Structure of Atom: common doubts

Question 1

What did Rutherford's gold foil experiment prove?

Accepted Answer

It showed that an atom's positive charge and almost all of its mass are packed into a tiny nucleus. Most alpha particles went straight through the foil, a few were deflected slightly, and roughly one in twenty thousand came almost straight back. So the atom is largely empty, with a nucleus about 10⁻¹⁵ m across inside an atom about 10⁻¹⁰ m across.

Question 2

Why did Rutherford's model of the atom fail?

Accepted Answer

It could not explain why atoms are stable. Classical physics says a charge moving in a circle is accelerating and must keep radiating energy, so an orbiting electron would lose energy and spiral into the nucleus almost instantly. The model also gave no account of how electrons are arranged around the nucleus or why atoms emit sharp line spectra rather than a continuous band.

Question 3

How do you calculate the number of neutrons in an ion?

Accepted Answer

Neutrons are always A − Z, the mass number minus the atomic number, for ions and neutral atoms alike. The charge changes only the number of electrons. For ²³Na⁺, Z = 11 and A = 23, so there are 12 neutrons and 10 electrons. Subtracting the electron count from A instead of Z is the usual mistake when ions appear in these questions.

Question 4

What is the difference between isotopes and isobars?

Accepted Answer

Isotopes are atoms of the same element, so they share an atomic number but have different mass numbers because their neutron counts differ; protium, deuterium and tritium are an example. Isobars share a mass number but have different atomic numbers, so they are different elements, like ¹⁴C and ¹⁴N. Isotopes behave alike chemically because chemistry depends on electrons, and the electron count follows Z.

Question 5

Why does brighter light not increase the kinetic energy of photoelectrons?

Accepted Answer

Brighter light only delivers more photons each second, so more electrons come out, but each electron still absorbs just one photon. That photon's energy hν first pays the work function and the remainder becomes kinetic energy, KE = hν − W₀. Since hν depends only on frequency, the only way to make each photoelectron faster is to use light of higher frequency.

Question 6

Why are no electrons emitted below the threshold frequency even with very intense light?

Accepted Answer

Below the threshold frequency ν₀ a single photon carries less energy than the work function, and an electron cannot add up energy from several photons. Each photon acts on its own, so no electron ever receives enough to escape, whatever the brightness or exposure time. This all-or-nothing behaviour could not be explained by treating light as a continuous wave, and it supported Planck's quantum idea.

Question 7

What is the work function in the photoelectric effect?

Accepted Answer

The work function W₀ is the least energy needed to pull an electron out of a metal surface, and it equals hν₀, where ν₀ is the threshold frequency. Its value depends on the metal: alkali metals such as caesium have small work functions and release electrons even under visible light. Any photon energy beyond W₀ appears as the electron's kinetic energy, ½mₑv² = hν − W₀.

Question 8

Which series of the hydrogen spectrum lies in the visible region?

Accepted Answer

Only the Balmer series is visible; its lines come from electrons falling to n = 2. The Lyman series ends on n = 1, involves larger energy gaps and lies in the ultraviolet. The Paschen, Brackett and Pfund series end on n = 3, 4 and 5 and fall in the infrared. A bigger energy gap always means higher frequency and shorter wavelength.

Question 9

Why does hydrogen show so many spectral lines when it has only one electron?

Accepted Answer

A sample of hydrogen holds an enormous number of atoms, and different atoms are excited to different levels and fall back by different routes. Every possible jump between two levels produces its own line, so together they give many lines even though each atom has a single electron. Electrons falling from level n down to the ground state can give at most n(n − 1)/2 different lines, so six from n = 4.

Question 10

Why is the energy of an electron in a Bohr orbit negative?

Accepted Answer

The negative sign shows that the electron is bound to the nucleus. Zero energy is assigned to an electron pulled completely out of the atom and at rest, and a bound electron has less energy than this, so its energy is below zero. The more negative the value, the more firmly it is held, which makes n = 1 the most stable level.

Question 11

How do you find the radius and energy of a Bohr orbit for He⁺ or Li²⁺?

Accepted Answer

Use rₙ = 52.9 n²/Z pm and Eₙ = −2.18 × 10⁻¹⁸ Z²/n² J, which hold for any one-electron species. For He⁺ (Z = 2) in n = 1, the radius is half of hydrogen's, 26.45 pm, and the energy is four times as negative, −8.72 × 10⁻¹⁸ J. These formulas do not work for neutral helium, which has two electrons.

