Question 1

Why did Bohr’s model fail for multi-electron atoms?

Accepted Answer

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.

Question 2

Why does the hydrogen atom show a line spectrum instead of a continuous spectrum?

Accepted Answer

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.

Question 3

Why is 4s filled before 3d in electronic configuration?

Accepted Answer

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.

Question 4

Why are half-filled and fully filled subshells more stable?

Accepted Answer

Half-filled and fully filled subshells have symmetrical electron distribution and maximum exchange energy, which lowers overall energy and increases stability.

Question 5

Why does effective nuclear charge increase across a period?

Accepted Answer

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.

Question 6

Why does atomic size decrease across a period?

Accepted Answer

Increasing effective nuclear charge pulls electrons closer to the nucleus without significant increase in shielding, reducing atomic radius.

Question 7

Why does atomic radius increase down a group?

Accepted Answer

Down a group, additional electron shells are added, increasing distance between nucleus and valence electrons. Shielding effect also increases, enlarging atomic size.

Question 8

Why is Heisenberg’s uncertainty principle important in atomic structure?

Accepted Answer

It establishes that position and momentum of an electron cannot be simultaneously determined with precision, proving that electrons cannot have fixed classical orbits.

Question 9

Why are d-orbitals five in number?

Accepted Answer

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.

Question 10

Why does ionization energy generally increase across a period?

Accepted Answer

Increasing effective nuclear charge holds valence electrons more strongly, requiring more energy to remove an electron.

Question 11

Why is the energy of electron in hydrogen negative?

Accepted Answer

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.

Question 12

Why does the Balmer series appear in the visible region?

Accepted Answer

Balmer series corresponds to transitions ending at n = 2. The energy differences for these transitions fall within visible wavelength range.

Question 13

Why does shielding effect reduce effective nuclear charge?

Accepted Answer

Inner shell electrons repel outer electrons, partially screening the nuclear attraction, thereby reducing the net effective nuclear charge experienced by valence electrons.

Question 14

Why are orbitals described as probability clouds?

Accepted Answer

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.

Question 15

Why does Pauli exclusion principle limit two electrons per orbital?

Accepted Answer

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.

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