Name:Date:

Period:Introduction to the Bohr Model of Atomic Structure.

In the early 1900’s, thanks in part to Rutherford’s famous gold foil experiment, scientistsknew several facts about the structure of atoms.

  • Most of the mass of an atom occupies a tiny but extremely dense and positively charged region in the center of the atom called the nucleus. This discovery led to the recognition of protons as nuclear material; neutrons were discovered some years later…
  • The much less massive electrons are found outside the nucleus, in a relatively much larger region of what is mostly empty space.

This image is not to scale. If the nucleus were really as big as you see it here, then the distance between the nucleus and the electrons would be 100 meters or more further away! Fact: The total volume of an atom is mostly empty space.

  • Rutherford’s nuclear atomic model was consistent with all of the available evidence known about atoms and chemical reactions at that point in history.
  • A major flaw in this model, however, was that it did not explain the periodic properties of the elements first identified by Russian scientist DmitriMendeleev(remember he’s the guy who created the first periodic table). It also seemed to defy known physical laws, which predicted that electrons should be drawn into the nucleus. Clearly a better model of atomic structure was needed.

In 1913, a scientist named Neil’s Bohr proposed an improved model of atomic structure, which was successful in describing the line spectrumof hydrogen, shown below.

The four colors represent different wavelengths of electromagnetic radiation, which are emitted by hydrogen gas when it is placed in a strong electric field. Bohr interpreted this observation by concluding that hydrogen’s single electron could exist only at a limited number of distances from the nucleus. Bohr assumed that the electron revolved around the nucleus in a circular orbit. When the electron absorbed a certain quantity of energy, it would jump to a higher allowed orbit. These different orbits are also called energy levels. When an electron is found in the smallest allowed orbit (closest to the nucleus), we say the electron is in the ground state. When an electron is found in a higher energy orbit, we say the electron is in an excited state. The colors we see in hydrogen’s line spectrum occur when an electron relaxes from a higher energy orbit to a lower energy orbit, emitting energy as photons of light.

Today chemists and physicists recognize that Bohr’s assumption of circular orbits is incorrect. The motion of electrons in atoms is much more complex than simple circular orbits. Indeed electrons, like light, have a dual nature exhibiting properties of both particles and of standing waves. For this reason, energy level is a better word choice than orbit. However we often still adopt this circular orbit assumption because its simplicity makes the Bohr model of the atom easy to work with.

The following diagramsshow the wavelengths of light (EMR) produced when the excited electron in a hydrogen atom relaxes to a lower energy level. Notice the some relaxations produce UV light, and infrared radiation, in addition to the four visible wavelengths.

Bohr’s conclusions about hydrogen’s line spectrum were applied to the atoms of other elements, but with less success. Soon an even better model of atomic structure was developed to replace the Bohr model. The most modern model of atomic structure is called the Quantum Mechanical Model, which builds off of Bohr’s ideas but is far more complex and mathematical (and beyond the scope of most first year high school chemistry classes).

Still, his concepts do help to understand some properties of the elements, especially those with atomic numbers 1-18. One additional conclusion that is still accepted today has to do with the number of electrons a given energy level can hold. As the energy levels increase in size (and energy, as well as distance from the nucleus), the number of electrons an energy level can hold increases according to the formula 2(N) 2, where N is the number of the energy level.

For example, the 1st lowest energy level (N=1) can hold 2(1) 2 = 2 electrons.

The second energy level (N=2) can hold 2(22) = 8 electrons.

How many electrons can the 3rd energy level hold?

When sketching Bohr diagrams of atoms with all of their electrons in the ground state, electrons always fill lower energy levels before occupying larger, more energetic energy levels.

Here is an example, the Bohr atomic diagram for Sodium: