhibit anomalous electronic configurations, particularly transition metals and lanthanides/actinides, due to the stability gained by half-filled or filled subshells. These anomalies occur to minimise energy and achieve greater stability.
The periodic table organised elements based on their electronic configurations and allows us to identify trends and patterns in their behaviour. The relationship between electronic configuration and the periodic table forms the foundation of chemistry, enabling scientists to predict and explain the properties and reactions of various elements.