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  2. Orbital hybridisation - Wikipedia

    en.wikipedia.org/wiki/Orbital_hybridisation

    Chemist Linus Pauling first developed the hybridisation theory in 1931 to explain the structure of simple molecules such as methane (CH 4) using atomic orbitals. [2] Pauling pointed out that a carbon atom forms four bonds by using one s and three p orbitals, so that "it might be inferred" that a carbon atom would form three bonds at right angles (using p orbitals) and a fourth weaker bond ...

  3. Isovalent hybridization - Wikipedia

    en.wikipedia.org/wiki/Isovalent_hybridization

    In chemistry, isovalent or second order hybridization is an extension of orbital hybridization, the mixing of atomic orbitals into hybrid orbitals which can form chemical bonds, to include fractional numbers of atomic orbitals of each type (s, p, d). It allows for a quantitative depiction of bond formation when the molecular geometry deviates ...

  4. Orbital overlap - Wikipedia

    en.wikipedia.org/wiki/Orbital_overlap

    In chemical bonds, an orbital overlap is the concentration of orbitals on adjacent atoms in the same regions of space. Orbital overlap can lead to bond formation. Linus Pauling explained the importance of orbital overlap in the molecular bond angles observed through experimentation; it is the basis for orbital hybridization.

  5. CRC Handbook of Chemistry and Physics - Wikipedia

    en.wikipedia.org/wiki/CRC_Handbook_of_Chemistry...

    Section 3: Physical Constants of Organic Compounds; Section 4: Properties of the Elements and Inorganic Compounds; Section 5: Thermochemistry, Electrochemistry, and Kinetics (or Thermo, Electro & Solution Chemistry) Section 6: Fluid Properties; Section 7: Biochemistry; Section 8: Analytical Chemistry; Section 9: Molecular Structure and Spectroscopy

  6. Homolysis (chemistry) - Wikipedia

    en.wikipedia.org/wiki/Homolysis_(chemistry)

    Orbital hybridization. The s-character of an orbital relates to how close electrons are to the nucleus. In the case of a radical, s-character more specifically relates to how close the single electron is to the nucleus. Radicals decrease in stability as they are closer to the nucleus, because the electron affinity of the orbital increases.

  7. Bent bond - Wikipedia

    en.wikipedia.org/wiki/Bent_bond

    The term itself is a general representation of electron density or configuration resembling a similar "bent" structure within small ring molecules, such as cyclopropane (C 3 H 6) or as a representation of double or triple bonds within a compound that is an alternative to the sigma and pi bond model.

  8. Hypervalent molecule - Wikipedia

    en.wikipedia.org/wiki/Hypervalent_molecule

    Using the language of orbital hybridization, the bonds of molecules like PF 5 and SF 6 were said to be constructed from sp 3 d n orbitals on the central atom. Langmuir, on the other hand, upheld the dominance of the octet rule and preferred the use of ionic bonds to account for hypervalence without violating the rule (e.g. " SF 2+

  9. Molecular geometry - Wikipedia

    en.wikipedia.org/wiki/Molecular_geometry

    Some common shapes of simple molecules include: Linear: In a linear model, atoms are connected in a straight line. The bond angles are set at 180°. For example, carbon dioxide and nitric oxide have a linear molecular shape. Trigonal planar: Molecules with the trigonal planar shape are somewhat triangular and in one plane (flat). Consequently ...