The Bohr model offers a simplified, two-dimensional depiction of an atom’s structure, useful for visualizing fundamental concepts. Developed by physicist Niels Bohr in 1913, this model represents electrons orbiting a dense, central nucleus in fixed, circular paths called energy shells. Each shell corresponds to a discrete energy level, similar to rungs on a ladder, meaning electrons cannot exist in the space between these levels. While the model is foundational for understanding the arrangement of subatomic particles, it is a historical representation and not an accurate modern view of atomic reality.
Calculating Protons, Neutrons, and Electrons
The process of drawing a Bohr model begins with determining the number of subatomic particles contained within the atom. This information is readily available on the Periodic Table, where the atomic number and atomic mass are listed for every element. The atomic number, the smaller of the two figures, directly indicates the number of protons found in the nucleus.
For a neutral atom, the count of negatively charged electrons orbiting the nucleus must precisely match the number of positively charged protons. Therefore, the atomic number provides both the proton and electron counts necessary for the model. The third required particle is the neutron, which is also located in the nucleus and contributes to the atom’s mass but carries no electrical charge.
To find the number of neutrons, the atomic mass must be used, which represents the combined total of protons and neutrons. Since atomic masses are typically given with decimals, round this figure to the nearest whole number. The number of neutrons is then calculated by subtracting the atomic number (the proton count) from this rounded atomic mass.
Step-by-Step Guide to Drawing the Model
The first physical step in constructing the model is to draw the nucleus, representing the atom’s center. Inside this circle, the calculated number of protons and neutrons should be written and labeled, typically using “P+” and “N” for simplicity. Next, begin drawing the energy shells, which are concentric circles around the nucleus, starting with the one closest to the center. These shells will house the total number of electrons.
The placement of electrons into these shells follows specific maximum capacity rules, ensuring the lowest energy levels are filled first. The innermost shell, designated as the first energy level, can hold a maximum of two electrons. Once this first shell is complete, the process moves outward to the second energy shell, which can accommodate up to eight electrons.
For most elements used in introductory models, the third energy shell also holds a maximum of eight electrons before the next shell begins to fill. Electrons are typically placed around the shell’s perimeter until the total required count is reached. The electrons placed in the outermost, partially filled shell are known as valence electrons, and their number dictates the atom’s basic chemical behavior.
When the Bohr Model Is Inaccurate
The Bohr model provides an excellent instructional tool but quickly becomes an inaccurate representation when applied beyond the simplest atoms. It was developed primarily to explain the spectral lines of hydrogen, an atom containing only a single electron. The model fails to accurately describe the behavior of atoms with two or more electrons, known as multi-electron atoms.
The primary reason for this failure is that the model ignores the complex interactions between multiple electrons. Electron-electron repulsion significantly affects energy levels and distribution, a factor that the simplified Bohr orbits do not account for. Furthermore, the model does not consider the existence of subshells, which are the more complex s, p, d, and f orbitals that house electrons in reality.
The fixed, circular orbits depicted in the model also fundamentally contradict the principles of modern quantum mechanics. According to the quantum mechanical view, an electron’s exact position and momentum cannot be known simultaneously, a concept known as the Heisenberg Uncertainty Principle. Instead of moving in predictable paths, electrons exist in three-dimensional probability regions, often visualized as electron clouds, which are far less defined than the rings of the Bohr model. This historical model is a gross oversimplification that cannot explain phenomena like the fine structure observed in spectral lines or the complexities of chemical bonding.