Atomic Structure and Fundamental Particles

Fundamental Particles

• The model for atomic structure has evolved over time, with the current model consisting of a small, dense, positively charged nucleus surrounded by orbiting electrons in electron shells.
• This model was discovered through the Rutherford scattering experiment in 1911.
• The nucleus contains protons and neutrons, giving it a positive charge and holding almost all of the atom's mass.
• In a neutral atom, the number of protons and electrons is equal.

Determining the Number of Fundamental Particles

• Mass Number (A): Represents the sum of protons and neutrons in an atom.
• Atomic Number (Z): Represents the number of protons in an atom.
• Example: If Atomic Number = 7 and Mass Number = 14, then: Proton number = 7 and Neutron number = 14 - 7 = 7

Relative Atomic Mass and Isotopes

• Relative Atomic Mass (Ar): The mean mass of an atom of an element, divided by one twelfth of the mean mass of an atom of the carbon-12 isotope.
• Isotopes: Atoms of the same element with the same atomic number, but a different number of neutrons, resulting in a different mass number.
• Isotopes react chemically in the same way because their proton number and electron configuration is the same.
• However, isotopes have different physical properties due to their differing mass numbers.

Ions and Mass Spectrometry

• Ions: Formed when an atom loses or gains electrons, resulting in an overall charge.
• Mass Spectrometry: Analytical technique used to identify isotopes and determine the relative atomic mass of an element.
• Time of Flight (TOF) Mass Spectrometry: Records the time taken for isotopes to reach a detector, generating spectra displaying each isotope present.
• Steps in TOF Mass Spectrometry:
  • Ionization: Sample vaporized and injected into a chamber, where high voltage causes electrons to be removed, creating +1 charged ions.
  • Acceleration: Positively charged ions accelerated towards a negatively charged detection plate.
  • Ion Drift: Ions deflected by a magnetic field into a curved path, with radius depending on charge and mass.
  • Detection: Ions hit the detection plate, gain an electron, and produce a flow of charge. The greater the abundance, the greater the current.
  • Analysis: Current values and flight times are used to generate a spectra printout showing isotope abundance.

Interpretation of Mass Spectra

• A 2+ charged ion will be affected more by the magnetic field, creating a path of smaller radius.
• This results in a halved m/z (mass to charge ratio) and is represented by a trace at half the expected m/z value on spectra.

Chlorine Spectra

• Mass spectra of chlorine display a characteristic pattern due to the presence of two isotopes:
  • Cl+ ions: 3:1 ratio
  • Cl2+ ions: 3:6:9 ratio
• This ratio arises from the different combinations of chlorine isotopes in the molecule.

Electron Configuration

• Electron Orbitals: Electrons are held in clouds of negative charge called orbitals. Different types of orbitals exist: s, p, d, and f, each with a unique shape.
• Electron Orbitals and the Periodic Table: These orbitals correspond to blocks on the Periodic Table, with each element in the block's outer electrons residing in that orbital.
• Electron Orbital Capacity: Each type of orbital can hold a specific number of electrons:
  • s-orbital: 2 electrons
  • p-orbital: 6 electrons
  • d-orbital: 10 electrons
• Orbitals fill in order of increasing energy, from s to d, with each orbital completely filled before the next one is used.
• Electron Spin: Within orbitals, electrons pair up with opposite spin to achieve stability. This spin is represented by arrows.

Rules for Electron Configuration

• Lowest energy orbitals are filled first.
• Electrons with the same spin fill an orbital before pairing begins.

• No single orbital holds more than two electrons.

Exceptions to Electron Configuration Rules

• Unpaired Spins: Unpaired electrons with unbalanced spins create instability due to natural electron repulsion.
• Electron Rearrangement: Electrons may rearrange to achieve stability, leading to changes in electron configuration.

Ionisation Energy

• Definition: Minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state.
• Units: kJmol-1
• Successive Ionisation Energies: Ionization energies for removing subsequent electrons, typically requiring more energy due to increased electrostatic attraction between the nucleus and remaining electrons.
• Trends in First Ionisation Energy:
  • Along a Period: Ionisation energy increases due to decreasing atomic radius and greater electrostatic forces of attraction.
  • Down a Group: Ionisation energy decreases due to increasing atomic radius and shielding effects, reducing electrostatic force.
• Sudden Large Increase in Successive Ionisation Energies: Indicates a change in energy level as the electron is removed from a closer orbital, requiring more energy.
• Supporting Evidence for Atomic Orbital Theory: This large energy increase provides evidence for the atomic orbital theory.
• Aluminum Exception: Aluminum's first ionisation energy is lower than expected due to a single pair of electrons with opposite spins, causing repulsion and reducing the energy required to remove the outer electron.