Electrons in an atom do experience an attractive force due to the positive charge of the nucleus. This force arises from the electrostatic interaction between the negatively charged electrons and the positively charged protons.
However, according to quantum mechanics, electrons in atoms do not behave like classical particles moving in well-defined orbits around the nucleus. Instead, they exist in certain energy levels or orbitals that are quantized and characterized by specific energy values. These orbitals define regions where the electrons are most likely to be found.
The behavior of electrons in atoms is described by wave functions, which give the probability distribution of finding an electron in a particular location. The wave functions are determined by solving the Schrödinger equation, a fundamental equation in quantum mechanics.
In these orbitals, electrons occupy specific energy levels, and there are restrictions on the number of electrons that can occupy each level. The Pauli exclusion principle states that no two electrons in an atom can have the same set of quantum numbers, which limits the number of electrons that can occupy a given orbital. This principle helps to prevent the electrons from all collapsing into the nucleus.
Additionally, the Heisenberg uncertainty principle plays a role. It states that there is a fundamental limit to how precisely both the position and momentum of a particle can be known simultaneously. This means that the exact location of an electron cannot be determined with certainty, but rather, it is described by a probability distribution.
In summary, electrons in atoms are governed by quantum mechanics, and their behavior is determined by the interplay of quantum principles, energy levels, and wave functions. These factors prevent the electrons from crashing into the nucleus and instead distribute them in specific orbitals or energy levels around the nucleus.