Atomic Models

Atomic Model

Key Features

Drawbacks

Thomson's Model (J. Thomson (1856-1940))

J. Thomson , a prominent British physicist, made a significant contribution to atomic understanding, earning him the Nobel Prize in Physics in 1906 for his pivotal work on the discovery of electrons.

Thomson envisioned the atom to resemble a watermelon, where the positive charge is dispersed throughout the atom like the red, edible part of the watermelon, while the electrons are interspersed within this positively charged sphere, akin to seeds in a watermelon.

Thomson's atomic model proposed the following:

1. The atom is composed of a positively charged sphere, with electrons embedded within it.

2. The negative and positive charges within the atom are of equal magnitude, resulting in the atom being electrically neutral.

Limitations of Thomson's Atomic Model:

Thomson's atomic model successfully accounted for the overall electrical neutrality of an atom. However, it encountered challenges when attempting to elucidate the specific arrangement and distribution of electrons within the atom.

One significant shortcoming was its inability to explain the distinct chemical properties observed among various elements. If we assume that electrons are uniformly embedded within the positively charged sphere, the opposite electric charges would cancel each other, rendering the charged sphere electrically neutral. This model failed to provide insights into why different elements exhibit unique chemical behaviors despite their similar atomic structures.

Rutherford's Model (Ernest Rutherford (1871-1937))

The α-Particle Scattering Experiment:

In this experiment, a stream of α particles from a radioactive source was directed on a thin (about 0.00004 cm thick) piece of gold foil.

However, the actual results of the experiment were quite surprising. It was observed that:

• The majority of fast-moving α-particles traversed the gold foil without significant deflection.
• Some α-particles experienced minor deflections upon interaction with the foil.
• Surprisingly, approximately one out of every 12000 particles appeared to rebound.

Rutherford deduced from the α-particle scattering experiment that:

• Most of the space within an atom is empty, as the majority of α-particles passed through the gold foil with minimal deflection.
• A tiny fraction of α-particles being deflected implied that the positive charge of the atom is concentrated in a minute volume.
• Calculations based on the data revealed that the nucleus, a positively charged center within the atom, holds almost all of the atom's mass and occupies a remarkably small volume.

These findings led Rutherford to propose the nuclear model of an atom, highlighting:

• A positively charged nucleus at the atom's center, containing nearly all of its mass.
• Electrons revolving around the nucleus in distinct orbits.
• The nucleus being substantially smaller in size compared to the overall atomic size, emphasizing its high density.

Limitations of Rutherford's Model of the Atom:

The Rutherford model, while revolutionary, faced significant challenges. One notable issue pertains to the stability of electron orbits. According to classical electromagnetic theory, an electron undergoing circular motion should constantly radiate energy due to acceleration. As a result, the electron would lose energy over time and eventually spiral into the nucleus. Such a scenario would render atoms highly unstable. Hence, he failed to explain stability of an atom.

It says nothing about the electronic structure of atoms i.e., how the electrons are distributed around the nucleus and what are the energies of these electrons.

Bohr's Model

(Niels Bohr (1885-1962))

Bohr's Model of an Atom:

1. Quantized Energy Levels: Electrons orbit the nucleus in specific, quantized energy levels, often referred to as shells or orbits. These orbits are stationary and have fixed energy values.

2. Stable Orbits (Stationary Orbits): Electrons in these orbits do not radiate energy and are stable. These orbits are sometimes called stationary shells due to the absence of energy radiation.

3. Energy Increases with Distance from Nucleus: The energy of an orbit increases as its distance from the nucleus increases. Orbits closer to the nucleus have lower energy levels.

4. Specific Energy Amounts: Electrons possess discrete amounts of energy that allow them to occupy particular orbits. The size of the orbit is inversely proportional to the energy it holds—smaller orbits have higher energies.

5. Energy Absorption and Emission: When an electron transitions from a lower energy level to a higher one, it absorbs energy. Conversely, when it moves from a higher energy level to a lower one, it emits energy.

6. Absorption and Emission in Electron Transition: The transfer of an electron from one orbit to another involves either the absorption or emission of energy, depending on the direction of the transition.

Limitations of Bohr theory –

• It could not explain the ability of atoms to form molecules by chemical bonds.
• It fails to account for the finer details (doublet, that is two closely spaced lines) of the hydrogen atom spectrum observed by using sophisticated spectroscopic techniques.
• This model is also unable to explain the spectrum of atoms other than hydrogen, for example, helium atom which possesses only two electrons.
• Further, Bohr’s theory was also unable to explain the splitting of spectral lines in the presence of magnetic field (Zeeman effect) or an electric field (Stark effect).