topic 2

Molecules and Compounds

• A molecule consists of two or more atoms.

• A compound consists of two or more elements.

• Examples of Molecules:

  • Element molecules:

    • Hydrogen (H₂)

    • Oxygen (O₂)

  • Compound molecules:

    • Water (H₂O)

    • Carbon Dioxide (CO₂)

Allotropes

• Definition: Allotropes are different forms of the same element.

• Examples:

  • Oxygen (O₂) vs. Ozone (O₃)

  • Phosphorus: Red phosphorus vs. White phosphorus

  • Carbon: Diamond vs. Graphite

Structure of Atoms

• Importance of studying subatomic particles and atomic structure.

• Focus on the following:

  • Subatomic Particles:

    • Basic building blocks of atoms

  • Isotopes:

    • Atoms of the same element with differing neutron numbers

    • Examples of isotopes and their significance.

Electron Configuration

• Emphasis on how electrons are arranged in atoms.

• Understanding the electron configuration is critical for predicting chemical behavior.

Electron Orbits and Chemical Bonding

• Understanding electron orbits is crucial for:

  • Chemical bonding

  • Structure and properties of materials

• Ancient Greek atomic theory:

  • Atoms viewed as solid particles of defined sizes.

• Discovery of electron emission:

  • Electrons can be emitted from atoms, leading to new models.

• Early models of the atom:

  • Plum pudding model:

    • Atoms = solid mass with electrons stuck on the surface.

Discovery of the Atomic Structure

• Ernest Rutherford's experiment (1924):

  • Fired alpha particles at gold foil:

    • Most particles passed through.

    • Some were deflected, indicating a small, dense nucleus.

• Rutherford's conclusion:

  • An atom consists of a small, positive nucleus surrounded by space containing electrons.

Bohr Model of the Atom

• Niels Bohr's contributions:

  • Electrons are restricted to certain orbits around the nucleus.

• Current atomic model:

  • More sophisticated, incorporating quantum mechanics to explain atomic structure.

• Electron properties:

  • Electrons are negatively charged particles, evidenced by experiments with magnetic fields.

Subatomic Particles

• Three types of subatomic particles:

  • Protons (positive charge, mass = 1 unit)

  • Neutrons (neutral charge, mass approximately = 1 unit)

  • Electrons (negative charge, mass = 1/2000 unit, considered negligible)

• Electrical neutrality in atoms:

  • Equal number of protons and electrons in neutral atoms.

• Example of atomic structure:

  • Hydrogen (1 proton, 0 neutrons) is stable without neutrons.

Role of Neutrons

• Neutrons provide stability in atomic nuclei:

  • Protons repulse each other due to positive charge.

  • Neutrons help hold the nucleus together, especially in heavier elements.

• Radioactive decay occurs when an isotope has an imbalanced ratio of protons and neutrons.

• Stable nuclei contain a proper balance of protons and neutrons:

  • Maximum stable size around atomic number 83 (Bismuth).

Isotopes

• Definition of isotopes:

  • Atoms of the same element with different numbers of neutrons.

• Example of hydrogen isotopes:

  • Protium (0 neutrons)

  • Deuterium (1 neutron, stable)

  • Tritium (2 neutrons, radioactive)

• Carbon isotopes:

  • Carbon-12, Carbon-13, Carbon-14 (used in radiocarbon dating).

Atomic Emission Spectrum

• Evidence for electron orbits comes from the atomic emission spectrum:

  • Electrons move between energy levels when energy is supplied.

  • Each element emits distinct wavelengths of light when energized.

• Historical experiments:

  • Newton's prisms demonstrated light spectrum.

The Periodic Table and Electron Configuration

• Organization of the periodic table reflects element electron configurations:

  • Elements in the same group have similar chemical properties due to outer shell electron patterns.

• Electron shells:

  • Each shell can hold a limited number of electrons (2, 8, etc.).

• Reactivity is based on outermost electron configuration:

  • Elements that need few electrons to fill their outer shell tend to react similarly.

Formation of Ions

• An ion is a charged atom formed by:

  • Removing electrons (forming cations) or

  • Adding electrons (forming anions).

• Cations:

  • Positively charged ions formed when electrons are lost (e.g., Na+ from sodium).

• Anions:

  • Negatively charged ions formed when electrons are gained (e.g., F- from fluorine, called fluoride).

• Isoelectronic species:

  • Ions that have the same number of electrons as a noble gas, leading to stability.

  • Example: Na+ has 10 electrons (same as Ne).

Conclusion

• Understanding atomic structure and electron configurations:

  • Key to predicting chemical behavior and bonding.

• Stable electron arrangements are crucial for an element's reactivity and chemical properties.