Bohr Atomic Model and Electron Configuration


This is part of preliminary HSC Chemistry course under the topic of Atomic Structure and Atomic Mass.

HSC Chemistry Syllabus

  • Model the atom's discrete energy levels, including electronic configuration and SPDF notation (ACSCH017, ACSCH018, ACSCH020, ACSCH021)
  • Investigate energy levels in atoms and ions through:
– Collecting primary data from a flame test using different ionic solutions of metals (ACSCH019)
– Examining spectral evidence for the Bohr model and introducing the Schrödinger model 

      Bohr Model - Flame Test, AES, Electron Configuration

      Bohr's Model of the Atom 

      Bohr's atomic model


      Bohr's model of the atom builds on the contribution by Rutherford. In this model, Bohr postulated:

      1. Electrons revolve around the nucleus in circular orbits with discrete radii and quantised energies. In these orbits, electrons exist in 'stationary states' and do not emit energy. 


      Bohr’s model of the atom describes electrons orbiting in stable energy levels as opposed to Rutherford's model in which electrons' motion was not described.


      Bohr model electron transition spectroscopy 


      2. An electron can transition between orbits by absorbing or releasing energy that is exactly equal to the difference in energy of orbits, consistent with the law of conservation of energy. 

      Electron excitation occurs when an electron absorbs energy to move to an orbit of higher energy.

      Electron relaxation occurs when an electron moves to a lower orbit, releasing energy in the form of electromagnetic radiation (photon). 


      If the photon's energy is less or greater than the difference in energy of orbits involved in the electron transition, no transition will occur. This interaction between electrons and energy is referred to as spectroscopy.

      When electrons absorb energy to move to orbits of higher energies, absorption spectroscopy occurs. Conversely, when electrons release energy while moving to orbits of lower energies, emission spectroscopy occurs.

      Electron Configuration in Bohr's Atomic Model

      Electron configuration refers to the arrangement of electrons in an atom's energy levels (orbits or shells). In Bohr's model, these configurations are determined by the principles of quantum mechanics (knowledge is not required in HSC Chemistry), with each shell capable of holding a certain maximum number of electrons:


      • The first shell (closest to the nucleus) can hold up to 2 electrons.
      • The second shell can hold up to 8 electrons.
      • The third and subsequent shells can hold more, but for most elements in the first few rows of the periodic table, the focus remains on the first and second shells.


      Electrons in each shell are drawn in pairs to represent that each orbital (see Schrödinger's model) contains up to two electrons.


      The following diagram shows the electron configuration of three noble gases: helium, neon and argon.

      • Electron configuration of helium = 2 (fully occupied first shell)
      • Electron configuration of neon = 2, 8 (fully occupied second shell)
      • Electron configuration of argon = 2, 8, 8


      Electron configuration of noble gases


      The number of electrons that can be accommodated in any shell of an atom can be determined using the formula `2n^2`, where `n` is the principal quantum number corresponding to the shell's number.

      • For the third shell (`n = 3`): maximum number of electrons 
      • For the fourth shell (`n = 4`): maximum number of electrons 

      It's important to note, however, that in actual atoms, especially those with lower atomic numbers, the electron configuration may not fully "fill up" these outer shells due to the distribution of electrons across different sub-shells (s, p, d, f orbitals) and the effects of electron-electron interactions that are not accounted for in the simpler models like Bohr's. The concept of orbitals is discussed separately in Schrödinger's model of the atom.

      Valence Shells and Valence Electrons 

      The way electrons are configured in an atom's shells significantly influences the atom's chemical behaviour and properties. Electrons in the outermost shell (valence shell), known as valence electrons, play a crucial role in chemical bonding and reactions. Atoms tend to gain, lose, or share electrons to achieve a full outer shell, leading to the formation of ions or molecules.

      For example, atoms of sodium and chlorine have 1 and 7 valence electrons respectively.

      • The electron configuration of sodium atom is 2, 8, 1
      • The electron configuration of chlorine atom is 2, 8, 7



      The arrangement of elements in the periodic table reflects their electron configuration. Elements in the same group (column) have similar valence electron arrangements, giving them similar chemical properties.

      Octet Rule

      The octet rule is a fundamental concept in chemistry that refers to the tendency of atoms to have eight electrons in their valence (outermost) shell to achieve a stable electronic arrangement. This rule is based on the observation that atoms of most elements interact in such a way that they can achieve a valence shell with eight electrons, similar to the electron configuration of noble gases, which are inherently stable due to their complete valence shells.

      The octet rule is derived from the electron configurations of the noble gases, which are elements in Group 18 of the periodic table (helium being an exception with only two electrons in its outer shell). These gases, including neon, argon, and helium, have complete outer electron shells and are noted for their lack of chemical reactivity. Their stable configuration is what many elements aim to achieve through chemical bonding.

      How Atoms Achieve an Octet

      Atoms can achieve a complete octet in three main ways:


      Formation of sodium cation


      1. Ion Formation: Atoms of some elements can achieve an octet by gaining or losing electrons to form ions. Metals, for example, tend to lose electrons to form positively charged cations, whereas non-metals tend to gain electrons to form negatively charged anions. For example, sodium (Na) can lose one electron to achieve an octet in its second shell, forming a Na⁺ ion, while chlorine (Cl) can gain an electron to complete its valence shell, forming a Cl⁻ ion.

      2. Covalent Bonding: Atoms can also achieve an octet by sharing electrons with other atoms, forming covalent bonds. This is common among nonmetals. For instance, two oxygen atoms can share two pairs of electrons, forming a double bond between them, and each oxygen atom ends up with an octet.

      3. Expansion of the Octet: Some elements in the third period and beyond can have more than eight electrons in their valence shell, due to the availability of d orbitals. This is known as the expansion of the octet and is observed in compounds like phosphorus pentachloride (PCl₅) and sulfur hexafluoride (SF₆).

      Limitations of the Octet Rule

      While the octet rule is a useful guideline for understanding chemical bonding and molecule formation, it has its limitations:

      • Exceptions: Not all elements follow the octet rule. Hydrogen (H) and helium (He) are stable with just two electrons in their outer shell, a duet, due to their small size. Transition metals and some heavier elements can have more than eight electrons in their valence shell.
      • Incomplete Octets: Some elements and compounds are stable with fewer than eight electrons in the valence shell, such as boron in BF₃ or beryllium in BeCl₂.

      Despite these exceptions, the octet rule remains a fundamental principle in chemistry, especially useful for predicting the behavior of elements in the main groups of the periodic table during chemical reactions and when forming molecules.