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In 1940, Sidgwick and H.M. Powell proposed the valence shell electron pair repulsion rule, which states that all bonding and non-bonding electron pairs around a central atom will occupy positions on a sphere that are as far apart from each other as possible in order to minimize mutual repulsion. Later, R.J. Gillespie and R.S. Nyholm developed this rule further by applying the concepts of charge correlation and spin correlation, turning it into a more practical rule. The spatial geometric configuration of a covalent molecule depends primarily on the geometric arrangement of the valence electron pairs surrounding the central atom, and this arrangement in turn depends mainly on the number of electron pairs in the central atom’s valence shell. The basic points of this rule are: (1) The valence electron pairs of the central atom include only those that form σ bonds and the lone electron pairs in the valence shell. (2) The valence shell electron pairs repel each other and tend to move as far apart as possible, but at the same time they are all attracted to the atomic nucleus of the central atom. When these two effects reach equilibrium, it determines the positions of the electron pairs in each valence shell; the orientation of some of these positions is the direction in which bonds are formed with other atoms. Thus, the spatial configuration of the molecule or its shape is determined. When the number of valence electron pairs on the central atom is 2, the geometric configuration is linear; it is trigonal planar when the number is 3; tetrahedral when it is 4; trigonal bipyramidal when it is 5; and octahedral when it is 6, and so on. (3) The valence shell lone pair is attracted only by the central atom and remains close to it, whereas the σ-bond electron pair is also attracted by the other bonding atom and is thus farther away from the central atom. The closer the valence electron pairs are to the central atom, the greater the repulsion between them, and the larger the angle between their directions (or bond angle). Therefore, the repulsion (angle) between lone electron pairs > the repulsion (angle) between lone pairs and bonding electron pairs > the repulsion (angle) between bonding electron pairs. (4) When the electronegativities of the atoms bonded to the central atom differ, it results in the bonding electron pair being at different distances from the central atom. The greater the electronegativity, the farther the bonding electron pair is from the central atom; this reduces the repulsion between the electron pairs, and accordingly, the angle between them also decreases. The number of σ-bond electron pairs in an ABn-type molecule or ion is equal to the number of coordinating atoms bonded to the central atom, which is the value of n in ABn. The number of lone electron pairs can be calculated using the following formula: Number of lone electron pairs = (valence electrons of the central atom – total number of unpaired electrons in the coordinating atoms) / 2. If the result is a decimal value (such as 1.5), it should be rounded up to an integer (in this case, 2). For ABn-type ions, if it is a negative ion, the molecular term must be increased by the magnitude of the negative charge; if it is a positive ion, the molecular term must be decreased by the magnitude of the positive charge, and then divided by 2. For example, in the SO2 molecule, the number of σ bond electron pairs is 2; the number of lone electron pairs is 6 – 2×2/2 = 1. Thus, the total number of valence shell electron pairs is 2 + 1 = 3. Due to the presence of one pair of lone electron pairs, the geometry of the SO2 molecule is trigonal planar, and its molecular shape is angular, which corresponds to unequal sp2 hybridization according to hybridization theory. The above method is simplified and applicable to general ABn-type molecules or ions. However, it cannot indicate the specific bonding situation; it is merely an empirical rule. Therefore, this method is usually used first to determine the spatial configuration, after which the corresponding hybridization type is used to analyze the bonding situation.