Structure of atom
Frequently Asked Questions
Bohr’s model assumes electrons move in fixed circular orbits and considers only Coulombic attraction between nucleus and one electron. In multi-electron atoms, electron–electron repulsion, shielding, and relativistic effects alter energy levels, which the Bohr model cannot explain.
Because electrons in hydrogen can occupy only quantized energy levels. When an electron transitions between these levels, it emits or absorbs photons of specific energies, producing discrete spectral lines.
According to the (n + l) rule, orbitals with lower (n + l) value fill first. For 4s, n+l = 4; for 3d, n+l = 5. Hence 4s has lower energy initially and is filled before 3d.
Half-filled and fully filled subshells have symmetrical electron distribution and maximum exchange energy, which lowers overall energy and increases stability.
Across a period, proton number increases while shielding remains nearly constant because electrons are added to the same shell. Thus, net attractive force on valence electrons increases.
Increasing effective nuclear charge pulls electrons closer to the nucleus without significant increase in shielding, reducing atomic radius.
Down a group, additional electron shells are added, increasing distance between nucleus and valence electrons. Shielding effect also increases, enlarging atomic size.
It establishes that position and momentum of an electron cannot be simultaneously determined with precision, proving that electrons cannot have fixed classical orbits.
For l = 2 (d-subshell), magnetic quantum number m? can take values -2, -1, 0, +1, +2. Thus there are five possible orientations, giving five d-orbitals.
Increasing effective nuclear charge holds valence electrons more strongly, requiring more energy to remove an electron.
Zero energy is defined when electron is infinitely far from nucleus. In bound state, electron possesses less energy than free state, hence energy is negative.
Balmer series corresponds to transitions ending at n = 2. The energy differences for these transitions fall within visible wavelength range.
Inner shell electrons repel outer electrons, partially screening the nuclear attraction, thereby reducing the net effective nuclear charge experienced by valence electrons.
Quantum mechanics defines electron position in terms of probability density derived from wave function \(\psi\). The square of wave function \((\psi)^2\) gives probability distribution.
Each orbital can have only one set of four quantum numbers. Since spin quantum number has only two values (+½ and -½), maximum two electrons can occupy one orbital with opposite spins.