Question
1. Why do atoms join together to form bonded species?2. Why can a molecule be more stable than its separated atoms?3. As two atoms approach, which attractive and repulsive interactions act between them?4. Describe how the potential energy of two hydrogen atoms changes as an H₂ bond forms.
Answer
1. Why do atoms form bonds?
Atoms form a stable chemical bond when the bonded arrangement has a lower total energy than the relevant separated particles. The lower-energy arrangement is more stable.
Electron-counting patterns such as the octet rule can be useful for predicting many main-group structures, but the more general physical reason for a stable bond is the existence of an energetically favorable arrangement.
2. Why can a molecule be more stable than free atoms?
When bond formation lowers the potential energy of the system, energy is released as the bonded species forms. Returning the atoms to the separated state then requires energy input.
That is why a genuinely stable molecule lies in an energy minimum relative to the corresponding separated atoms.
3. What interactions act as two atoms approach?
Several electrostatic interactions occur at the same time:
- each electron is attracted to positively charged nuclei;
- the two nuclei repel one another;
- electrons repel one another.
At suitable separations, attractive interactions can outweigh the repulsive contributions enough to lower the system's energy. At very short separations, the repulsions rise strongly.
4. What happens to the potential energy as H₂ forms?
When two hydrogen atoms are very far apart, their interaction energy is conventionally taken as approximately zero. As they approach, each electron begins to interact with both nuclei and the potential energy decreases.
The energy reaches a minimum at an H–H separation of about 74 pm, the equilibrium bond length. That is the most stable separation.
If the nuclei are pushed closer than this, nucleus–nucleus and electron–electron repulsions increasingly outweigh the additional attraction, so the potential energy rises steeply.
Forming the H–H bond releases energy; breaking it requires energy. OpenStax gives an H–H bond dissociation energy of about 436 kJ/mol.
Evidence boundary
The detailed potential-energy curve in Part 4 describes H₂ covalent bonding. The broader lower-energy bonded state idea is general, but different bonding types require different electronic models. The octet rule is a useful main-group heuristic rather than a universal law of chemical bonding.
Sources
These references support the concepts and methods used in the explanation above.