Atomic Structure

Structure of the Atom

Atoms are the fundamental units of matter, composed of three main subatomic particles: protons, neutrons, and electrons. Each particle has distinct properties and occupies specific regions within the atom.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle also located in the nucleus.

• Electron: Negatively charged particle found in energy levels surrounding the nucleus.

Atomic Number and Mass Number

The atomic number (Z) is the number of protons in the nucleus and defines the element. The mass number (A) is the sum of protons and neutrons in the nucleus.

• Atomic Number (Z): Number of protons (and electrons in a neutral atom).

• Mass Number (A): Number of protons + number of neutrons.

Isotopes

Isotopes are atoms of the same element with the same atomic number but different mass numbers due to varying numbers of neutrons.

• Example: Carbon-12 and Carbon-13 are isotopes of carbon.

Mass Spectrometry and Relative Masses

Relative Atomic and Isotopic Mass

The relative atomic mass (Ar) is the weighted mean mass of an atom compared to 1/12 of the mass of a carbon-12 atom. The relative isotopic mass is the mass of a specific isotope relative to 1/12 of the mass of carbon-12.

Mass Spectrometer: Principle and Operation

A mass spectrometer measures the masses of positive ions formed from atoms and molecules. The sample must be in the gaseous state and undergoes several steps:

1. Sample is vaporized and injected.

2. Ionization: Bombardment with electrons forms positive ions.

3. Acceleration: Positive ions are accelerated by an electric field.

4. Deflection: Ions are deflected by a magnetic field according to their mass-to-charge ratio (m/z).

5. Detection: Ions are detected and counted at each m/z value.

Interpreting Mass Spectra

The mass spectrum displays the relative abundance of ions (y-axis) versus their mass/charge ratio (x-axis). The charge is usually +1, so the ratio is typically the mass.

• Relative abundance: Percentage of each isotope or ion detected.

• Calculation of Ar: Weighted average based on abundance and mass of each isotope.

Atomic Orbitals and Electronic Configurations

Energy Levels and Quantum Shells

Electrons are arranged in energy levels (shells) around the nucleus, each defined by a principal quantum number (n). The first shell (n=1) is closest to the nucleus and has the lowest energy.

Sub-shells and Orbitals

Each shell is divided into sub-shells (s, p, d, f), which contain orbitals. An orbital is a region where there is a high probability (90%) of finding an electron.

• s sub-shell: 1 orbital, spherical shape, holds 2 electrons.

• p sub-shell: 3 orbitals, dumbbell shape, holds 6 electrons.

• d sub-shell: 5 orbitals, clover shape, holds 10 electrons.

• f sub-shell: 7 orbitals, double dumbbell, holds 14 electrons.

Orbital Shapes

The s orbital is spherical, p orbitals are dumbbell-shaped, d orbitals are clover-shaped, and f orbitals are more complex. The size and energy of orbitals increase with quantum number.

Electronic Configuration

Electronic configuration describes the arrangement of electrons in each sub-shell and energy level. Electrons fill the lowest energy orbitals first (Aufbau principle), but there are exceptions due to energy differences between orbitals.

Hund's Rule and Pauli Exclusion Principle

Hund's rule states that electrons occupy orbitals singly before pairing. The Pauli Exclusion Principle states that two electrons in the same orbital must have opposite spins.

• Example: Electronic configurations for B, C, N, O, F, Ne.

Ionisation Energies

Definition and Types

Ionisation energy (\( \Delta H_i \)) is the energy required to remove an electron from an atom in the gaseous state. It is measured in kJ/mol.

• First ionisation energy: Energy to remove one electron from each atom in one mole of gaseous atoms.

• Second ionisation energy: Energy to remove one electron from each singly charged positive ion.

• Third ionisation energy: Energy to remove one electron from each doubly charged positive ion.

Successive Ionisation Energies

Successive ionisation energies increase as more electrons are removed, especially when moving to a new shell closer to the nucleus.

Factors Affecting Ionisation Energy

• Atomic radius: Larger radius lowers ionisation energy.

• Nuclear charge: Higher charge increases ionisation energy.

• Number of inner shells (shielding): More inner shells decrease ionisation energy due to increased shielding.

Trends in Ionisation Energies

Across a Period

Ionisation energy generally increases across a period due to increasing nuclear charge and decreasing atomic radius, with no change in shielding.

• Rapid decrease between last element of one period and first of next due to increased distance and shielding.

• Drop between Group 2 and 3 elements (e.g., Be and B) due to higher energy and shielding of 2p electrons in B.

• Drop between Group 5 and 6 elements (e.g., N and O) due to spin-pair repulsion in O.

Down a Group

First ionisation energy decreases down a group due to increasing atomic radius and shielding effect, despite increasing nuclear charge.

• This trend is repeated in Groups 2, 5, 6, 7, and 8.

The Periodic Table

Structure of the Periodic Table

The Periodic Table is organized into vertical columns called groups and horizontal rows called periods. Elements in the same group have similar chemical properties due to the same number of outer electrons.

• Groups: Same number of electrons in outer shell.

• Periods: Period number equals number of occupied shells.

• Blocks: s-block (groups 1, 2), p-block (groups 3-8), d-block (transition metals).

Periodic Properties

Periodic properties are regularly repeating patterns of atomic, physical, and chemical properties explained by electron configurations.

Atomic Radii

Atomic radius is the distance from the nucleus to the boundary of the electron cloud. It can be measured as covalent radius (diatomic molecules), van der Waals radius (noble gases), or metallic radius (adjacent atoms in metals).

Atomic radius decreases across a period due to increasing nuclear charge, which pulls electrons closer, counterbalancing electron-electron repulsion.

Melting and Boiling Points

Elements with giant lattice structures (e.g., metals, diamond) have high melting and boiling points, while those with simple molecular structures have low values.

First Ionisation Energies

First ionisation energies show periodic trends, increasing across a period and decreasing down a group.

Additional info: Where tables or diagrams were incomplete, logical academic context was added to ensure clarity and completeness. All equations are provided in LaTeX format as required.