Electron Shells and Orbitals

Electron Shells & Orbitals

• First shell holds a maximum of 2 electrons.
• Second shell holds a maximum of 8 electrons.
• Third shell holds a maximum of 18 electrons.
• Fourth shell holds a maximum of 32 electrons.
• 's' subshells have 1 orbital and hold 2 electrons.
• 'p' subshells have 3 orbitals and hold 6 electrons.
• 'd' subshells have 5 orbitals and hold 10 electrons.
• 'f' subshells have 7 orbitals and hold 14 electrons.

Electron Configuration Rules

• The Aufbau Principle dictates filling the lowest energy orbitals first.
• The Madelung Rule determines the orbital filling order.
• Hund’s Rule states that electrons fill empty orbitals before pairing up to minimize repulsion.

Order of Filling Orbitals

• The filling order is 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p...
• Note that 4s fills before 3d because it has lower energy.

Shorthand Notation

• Use noble gases as a shortcut, e.g., Na (11 electrons) is [Ne] 3s¹.

Electrons in Reactions

• Valence electrons (outermost shell) determine reactivity.
• Inner shell electrons are not involved in bonding.
• Electrons in the same subshell fill one per orbital before pairing, according to Hund’s Rule.

Atomic Structure & Electron Shells

• Atoms consist of protons (+), electrons (-), and neutrons (0).
• Electrons reside in shells around the nucleus.
• Electron shells fill in a specific order, with the 1st shell holding 2 electrons, the 2nd shell holding 8 electrons, and the 3rd shell holding 8 (stable) or up to 18 electrons.

Electron Configuration Rules

• Lower energy orbitals are filled first, following the Aufbau Principle.
• Electrons fill empty orbitals before pairing up (Hund's Rule).
• The electron configuration format is, e.g., Carbon: 1s² 2s² 2p².
• Chromium and Copper break the Aufbau principle for extra stability.
• Electrons are lost from higher energy orbitals first when forming ions.

Bonding Basics

• Full valence shells provide stability like noble gases.
• Atoms gain, lose, or share electrons to achieve stability.
• Gaining electrons results in a negative ion (anion).
• Losing electrons results in a positive ion (cation).
• Valency refers to the number of bonds an atom can form, e.g., Carbon = 4.

Electronegativity & Bonding

• Electronegativity measures an atom’s ability to attract electrons.
• Electronegativity increases across a period (left to right).
• Electronegativity decreases down a group (top to bottom).
• Polar bonds occur when electrons are unequally shared, e.g., Water (O-H bond).

Easy Tricks

• "2-8-8 Rule" for the first 3 shells.
• Aufbau relates to order, and Hund relates to filling empty orbitals first.
• Fluorine has the highest electronegativity.
• Noble gases are stable due to full valence shells and tend not to form bonds.

First Ionisation Energy

• First Ionisation Energy is the minimum energy to remove one electron from each atom in one mole of gaseous atoms, represented as X(g) → X⁺(g) + e⁻.
• The distance from the nucleus affects ionisation energy; greater distance means weaker attraction and lower ionisation energy.
• Electron shielding (inner electrons blocking nuclear attraction) lowers ionisation energy.
• Greater nuclear charge (more protons) increases ionisation energy.

Trends in Ionisation Energy

• More electron shells result in increased shielding and distance, leading to lower ionisation energy down a group.
• More protons with the same shielding result in stronger nuclear attraction and higher ionisation energy across a period.

Orbitals & Ionisation Energy

• Electrons exist in orbitals, not fixed paths.
• Electrons in s-orbitals (closer to the nucleus) have higher ionisation energy than those in p-orbitals.
• E.g., Boron has lower ionisation energy than Beryllium because Boron's electron is in a p-orbital.

Second Ionisation Energy

• The energy needed to remove a second electron from a X⁺ ion, represented as X⁺(g) → X²⁺(g) + e⁻.
• Removing a second electron requires more energy due to stronger nuclear attraction.
• Ionisation energy greatly increases when removing an electron from a new shell.
• A large jump in ionisation energy happens after the 1st IE for Na and after the 2nd IE for Mg due to removing from a stable noble gas core.

Reactivity & Ionisation Energy

• Metals lose electrons easily.
• Group 1 metals are very reactive due to low 1st IE.
• Group 2 metals are less reactive than Group 1 due to higher 1st IE.
• Non-metals hold electrons tightly due to high ionisation energy.
• Oxygen has a much higher IE than calcium.

Atomic & Ionic Radius

• Atomic radius increases down a group due to more shells.
• Atomic radius decreases across a period due to stronger attraction.
• Cations (+) are smaller due to losing electrons and stronger attraction.
• Anions (-) are larger due to gaining electrons and weaker attraction.

Electron Shielding

• Inner electrons repel outer electrons, reducing nuclear attraction.
• More shielding leads to lower ionisation energy.
• Shielding stays constant across a period, but nuclear charge increases.
• Shielding increases down a group, reducing nuclear pull.

Key Mnemonics

• "Down is down, Across is up" for ionisation energy trends.
• "More layers, less pull" for electron shielding.
• "Bigger down, smaller across" for atomic radius.
• "Each electron removed = stronger hold" for ionisation energy.

Successive Ionisation Energies

• Each electron removed increases the attraction to the nucleus, making each successive ionisation energy higher.
• Large jumps occur when removing an electron from a new inner shell.
• Use ionisation energy graphs to determine the group number of an element.

Effective Nuclear Charge (Z_eff)

• Z_eff = Nuclear charge - Shielding electrons.
• Z_eff increases across a period (smaller atomic radius, higher ionisation energy).
• Z_eff stays the same down a group (larger atomic radius, lower ionisation energy).

Electron Configurations & Ionisation Energy

• Stable configurations lead to higher IE.
• Full & half-full subshells are more stable and harder to ionise.
• Nitrogen has a half-full p-subshell, resulting in higher ionisation energy than oxygen. And magnesium (3s²) > aluminum (3s² 3p¹) cause 3p is easier to remove.

Periodic Trends: Beyond Ionisation Energy

• Electronegativity is the ability to attract electrons in a bond.
• Electronegativity increases across a period and decreases down a group.
• Fluorine is the most electronegative element.
• Metals lose electrons easily, while non-metals hold electrons tightly.
• Across a period, elements transition from metals to metalloids to non-metals.
• Electron affinity measures the energy change when an atom gains an electron.

Exceptions to Ionisation Energy Trends

• Some exceptions exist due to subshell configurations.
• Group 2 elements (Be, Mg) have higher IE than Group 13 (B, Al) because the p-orbital is higher energy and easier to remove.
• Group 15 (N, P) have higher IE than Group 16 (O, S) because of paired electrons in the p⁴ orbital.