26.2 Problems with Rutherford's Model

According to Rutherford, the scattering alpha particles indicated that within the atom there existed a tiny, but very dense, positively charged core.  Rutherford concluded that the atom was not filled with a positively charged substance (as Thomson had described); rather, all the positive charge of the atom was located in a nucleus at the centre of the atom.  This nucleus was small but contained almost all the mass of the atom.  Thus, Rutherford proposed a nuclear model.  In this model the atom has a dense nucleus with relatively vast amounts of empty space through which the electrons can pass.  The negative charge of the orbiting electrons was the opposite of the positive charge of the nucleus; so, overall, the atom is still electrically neutral, as Dalton determined.

There was a problem with the nuclear model of the atom.  Recall that positive and negative charges attract.  If the nucleus were positively charged, then what stopped the electrons from being sucked into the nucleus?  To address this problem Rutherford suggested that the electrons orbit the nucleus, much like the moon orbits Earth or like Earth orbits the sun.  The force of attraction between the electrons and the nucleus provides the force necessary to keep the electrons in orbit.  Hence, Rutherford proposed a planetary model of the atom.

Failure of the Planetary Model

The planetary model was also severely flawed.  According to Rutherford's model, atoms are not stable and will collapse in on themselves.  According to Maxwell's electromagnetic theory, when charged particles like electrons are accelerated, they emit electromagnetic radiation.  Electrons orbiting a nucleus undergo inward acceleration; thus, they should continuously emit electromagnetic radiation.  And, if the electrons emit electromagnetic radiation, they should be losing energy.  If the electrons lose energy, they will eventually spiral into the nucleus.  Why wasn't this happening?  By the end of the 19th century an adequate model of the atom had not yet surfaced.

The Role of Atomic Spectra

In developing a model of the atom, scientists also had to account for atomic spectra.  By the early 1800s, scientists knew that every element emits unique line spectra.  For example, when an evacuated bulb is filled with neon gas and a voltage is applied to the electrodes, a characteristic red glow is emitted.

Separating this light by wavelength reveals that neon gas is actually emitting a small collection of unique wavelengths that fall in the red to yellow region of the visible light spectrum.  You can see these unique wavelengths identified on the line spectrum below the bulbs.  If the bulb were filled with a different gas, argon for example, a blue colour would be produced with a different and unique line spectrum.  A line spectrum is like a chemical fingerprint as each element has its own distinct pattern.

Why does each element have unique spectra, and why is it a line spectra?  Somehow, these phenomena must be tied into the mystery of atomic structure.

Understanding line spectra is important.  Continue to page 4 of this lesson to learn about the three types of spectra!