Atomic Structure and Subatomic Particles

Fundamental Particles of the Atom

Atoms are composed of three fundamental subatomic particles: protons, neutrons, and electrons. Each particle has distinct properties and plays a unique role in atomic structure.

• Proton: Located in the nucleus, has a relative mass of 1 and a charge of +1.

• Neutron: Located in the nucleus, has a relative mass of 1 and no charge (0).

• Electron: Located in orbitals around the nucleus, has a relative mass of 1/1840 and a charge of -1.

Atomic Number (Z): The number of protons in the nucleus, which defines the element.

Mass Number (A): The total number of protons and neutrons in the atom.

Number of Neutrons: Calculated as A - Z.

Isotopes

Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons. Isotopes have similar chemical properties due to identical electronic structures, but their physical properties may vary because of differences in mass.

• Relative Isotopic Mass: The mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12.

• Relative Atomic Mass (R.A.M): The average mass of one atom compared to one twelfth of the mass of one atom of carbon-12.

• Relative Molecular Mass: The average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12.

Mass Spectrometry and Isotope Analysis

Principles of Mass Spectrometry

Mass spectrometry is a technique used to determine the isotopic composition of elements and molecules. It measures the mass-to-charge ratio (m/z) and abundance of each isotope, allowing calculation of the relative atomic mass.

• Calculation of R.A.M:

• If relative abundance is used instead of percentage:

Mass spectrometry is used in planetary exploration, drug testing, quality control in pharmaceuticals, and radioactive dating.

Mass Spectra of Diatomic Molecules

Elements with multiple isotopes (e.g., Cl and Br) produce characteristic spectra for their diatomic molecules, reflecting combinations of isotopes.

Measuring Molecular Mass (Mr)

When molecules are analyzed, the peak with the largest m/z corresponds to the molecular ion, representing the Mr of the molecule.

Ionisation Energies

Definition and Patterns

Ionisation energy is the energy required to remove an electron from an atom or ion in the gaseous state. The first ionisation energy refers to removing one electron from a neutral atom, while successive ionisation energies involve removing additional electrons.

• First Ionisation Energy:

• Second Ionisation Energy:

Three main factors affect ionisation energy:

1. Nuclear Attraction: More protons increase attraction.

2. Distance from Nucleus: Electrons further from the nucleus are less tightly held.

3. Shielding: Inner electrons repel outer electrons, reducing nuclear attraction.

Successive Ionisation Energies

Successive ionisation energies increase as more electrons are removed, due to increased attraction on remaining electrons. Large jumps in ionisation energy indicate removal from a new, inner shell.

Trends in Ionisation Energy Across Periods and Groups

Ionisation energy shows periodicity, with repeating patterns across periods. It generally increases across a period due to increasing nuclear charge and decreases down a group due to increased distance and shielding.

• Helium: Highest first ionisation energy due to proximity to nucleus and no shielding.

• Down a Group: Ionisation energy decreases as electrons are further from the nucleus and more shielded.

• Across a Period: Ionisation energy increases due to greater nuclear charge and similar shielding.

• Small Drops: Occur between Mg and Al (3p subshell starts filling) and P to S (electron pairing causes repulsion).

Electronic Structure and Models of the Atom

Energy Levels, Sublevels, and Orbitals

Electrons occupy principal energy levels (numbered 1, 2, 3, 4...) and sublevels (s, p, d, f). Each sublevel contains orbitals, which can hold up to two electrons of opposite spin.

• s sublevel: Holds up to 2 electrons, spherical shape.

• p sublevel: Holds up to 6 electrons, dumbbell shape.

• d sublevel: Holds up to 10 electrons.

• f sublevel: Holds up to 14 electrons.

Electrons fill sublevels in order of increasing energy: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p.

Electronic structure is written using numbers and letters (e.g., Oxygen: 1s2 2s2 2p4).

Spin diagrams use arrows to represent electrons with opposite spins in orbitals.

Electronic Structure for Ions

• Positive ions: Electrons are lost (e.g., Mg2+: 1s2 2s2 2p6).

• Negative ions: Electrons are gained (e.g., O2-: 1s2 2s2 2p6).

Periodicity and the Periodic Table

Classification of Elements: s, p, d, and f Blocks

The periodic table is divided into blocks based on the type of orbital being filled by the highest energy electrons:

• s block: Groups 1 and 2, plus helium.

• p block: Groups 13 to 18.

• d block: Transition metals (groups 3 to 12).

• f block: Lanthanides and actinides.

Trends Across Periods

Several properties show clear trends across periods:

• Atomic Radius: Decreases across a period due to increased nuclear charge pulling electrons closer.

• First Ionisation Energy: Generally increases across a period, with small drops due to subshell filling and electron pairing.

• Melting and Boiling Points: Vary depending on bonding type:

  • Na, Mg, Al: Metallic bonding, high melting/boiling points.

  • Si: Macromolecular, very high melting/boiling point.

  • Cl2, S8, P4: Simple molecular, low melting/boiling points due to weak London forces.

  • Ar: Monoatomic, very low melting/boiling point.

Similar trends are observed in period 2 (Li, Be, B, C, N, O, F, Ne).

Summary Table: Subatomic Particles

Summary Table: Periodic Table Blocks

Example: Magnesium (Mg) is an s block element with electronic configuration 1s2 2s2 2p6 3s2.

Additional info: The periodicity in ionisation energy, atomic radius, and melting/boiling points is fundamental to understanding chemical reactivity and bonding in elements.