Changes in the Atomic Model

Over time, scientists have developed several models of the atom to explain experimental evidence and improve our understanding of atomic structure. This note covers the main changes from Dalton's solid sphere model to Bohr's model, including the discovery of protons and neutrons.

Early Atomic Models

The understanding of the atom has changed significantly over time as new evidence emerged.

Dalton's Solid Sphere Model

In the early 19th century, John Dalton proposed that atoms were tiny, solid, indivisible spheres. He suggested that each element was made up of identical atoms, different from those of other elements. This model explained the idea of elements combining in fixed ratios but did not include any internal structure.

Thomson's Plum Pudding Model

In 1897, J.J. Thomson discovered the electron, a tiny negatively charged particle much smaller than an atom. This led to the plum pudding model, where the atom was thought to be a positively charged 'pudding' with negatively charged electrons scattered like 'plums' inside it.

This model explained the presence of electrons but could not explain later experimental results showing the atom's structure was more complex.

Rutherford's Nuclear Model

In 1909, Ernest Rutherford and his team performed the gold foil experiment, which changed the atomic model dramatically.

Gold Foil Experiment

Alpha particles (positively charged helium nuclei) were fired at a very thin sheet of gold foil. Most passed straight through, but some were deflected at large angles, and a few bounced straight back.

Nucleus Discovery

Rutherford concluded that the atom must have a tiny, dense, positively charged centre called the nucleus, where almost all the mass is concentrated. The rest of the atom is mostly empty space, with electrons orbiting this nucleus.

This explained why most alpha particles passed through (empty space) and why some were deflected (repelled by the positive nucleus).

For instance, if alpha particles are fired at gold foil and 1 in 8000 bounces back, this shows the nucleus is very small compared to the atom’s size.

Bohr's Model of the Atom

Niels Bohr improved on Rutherford’s model by suggesting that electrons orbit the nucleus in fixed paths or energy levels, rather than randomly.

Electrons in Fixed Orbits

Electrons can only exist in certain allowed orbits with fixed energies. They do not spiral into the nucleus because they can only occupy these specific energy levels.

Energy Levels and Emission Spectra

When electrons jump between energy levels, they absorb or emit specific amounts of energy as electromagnetic radiation. This explains the line spectra seen in elements, where only certain colours (wavelengths) are emitted.

For example, when an electron in a hydrogen atom falls from a higher orbit to a lower one, it emits light of a specific wavelength, producing a line in the emission spectrum.

Discovery of Protons and Neutrons

Proton Identification

After the nucleus was discovered, scientists found it contained positively charged particles called protons. Each proton has a charge of +1 and a mass of about 1 atomic mass unit (amu).

Neutron Discovery

The total mass of the nucleus was found to be greater than the total mass of the protons alone. This led to the discovery of neutrons in 1932 by James Chadwick. Neutrons have no charge but have a similar mass to protons.

Nucleus Composition

The nucleus is made up of protons and neutrons, which together account for nearly all the atom’s mass. Electrons orbit the nucleus and have negligible mass compared to protons and neutrons.

For example, the nucleus of a carbon atom contains 6 protons and 6 neutrons, giving it a mass number of 12.

Example: Calculate the mass number of an atom with 8 protons and 8 neutrons.

Mass number = number of protons + number of neutrons = 8 + 8 = 16.

Example: Understanding the Gold Foil Experiment

Alpha particles were fired at a thin gold foil. Most passed straight through, but 1 in 8000 were deflected back. What does this tell us about the atom’s structure?

Since most alpha particles passed through, the atom must be mostly empty space. The few deflected particles indicate a small, dense, positively charged nucleus that repels the alpha particles.