The concept of electrons orbiting the nucleus of an atom in a manner similar to planets orbiting the sun is an outdated model known as the "planetary model" or "Rutherford model" of the atom. According to this model, electrons would indeed be expected to spiral into the nucleus due to the electromagnetic attraction between the negatively charged electrons and the positively charged protons.
However, this classical view of atomic structure was superseded by the development of quantum mechanics in the early 20th century. Quantum mechanics describes the behavior of particles at the atomic and subatomic level more accurately and explains why electrons do not simply collapse into the nucleus.
In the quantum mechanical model, electrons are not thought of as orbiting the nucleus in a classical sense. Instead, they exist in specific regions around the nucleus known as "atomic orbitals" or "electron shells." Each electron shell corresponds to a specific energy level, and within each shell, there are subshells and orbitals where electrons can be found.
Electrons occupy these orbitals based on their energy, and they exhibit wave-like properties. The behavior of electrons is described by mathematical equations called wave functions, and these functions determine the probability distribution of finding an electron in a particular region around the nucleus.
The stability of electrons in their respective orbitals is due to a balance between the attractive force of the positively charged nucleus and the repulsive forces between electrons themselves. Electrons occupy the lowest available energy levels first, and they tend to distribute themselves in a way that minimizes their energy.
It's important to note that the concept of electron behavior can be quite complex and involves quantum mechanical principles that can be challenging to visualize. The quantum mechanical model provides a more accurate description of the behavior of electrons in atoms compared to the classical view of orbiting particles.