Chemical Bonding Concepts (Part 1)

Ian introduces the concept of chemical bonding, explaining it as the attractive forces between particles such as atoms, ions, or molecules that result in a lower energy state. He describes the process of bond formation involving the redistribution of valence electrons and uses the example of two hydrogen atoms to illustrate how potential energy changes with the distance between atoms, defining the ideal bond length or equilibrium bond length. Ian also introduces arrow pushing diagrams to depict bond formation, particularly in the context of organic chemistry, and discusses electronegativity as a measure of an atom's ability to attract shared electrons, noting its increase across the periodic table.

This paragraph delves into the nature of covalent and ionic bonds, emphasizing the spectrum between purely covalent to purely ionic bonds. Ian explains that covalent bonds occur when there is no difference in electronegativity, as seen in homonuclear diatomic molecules, while ionic bonds form when there is a significant difference, leading to the transfer of electrons from one atom to another. Polar covalent bonds fall in the middle of this spectrum, with electrons being more attracted to the more electronegative atom, resulting in partial charges. The paragraph also covers the concept of polarizing power and how it influences the degree of covalent character in ionic bonds, using the transition from sodium chloride to aluminum chloride as an example.

Ian discusses exceptions to the octet rule, such as molecules with fewer or more than eight electrons in their valence shell. He explains that molecules like boron trichloride and aluminum chloride can have less than an octet due to their inability to form enough covalent bonds, leading to unique chemical reactivity. On the other hand, hypervalent molecules, which are possible for elements in period three and beyond, have more than eight electrons in their valence shell. This is achieved by the expansion of the octet using energetically low-lying d orbitals, as exemplified by PCl5, where the phosphorus atom promotes an electron to form five bonds instead of the typical four.

The VSEPR theory is introduced as a method to predict molecular geometries based on the arrangement of electron pairs around a central atom. Ian outlines the steps to deduce molecular geometry using VSEPR, which involves counting the regions of electron density and arranging them to minimize repulsion. He describes various geometries such as linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral, depending on the number of electron pairs. The paragraph also touches on the influence of lone pairs and the electronegativity of the central atom on bond angles, providing examples to illustrate these concepts.

Ian shifts the perspective to orbitals, explaining that bonding can be viewed as the overlap of atomic orbitals, leading to the formation of sigma and pi bonds. He discusses the formation of these bonds through the overlap of s, p, and hybrid orbitals, and how this relates to the observed bond angles and molecular geometries. The concept of hybridization is explored in the context of methane (CH4), ethene (C2H4), and acetylene (C2H2), detailing how carbon atoms use different hybrid orbitals (sp3, sp2, and sp) to form the appropriate number of bonds and achieve the observed molecular shapes.