The statement that "electrons cannot be found between orbitals" is not entirely accurate. In the context of atomic orbitals in an atom, electrons can exist in regions of space between the orbitals, but the probability of finding an electron in those regions is extremely low.
Atomic orbitals describe the probability distribution of finding an electron in a particular region around the nucleus. These regions are often depicted as three-dimensional shapes such as spheres (s orbitals), dumbbells (p orbitals), or more complex shapes (d and f orbitals). The shapes of the orbitals are determined by mathematical solutions to the Schrödinger equation, which describes the behavior of electrons in quantum mechanics.
However, it's important to note that the actual location of an electron in an atom is not fixed or well-defined. According to quantum mechanics, electrons exhibit wave-particle duality, meaning they can behave both as particles and waves. The orbitals represent the most probable locations of finding an electron, but electrons can also exhibit wave-like characteristics and exist in regions of space between the orbitals. These regions are known as electron clouds or electron densities.
The electron density describes the likelihood of finding an electron at a given point in space. The higher the electron density, the greater the probability of finding an electron in that region. While the electron density is generally concentrated around the orbitals, it can extend into the regions between them. However, the probability of finding an electron in these inter-orbital regions is extremely low due to the mathematical nature of the orbitals themselves.
In summary, electrons can exist in regions between orbitals, but the probability of finding them in those regions is very low. The concept of electron density helps us understand the distribution of electrons in space, and while the orbitals provide a good approximation of their locations, they do not represent the exact positions of the electrons.