Structure of Atom Quiz

Structure of Atom

Electron Configurations

• Electron configuration determines how electrons are distributed in an atom's orbitals.
• Follow the Aufbau principle (fill lower energy levels first).
• Use the Pauli exclusion principle (no two electrons can have the same set of quantum numbers).
• Hund’s rule states that electrons occupy degenerate orbitals singly before pairing up.
• Notation example: Carbon (C) has an electron configuration of 1s² 2s² 2p².

Subatomic Particles

• Protons:

  • Positively charged, found in nucleus.
  • Define the atomic number (Z) of an element.

• Neutrons:

  • Neutral charge, also located in nucleus.
  • Contribute to atomic mass but not charge.

• Electrons:

  • Negatively charged, orbiting the nucleus in electron clouds.
  • Equal in number to protons in a neutral atom.

Atomic Models

• Dalton’s Model:

  • Atoms as indivisible spheres; different elements have different atoms.

• Thomson’s Model (Plum Pudding Model):

  • Electrons embedded in a positively charged "soup."

• Rutherford’s Model:

  • Atoms have a dense nucleus with electrons orbiting around it; led to the understanding of the nucleus.

• Bohr’s Model:

  • Electrons in fixed orbits with quantized energy levels, explaining spectra of hydrogen.

• Quantum Mechanical Model:

  • Electrons exist in probabilistic orbitals rather than fixed paths; uses wave functions.

Isotopes and Ions

• Isotopes:

  • Atoms of the same element with different numbers of neutrons.
  • Same atomic number (Z) but different mass numbers (A).
  • Example: Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons).

• Ions:

  • Atoms that have gained or lost electrons, resulting in a net charge.

  • Cations: Positively charged ions (loss of electrons).

    • Example: Na⁺ (sodium ion).

  • Anions: Negatively charged ions (gain of electrons).

    • Example: Cl⁻ (chloride ion).

Electron Configurations

• Electron configuration describes the distribution of electrons within an atom's energy levels and orbitals.
• The Aufbau principle dictates filling lower energy levels before higher ones.
• The Pauli exclusion principle states that no two electrons can have the same four quantum numbers.
• Hund's rule says electrons occupy degenerate orbitals singly before pairing up.
• Example: Carbon (C) has an electron configuration of 1s² 2s² 2p².

Subatomic Particles

• Protons: Positively charged particles found in the nucleus, define the atomic number (Z) of an element.
• Neutrons: Neutral particles also located in the nucleus, contribute to the atom's mass but not its charge.
• Electrons: Negatively charged particles orbiting the nucleus in electron clouds, equal in number to protons in a neutral atom.

Atomic Models

• Dalton's Model: Atoms are indivisible spheres, with different elements having different atoms.
• Thomson's Model (Plum Pudding Model): Electrons are embedded in a positively charged "soup".
• Rutherford's Model: Atoms have a dense, positively charged nucleus with electrons orbiting around it, leading to the understanding of the nucleus.
• Bohr's Model: Electrons occupy fixed orbits with quantized energy levels, explaining the spectra of hydrogen.
• Quantum Mechanical Model: Electrons exist in probabilistic orbitals rather than fixed paths, described by wave functions.

Isotopes and Ions

• Isotopes: Atoms of the same element with different numbers of neutrons.
• Same atomic number (Z) but different mass numbers (A).
• Example: Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons).
• Ions: Atoms that have gained or lost electrons, resulting in a net charge.
• Cations: Positively charged ions (loss of electrons).
• Example: Na⁺ (sodium ion).
• Anions: Negatively charged ions (gain of electrons).
• Example: Cl⁻ (chloride ion).