Early Theories of Matter

Early Theories of Matter

• Democritus believed matter was composed of indivisible particles called atoms, which were uniform, solid, hard, incompressible, and indestructible. Atoms moved in infinite numbers through empty space until stopped.
• Aristotle proposed that all materials on Earth were made of four elements: Earth, Fire, Water, and Air. He believed all substances were made of small amounts of these four elements.
• Dalton's Atomic Theory proposed that:
  • All matter is composed of extremely small particles called atoms.
  • Atoms of a given element are identical in size, mass, and other properties.
  • Atoms of different elements differ in size, mass, and other properties.
  • Atoms cannot be subdivided, created, or destroyed.
• Thomson discovered electrons, which were negatively charged subatomic particles found within all atoms.
• Nagaoka proposed an atomic model with a small nucleus surrounded by a ring of electrons.
• Rutherford discovered that the positive charge and most of the mass of an atom are concentrated in an extremely small volume called the nucleus.
• Chadwick discovered neutrons, which are neutral subatomic particles found in the nucleus. He shot particles through gold to make this discovery

Subatomic Particles

• Proton: Located in the nucleus, with a relative mass of 1 and a charge of 1+.
• Neutron: Located in the nucleus, with a relative mass of 1 and a charge of 0.
• Electron: Located in orbitals around the nucleus, with a relative mass of 1/1838.65 and a charge of 1-.

Atomic Structure and Terminology

• Atomic Number (Z): The number of protons in an atom's nucleus.
• Mass Number (A): The sum of protons and neutrons in an atom's nucleus.
• Isotopes: Atoms of the same element that have the same number of protons but different numbers of neutrons.
• Relative Atomic Mass (Aᵣ): The weighted average of the masses of all naturally occurring isotopes of an element.
• Unified Atomic Mass Unit (u): A unit of mass used to express atomic masses, defined as 1/12 the mass of a carbon-12 atom.
• Atomic radius: The distance from an atom's nucleus to the furthest orbital of electrons.
• Shielding effect: The decrease in the nucleus's force of attraction on valence electrons due to the existence of electrons in the inner shells.
• Ionic radius: The distance from the nucleus of an ion up to which it has an influence on its electron cloud.

Properties of Ionic and Covalent Compounds

Intermolecular forces

• London dispersion force: Weakest intermolecular force, occurs between all molecules, involves temporary dipoles.
• Dipole-dipole force: Occurs between polar molecules due to permanent dipoles.
• Hydrogen bonding: Strongest intermolecular force, occurs between molecules containing a hydrogen atom bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine).

Ionization Energy

• Ionization Energy: The amount of energy required to remove an electron from an isolated atom or molecule.
• First ionization energy: The energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gas phase.
• Second ionization energy: The energy required to remove the outermost, or least bound, electron from a 1+ ion of the element.

Electron Affinity

• Electron affinity: The change in energy (in kJ/mol) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion.

Electronegativity

• Electronegativity: A chemical property that describes the tendency of an atom or a functional group to attract electrons toward itself.
• Trends in electronegativity:
  • Increases across a period due to increasing nuclear charge.
  • Decreases down a group due to increasing atomic radius.

Other Important Concepts

• Delta EN: The difference in electronegativity between two atoms in a bond.
  • ΔEN = 0 (non-polar covalent bond)
  • ΔEN < 1.7 (polar covalent bond)
  • ΔEN > 1.7 (ionic bond)
• Diatomic molecule: A molecule composed of two atoms.
• Polyatomic molecule: A molecule composed of three or more atoms.
• Lone pair: A pair of valence electrons that are not involved in bonding.
• Octet rule: The tendency of atoms to gain, lose, or share electrons to achieve a stable configuration with eight valence electrons.
• Structural formula: A diagram showing the arrangement of atoms and bonds within a molecule.
• Bonding capacity: The number of chemical bonds an atom can form.

Explore the foundational theories of matter through the insights of key philosophers and scientists. From Democritus's indivisible atoms to Dalton's atomic theory and the discoveries by Thomson and Rutherford, this quiz covers the evolution of our understanding of matter. Test your knowledge on these historical concepts and their significance in science.