Bohr Model of the Atom

Development leading to the Bohr Model of the Atom

Our understanding of the atom went through several distinct models throughout the last two centuries, improving with accuracy as new evidence was obtained and more fundamental physical principles were revealed.

Plum Pudding Model

Before the 20th century, we had no idea what the subatomic structure of the atom was like. In 1803 John Dalton theorized that the atom was indivisible and could not be broken down into anything smaller. However, when the electron was discovered in 1897 by J. J. Thomson, this all changed. After much deliberation and research, he proposed the plum pudding model or the Thomson Model.

The model attempted to explain how an atom had no net electric charge, yet individual electrons possessed negative charges. Thomson proposed that negatively charged "plums" (electrons) were surrounded by a positively charged "pudding," as an atom must contain some positive charge to cancel out the negative charge of the electrons.

Rutherford Model

In 1905, Thomson’s student Ernst Rutherford tested the plum pudding model by directing a beam of alpha particles at a strip of gold foil. Alpha particles are a form of radiation with a large positive charge. He expected the alpha particles to pass through the gold with no deflection as the positively charged "pudding" should be evenly spread out. However, a very small number of the alpha particles were deflected, sometimes being reflected completely.

In the experiment with alpha particles, the expectation was for the alpha particles to go through without any deviation, in accordance with Thomson's model. However, a small number of the particles were deflected, which can be explained by Rutherford's Model, Wikimedia Commons CC-BY-SA-3.0

It took several years for Rutherford to correctly interpret the results of his experiment. He proposed that the atom actually consisted of a small, compact, and positively charged nucleus surrounded by a cloud of electrons, called the Rutherford model. The proton and the neutron (the subatomic particles that make up the nucleus) were discovered later, in 1917 and 1932, respectively.

Bohr Model

Unfortunately, the Rutherford model was still flawed. The leading theory at the time was that the electrons revolved around the nucleus in arbitrary circular orbits, like how a planet orbits a star. However, electrons lose energy when they are accelerated, and therefore they should collapse into the nucleus under the Rutherford model. In 1913, Niels Bohr proposed his own structure of the atom to explain this.

Subatomic structure of a Hydrogen atom. An electron absorbs and emits energy to move between energy levels, Wikimedia Commons CC-BY-SA-3.0

In Bohr's model of the atom, electrons orbit around the nucleus in fixed energy levels called shells. Electrons can only exist in these shells and they move between them by gaining and losing certain amounts of energy corresponding to the energy difference between the energy levels. This is how Niels Bohr was able to explain why the atom does not collapse. His conclusions are based on quantum physics, which you won't have to worry too much about at this level.

Atomic Structure of the Bohr Model

Energy and Color

Almost all light in the universe comes from atoms. When scientists study a pure element, they observe that only a few specific colors are ever emitted. Different elements emit different patterns of colors, called an emission spectrum. For example, hydrogen has an orange line, two blue lines, and a purple line. The colors and patterns of hydrogen are distinct, even from helium.

Emission Spectrum comparison of Hydrogen and Helium, Wikimedia Commons CC-BY-SA-4.0

Remember that light is pure energy in the form of an electromagnetic wave or photon. The energy of an EM wave determines its color. Redder light has lower energies, while bluer light has higher energies. As an atom can only absorb and emit certain colors of light, we know that something within the structure of the atom must have specific energy levels.

An electron absorbs and emits photons to move energy levels, Wikimedia Commons CC BY 3.0

In Bohr’s model, electrons orbit the nucleus at discrete energy levels called shells. They cannot exist between energy levels unless they are passing from one level to another. The lowest level (E1) is the ground state, while all higher energy levels are excited states. It’s important to note that the energy of an electron in an atom is always negative. The negative symbol denotes that the electron must be given energy to be ejected from the atom entirely. This energy is defined as ionization energy or ionization state. Consider a hydrogen atom with a ground state of $-13.6\mathrm{eV}$, giving the electron ionization energy of $13.6\mathrm{eV}$ will force it to be ejected.

When an electron moves from a higher energy level to a lower one a photon is emitted with an energy corresponding to the energy difference between the two shells. If an electron moves between two energy levels, the energy of the emitted photon can be calculated with the Planck equation:

$E=hf=E2-E1$

Where$E$is Energy,$h$is Planck’s constant, and$f$is the frequency of light produced.

Similarly, this equation can also be used when an electron absorbs energy from a photon and moves into a more excited state. Electrons in an atom can only absorb certain energies of light, which would move it to a more excited state. Electrons are able to move between multiple orbits in one transition, for example, between$E4$and$E2$.

How Electrons fill Energy Levels

There is a limit to how many electrons can physically exist in any shell at a given point in time. The first shell can only hold two electrons, while shells 2 and 3 can contain 8 electrons each. The diagram below shows the electron configuration of ten elements on the periodic table, assuming no electron has been excited to a higher state.

The electron configuration of ten elements, Wikimedia Commons CC-BY-4.0

As electrons are added to an atom, they will fill the lowest unfulfilled energy levels available. Each shell must be filled before the next starts to fill. The only exception is when an electron has been excited. However, excited electrons will very quickly undergo spontaneous emission of a photon to return to a lower energy level. This is because the negatively-charged electrons are attracted to the positively-charged nucleus, like how gravity attracts a skydiver to the ground.

Limitations of the Bohr Atomic Model

Bohr's model was essentially flawless, as long as the atom you were studying had only one electron. Unfortunately, every single atom on the periodic table, with the exception of hydrogen, has more than one electron. When Bohr attempted to use his atomic model to predict the spectral lines of other elements, his calculations drifted further and further from the actually observed emission spectrums as the number of electrons in a sample element increased.

One problem is that multiple electrons in the 'orbit' of a nucleus will begin to interact with each other, complicating the energy structures of the electron shells. Erwin Schrödinger, in 1926, determined that electrons actually move around the nucleus in different clouds according to their energy level. The white regions on the image below show a higher probability of finding an electron in that space, while darker regions demonstrate the opposite.

In Schrödinger's Quantum Model of the atom, electrons move around the nucleus in different clouds, Pixabay

Advantages and Disadvantages of Bohr's Atomic Model

The main reason Bohr's model is useful is that it clearly demonstrates the underlying structure of the atom and how electrons can gain and lose energy through the absorption and emission of photons.

Unfortunately, Bohr's model isn't completely correct due to quantum mechanical principles. Electrons do not "orbit" the nucleus at neat, distinct lines but instead in indistinct clouds. However, Bohr's model is relatively neat and accurate and is an important fundamental step at high school level to understanding the physics that governs the world.