What Is Molecular Geometry?

The relative arrangement of each atom in a covalent compound molecule in space is called the spatial configuration or geometric configuration of the molecule. Covalent bonds have directionality. In molecules of the same composition, if the order and arrangement of atoms are different, different molecules can be formed. Through physical methods such as dipole moment measurement and electron diffraction, the relative positions of atoms in a molecule can be determined, and the bond length and angle can be determined to determine the molecular configuration.

Molecular geometric configuration is one of the basic problems of molecular structure and chemical bond theory. Understanding the molecular geometry is very important for understanding the relationship between the properties of a substance and its internal structure. Since 1916, from the point of view of electronics, Lewis first proposed that chemical bonds are caused by electron pairs. This electron theory is widely used to elucidate the molecular structure of matter. However, this theory treats electrons as stationary negative charges according to classical electrostatic theory, and will inevitably encounter many unsolvable contradictions. In 1927, Heitter (W.) and London (F.) first applied quantum mechanics theory to molecular structure. Later, Pauling (L.) and others developed this result. Established modern price key theory. In 1932. Millikan and Hundt proposed the molecular orbital theory from a different perspective than the valence bond theory. Later, in order to better explain the actual spatial configuration and stability of the molecule, Pauling (L.) proposed the "orbital hybrid theory" based on the "electron pairing" assumption. ^[1]

The method to determine the geometric configuration of simple polyatomic molecules by theory is to calculate the total energy of the molecule under different bond length bond angles. When the total energy is the lowest, the corresponding bond length bond angle is the stable geometric configuration of the molecule.

Lone pair effect

In the molecule, since the bonding electron pair is attracted by both the central atom and the coordinating atom, the electron cloud is relatively tight, while the lone electron pair is attracted by only one atomic nucleus, and the occupied orbit is closer to the nucleus, near the core Occupying a large space, the electronic cloud is more "hypertrophic". In this way, the repulsion effect of the lone electron pair on the electrons in the adjacent orbits is stronger than that in the bonding orbitals. Therefore, lone electron pairs tend to be as far away from each other as possible, and are squeezed into bonded electron pairs.

The lone electron pair effect can help us choose the stable configuration of the molecule, on the other hand, it can also be used to explain the cause of the molecular spatial configuration distortion. For example, in the CH ₄ , NH ₃ , and H ₂ O series, the solitary electron pairs are 0, 1, 2, and the bond angles are gradually reduced to 109.5 °, 107.5 °, and 104.5 °, respectively. The reason is the increase in the number of lone pairs of electron pairs, which are extruded into bond electron pairs.

Geometrical configuration

Heavy-bond orbits are more repulsive to other orbits than single-bond orbits. Due to the large space occupied by four electrons in a double bond or six electrons in a triple bond

Figure (2) Bond angles of molecules with heavy bonds

The space occupied by the two electrons in a single bond also has a large repulsive force against adjacent bonded electron pairs. Therefore, the angle of the bond angle containing heavy bonds is relatively large. For example, the single bond-double bond bond angle in COCl ₂ molecule is 124.3 °, while the single bond-single bond bond angle is 111.3 °.

Electronegative effect of molecular geometric configuration

The electronegativity between the central atom and the coordinating atom will also affect the molecular configuration data (P202-205). The greater the electronegativity of the coordinating atom bonded to the central atom, the stronger the ability to attract valence electrons, and the valence electrons will move toward the ligand, farther from the central atom, reducing the charge density near the central atom, thereby The repulsion between the keyed orbit and the adjacent orbit is reduced. That is, the repulsive force between the valence electron pair decreases with the increase of the electronegativity of the coordinating atom, and the generated bond angle is also smaller. For example, in the NF ₃ molecule, the bond angle is 102.1 °; in the NH ₃ molecule, the bond angle is 107.5 °. In the same way, when the coordination atoms are the same and the central atom is different, as the electronegativity of the central atom becomes smaller, the bond angle also decreases.

However, there are exceptions to this rule, such as the bond angle of PF ₃ (104 °)> the bond angle of PX ₃ (X = Cl, Br, I). The reason for this anomaly may be that F is a small atom of the second period and the charge density itself is already very high. The bonding electron pair of the P-F bond is biased towards the F atom when the bond is formed, which causes the charge density of F to increase sharply. Big. Therefore, the F atom may feed back P _y or P _z electrons to the d _xy or d _xz orbital of the P atom, reducing the charge density of the F atom, increasing the charge density of the P atom, and thereby increasing the P-F bond. The repulsive force between them makes the P-F key bond angle larger.

Number of adjacent electron pairs in the geometric configuration of the molecule

The number of adjacent electron pairs in the valence layer pairs has an effect on the bond length. Valence layer electron pairs with a large number of adjacent electron pairs are ranked by other surrounding electron pairs.

image 3)

The repulsive force is large, and it is far from the nucleus, that is, the generated bond is longer. For example, each upright electron pair of the AL ₅ molecule has three neighboring pairs at an angle of 90 °, while each flat electron pair has two neighboring pairs at an angle of 90 ° and two adjacent pairs at an angle of 120 °. Because the repulsive force between electron pairs at an angle of 120 ° is much smaller than the repulsive force of electron pairs at an angle of 90 °, the total repulsive force on an upright electron pair is greater than that of a flat electron pair, so the upright bond is longer than the flat bond, such as image 3).

The symmetry of AL ₂ , AL ₃ , AL ₄ , and AL ₆ molecules makes the number of adjacent electron pairs of each electron pair the same, so all bond lengths of these types of molecules are equal. For example: CO _2, BCl _3, CCl _4, SF ₆ and other molecules belong to this type of molecule. ^[3]