Question 12

What are the limitations of Bohr's model?

Accepted Answer

It succeeds only for one-electron species like H, He⁺ and Li²⁺. It cannot account for the spectra of atoms with more electrons, the splitting of spectral lines in magnetic or electric fields, or the way atoms join into molecules. Its fixed circular orbits also contradict the uncertainty principle, since a definite path would need exact position and momentum at the same time.

Question 13

Why don't we observe the wave nature of a cricket ball?

Accepted Answer

Its wavelength is far too tiny to detect. By de Broglie's relation λ = h/mv, the wavelength shrinks as mass and speed grow, and Planck's constant is only 6.626 × 10⁻³⁴ J s. A ball of about a tenth of a kilogram gets a wavelength vastly smaller than an atom, whereas an electron's minute mass gives a wavelength close to atomic sizes, which shows up in diffraction.

Question 14

What does Heisenberg's uncertainty principle mean for electrons?

Accepted Answer

It means an electron's position and momentum cannot both be fixed exactly at the same instant: Δx·Δp ≥ h/4π. Locating the electron more sharply makes its momentum less certain, and the other way round. For something as light as an electron the uncertainty is large on the scale of an atom, so a precise orbit makes no sense and we speak only of probabilities.

Question 15

What is the difference between an orbit and an orbital?

Accepted Answer

An orbit is Bohr's definite circular track at a fixed distance from the nucleus, while an orbital is a region described by a wave function ψ, with ψ² giving the probability of finding the electron there. Orbits assume exact position and momentum, which the uncertainty principle forbids. Orbitals come in different shapes, such as spherical s and dumbbell-shaped p, and each holds at most two electrons.

Question 16

What do the four quantum numbers tell us?

Accepted Answer

The principal number n gives the shell and largely fixes size and energy. The azimuthal number l, from 0 to n − 1, gives the subshell and its shape. The magnetic number mₗ, from −l to +l, gives the orbital's orientation in space. The spin number mₛ, +½ or −½, gives the electron's spin. Three numbers label an orbital; all four identify one electron.

Question 17

How do you calculate radial and angular nodes of an orbital?

Accepted Answer

Radial nodes equal n − l − 1 and angular nodes equal l, so the total is n − 1. A 3p orbital (n = 3, l = 1) has one radial and one angular node, and a 4d orbital (n = 4, l = 2) has one radial and two angular nodes. The usual slip is to take n − 1 as the radial count, which gives 3p two radial nodes.

Question 18

Why do 2s and 2p have the same energy in hydrogen but not in other atoms?

Accepted Answer

In hydrogen there is only one electron, so its energy depends on n alone. In atoms with several electrons, inner electrons screen the nucleus, and an s electron penetrates closer to the nucleus than a p electron of the same shell. It feels a larger effective nuclear charge and is held more strongly, so 2s lies below 2p, and in the third shell 3s < 3p < 3d.

Question 19

Why does 4s fill before 3d?

Accepted Answer

Because 4s has the lower (n + l) value. For 4s it is 4 + 0 = 4, and for 3d it is 3 + 2 = 5; the subshell with the smaller sum is lower in energy and fills first. If two subshells tie, the one with the smaller n fills first, which is why 4f comes before 5d even though both have n + l = 7.

Question 20

Why are chromium and copper exceptions in electronic configuration?

Accepted Answer

Chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹ because half-filled and completely filled subshells are especially stable. Such sets are symmetrical, and having more electrons with parallel spins raises the exchange energy that lowers the total energy. Shifting one 4s electron into 3d buys this stability, so 3d⁴ 4s² and 3d⁹ 4s² are not the ground states.

Question 21

What is the difference between Pauli's exclusion principle and Hund's rule?

Accepted Answer

Pauli's principle says no two electrons in one atom can share all four quantum numbers, so an orbital holds at most two electrons, and they must have opposite spins. Hund's rule says orbitals of equal energy are first filled singly, with parallel spins, before any pairing begins. That is why nitrogen's three 2p electrons occupy separate orbitals and pairing starts only with oxygen's fourth p electron.

Question 22

How is Structure of Atom usually tested in NEET?

Accepted Answer

Expect numericals on photon energy, the photoelectric equation, de Broglie wavelength, and Bohr radius and energy for one-electron species, along with questions on which transition gives a line in a given series. Conceptual items cover quantum numbers, counting nodes, the (n + l) rule and the Cr and Cu exceptions. Common slips are unit errors and confusing radial nodes with total nodes.

Structure of Atom: common doubts, answered

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