Inorganic chemistry

Avogadro's Law

• States that equal volumes of gases at the same temperature and pressure contain equal numbers of moles or molecules.
• The volume of a gas is directly proportional to the number of moles of gas present (V & n) when temperature and pressure are constant.
• One mole of any gas at a given temperature and pressure has the same fixed volume, called the molar gas volume.
• At standard temperature and pressure (STP), one mole of any gas occupies a volume of 22.4 liters.

Dalton's Law of Partial Pressures

• States that the total pressure of a mixture of gases is equal to the sum of the partial pressures of all the gases present.
• Each component gas in a mixture exerts a pressure as if it were alone in the container.
• The individual pressure of each gas in the mixture is defined as its partial pressure.
• The law is expressed as P = P₁ + P₂ + P₃ + ... (V and T are constant), where P is the total pressure and P₁, P₂, P₃, etc. are the partial pressures of each gas.

Democritus' Atomic Theory

• Proposed by the Greek philosopher Democritus around 400 BC.
• Suggested that all matter consists of indivisible particles called atoms.
• Believed that atoms were solid, had no internal structure, and differed in size, weight, and shape.

John Dalton's Atomic Model

• The first modern atomic model, building upon Democritus' ideas.
• Key postulates:
  • All matter consists of tiny particles called atoms.
  • Atoms are indestructible and unchangeable.
  • Elements are characterized by the weight of their atoms.
  • When elements react, their atoms combine to form new compounds.

J.J. Thomson's Discovery of the Electron

• Using cathode ray tube technology, Thomson discovered the electron in the 1890s.
• His model suggested that atoms were positively charged spheres with negatively charged electrons embedded within.

Erwin Schrödinger's Atomic Model

• Utilizing Heisenberg's uncertainty principle, Schrödinger developed the modern atomic model in the 1920s.
• Known as the cloud model, it suggests that electrons do not travel in exact orbits, but rather in regions of probability where they are likely to be found.

Atomic Structure Summary

• An atom is the smallest part of an element.
• Each atom of an element has a unique structure.
• Atoms can be divided into smaller subatomic particles: protons, electrons, and neutrons.

Quantum Numbers

• Describe the location and properties of electrons within an atom or ion.
• Four types of quantum numbers:
  • n (principal): determines the energy level of an electron.
  • l (orbital): determines the shape of the orbital.
  • ml (magnetic): designates a particular suborbital.
  • s (spin): describes the intrinsic angular momentum of an electron (spin up or spin down).

Pauli Exclusion Principle

• States that no two electrons within an atom can have the same four quantum numbers.
• This principle explains why electrons occupy different energy levels, orbitals, and spin states.

Types of Atomic Orbitals

• The most probable area to find an electron takes on a specific shape.
• Four main types of atomic orbitals:
  • s: spherical.
  • p: dumbbell shaped.
  • d: more complex shapes.
  • f: even more complex shapes.

Writing Lewis Structures with the Octet Rule

• Lewis structures are diagrams representing the bonding and lone pairs of electrons surrounding an atom.
• The octet rule states that atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight electrons in their outermost shell.
• Steps to write Lewis structures:
  • Determine the total number of valence electrons.
  • Draw a skeleton structure with single bonds between atoms.
  • Distribute remaining electrons as lone pairs to complete octets.

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) Theory predicts the shape of molecules based on the repulsion between electron pairs around a central atom.
• The steric number refers to the total number of electron pairs surrounding an atom, including bonding and lone pairs.
• The AXE method uses the symbols A, X, and E to represent the central atom, sigma bonds, and lone pairs, respectively.
• VSEPR theory predicts molecular geometries based on the steric number and the distribution of X and E.

Exceptions to VSEPR Theory

• Some compounds, especially transition metal complexes, do not have geometries predicted by VSEPR.
• This is due to the interaction of core d-electrons with ligands.
• Other theories, like VALBOND Theory and Crystal Field Theory, are used to predict the geometry of these exceptions.

Hybridization involving d-orbitals

• Hybridization refers to the mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies.
• d-orbitals can participate in hybridization, resulting in different geometries and bonding characteristics.
• Examples include sp3d-hybridization in PF5 and sp3d2-hybridization in SF6.

Octahedral Geometry

• Occurs in molecules with sp3d2-hybridization of the central atom.
• Six identical ligands are arranged around the central atom in a symmetric octahedral shape.

Square Planar Geometry

• Occurs in molecules with dsp2-hybridization of the central atom.
• Four identical ligands are arranged around the central atom in a square planar shape.