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Why Hydrogen Bonds Stronger Than Dipole Dipole
An Analysis of Hydrogen Bonds Stronger than Dipole-Dipole Interaction
On the theory of intermolecular forces, there are many discussions on the strength of hydrogen bonds and dipole-dipole interactions. The strength of the hydrogen bond is stronger than that of the dipole-dipole interaction. The analysis is as follows.
First on hydrogen bonds. For hydrogen atoms, their electron clouds tend to be biased towards electronegative atoms connected to them, such as fluorine, oxygen, nitrogen, etc. When hydrogen bonds with electronegative large atoms, hydrogen is almost in a bare proton state with a partial positive charge, and the exposed hydrogen then has a strong attraction to the lone pair electrons of electronegative large atoms in adjacent molecules. This attraction force is called hydrogen bond.
Subscopic dipole-dipole interaction. In a molecule, if the positive and negative charge centers do not coincide, it becomes a dipole. Between adjacent dipole molecules, the positive and negative terminals attract each other, which is a dipole-dipole interaction. However, this effect is only based on the uneven charge distribution of the molecule as a whole, and its mutual attraction force is weaker than that of hydrogen bonds.
The source of the difference in strength is that the nearly naked proton state of hydrogen in the hydrogen bond makes some of its positive charges concentrated and exposed, and it has a strong attraction to electronegative large atoms to electrons. In the dipole-dipole interaction, the charge distribution is relatively scattered and the attraction is weak. Second, when hydrogen bonds are formed, the interaction between specific atoms (fluorine, oxygen, nitrogen, etc.) and hydrogen has a certain directionality and saturation, which makes the hydrogen bond more stable; the directionality and saturation of the dipole-dipole interaction are not so obvious, so they are weaker.
In summary, hydrogen bonds are stronger than dipole-dipole interactions due to their unique charge distribution and interaction characteristics. Knowing this, the understanding of intermolecular forces is further advanced.
On the theory of intermolecular forces, there are many discussions on the strength of hydrogen bonds and dipole-dipole interactions. The strength of the hydrogen bond is stronger than that of the dipole-dipole interaction. The analysis is as follows.
First on hydrogen bonds. For hydrogen atoms, their electron clouds tend to be biased towards electronegative atoms connected to them, such as fluorine, oxygen, nitrogen, etc. When hydrogen bonds with electronegative large atoms, hydrogen is almost in a bare proton state with a partial positive charge, and the exposed hydrogen then has a strong attraction to the lone pair electrons of electronegative large atoms in adjacent molecules. This attraction force is called hydrogen bond.
Subscopic dipole-dipole interaction. In a molecule, if the positive and negative charge centers do not coincide, it becomes a dipole. Between adjacent dipole molecules, the positive and negative terminals attract each other, which is a dipole-dipole interaction. However, this effect is only based on the uneven charge distribution of the molecule as a whole, and its mutual attraction force is weaker than that of hydrogen bonds.
The source of the difference in strength is that the nearly naked proton state of hydrogen in the hydrogen bond makes some of its positive charges concentrated and exposed, and it has a strong attraction to electronegative large atoms to electrons. In the dipole-dipole interaction, the charge distribution is relatively scattered and the attraction is weak. Second, when hydrogen bonds are formed, the interaction between specific atoms (fluorine, oxygen, nitrogen, etc.) and hydrogen has a certain directionality and saturation, which makes the hydrogen bond more stable; the directionality and saturation of the dipole-dipole interaction are not so obvious, so they are weaker.
In summary, hydrogen bonds are stronger than dipole-dipole interactions due to their unique charge distribution and interaction characteristics. Knowing this, the understanding of intermolecular forces is further advanced.
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