Q.Which of the following conclusions could not be derived from Rutherford's α-particle scattering experiement?

Concept: Experimental scope vs. theoretical model

Rutherford's α-particle scattering experiment (1911) involved firing alpha particles at thin gold foil and observing their deflection patterns. The key observations were:

1. Most particles passed straight through → the atom is mostly empty space (A is a direct conclusion).

2. A tiny fraction scattered at large angles → a small, dense, positively charged nucleus exists. Comparing the scattering probability with atomic dimensions gave the nuclear radius ~10⁻¹⁵ m versus the known atomic radius ~10⁻¹⁰ m (B follows directly).

3. Large-angle deflections → strong repulsive forces act at close range, consistent with electrostatic repulsion between positive charges. Since the atom is neutral overall, negatively charged electrons must be attracted to the nucleus electrostatically (D is a valid inference).

However, the experiment revealed nothing about electron motion or energy quantization. The idea that electrons occupy fixed circular orbits with discrete energies came from Bohr's 1913 model, which was introduced later to explain atomic spectra—not from scattering data.

The Scale Analogy: What the Experiment Actually Told Us

Imagine throwing tennis balls at a large curtain. If most pass straight through, you know the curtain is mostly empty space. If a few bounce back sharply, something small and very hard must be hidden inside. That's exactly what Rutherford's α-particle scattering revealed about atomic structure.

When Rutherford fired positively charged α-particles at a thin gold foil, he observed three key patterns:

• Most particles passed through undeflected
• Some deflected at small angles
• A tiny fraction (about 1 in 8000) bounced back at large angles, even straight back

These observations were shocking because the prevailing "plum pudding" model predicted only minor deflections. Rutherford famously said it was like firing a naval shell at tissue paper and having it bounce back.

What the Experiment Could and Could Not Tell Us

Let's examine each conclusion:

1. Most of the space in the atom is empty — This follows directly from the observation that the vast majority of α-particles passed through the foil without deflection. If the atom were a uniform distribution of matter, every particle would interact significantly. The unimpeded passage proves the atom is mostly void.

2. The radius of the atom is about $10^{-10}$ m while that of nucleus is $10^{-15}$ m — The scattering angles and the fraction of particles deflected allowed Rutherford to estimate the size of the deflecting center (the nucleus). By comparing this to the known atomic spacing in the foil (from X-ray crystallography and other methods), he could establish the scale difference: the nucleus is about 100,000 times smaller than the atom itself.

3. Electrons move in a circular path of fixed energy called orbits — Here's the critical point. The α-particle scattering experiment told Rutherford where the positive charge and mass were concentrated (in a tiny nucleus), but it revealed nothing about electron motion. The scattered α-particles interacted with the nucleus through electrostatic repulsion; they didn't map out electron trajectories. The concept of fixed circular orbits came later from Bohr's model (1913), which was built to explain atomic spectra, not scattering data.

4. Electrons and the nucleus are held together by electrostatic forces of attraction — Once Rutherford established that the nucleus is positive and tiny, and knowing that atoms are electrically neutral overall, the negative electrons must be distributed in the surrounding space. The only force that could keep them bound to the positive nucleus over atomic distances is the electrostatic (Coulomb) attraction. This is a logical inference from the nuclear model.