Reading: Electrons

Electron Shells and the Bohr Model

Three concentric circles around the nucleus of a hydrogen atom represent principal shells. These are named 1n, 2n, and 3n in order of increasing distance from the nucleus. An electron orbits in the shell closest to the nucleus, 1n.

Figure 1. The Bohr model.

It should be stressed that there is a connection between the number of protons in an element, the atomic number that distinguishes one element from another, and the number of electrons it has. In all electrically neutral atoms, the number of electrons is the same as the number of protons. Thus, each element, at least when electrically neutral, has a characteristic number of electrons equal to its atomic number.

An early model of the atom was developed in 1913 by Danish scientist Niels Bohr (1885–1962). In this model, electrons exist within principal shells. An electron normally exists in the lowest energy shell available, which is the one closest to the nucleus. Energy from a photon of light can bump it up to a higher energy shell, but this situation is unstable, and the electron quickly decays back to the ground state. In the process, a photon of light is released.

The Bohr model shows the atom as a central nucleus containing protons and neutrons, with the electrons in circular orbitals at specific distances from the nucleus, as illustrated in Figure 1. These orbits form electron shells or energy levels, which are a way of visualizing the number of electrons in the outermost shells. These energy levels are designated by a number and the symbol “n.” For example, 1n represents the first energy level located closest to the nucleus.

Electrons fill orbitals in a consistent order: they first fill the orbitals closest to the nucleus, then they continue to fill orbitals of increasing energy further from the nucleus. If there are multiple orbitals of equal energy, they will be filled with one electron in each energy level before a second electron is added. The electrons of the outermost energy level determine the energetic stability of the atom and its tendency to form chemical bonds with other atoms to form molecules.

Under standard conditions, atoms fill the inner shells first, often resulting in a variable number of electrons in the outermost shell. The innermost shell has a maximum of two electrons but the next two electron shells can each have a maximum of eight electrons. This is known as the octet rule, which states, with the exception of the innermost shell, that atoms are more stable energetically when they have eight electrons in their valence shell, the outermost electron shell. Examples of some neutral atoms and their electron configurations are shown in Figure 2. Notice that in this figure, helium has a complete outer electron shell, with two electrons filling its first and only shell. Similarly, neon has a complete outer 2n shell containing eight electrons. In contrast, chlorine and sodium have seven and one in their outer shells, respectively, but theoretically they would be more energetically stable if they followed the octet rule and had eight.

Practice

Bohr diagrams of elements from groups 1, 14, 17 and 18, and periods 1, 2 and 3 are shown. Period 1, in which the 1n shell is filling, contains hydrogen and helium. Hydrogen, in group 1, has one valence electron. Helium, in group 18, has two valence electrons. The 1n shell holds a maximum of two electrons, so the shell is full and the electron configuration is stable. Period 2, in which the 2n shell is filling, contains lithium, carbon, fluorine, and neon. Lithium, in group 1, has 1 valence electron. Carbon, in group 14, has 4 valence electrons. Fluorine, in group 17, has 7 valence electrons. Neon, in group 18, has 8 valence electrons, a full octet. Period 3, in which the 3n shell is filling, contains sodium, silicon, chlorine, and argon. Sodium, in group 1, has 1 valence electron. Silicon, in group 14, has 4 valence electrons. Chlorine, in group 17, has 7 valence electrons. Argon, in group 18, has 8 valence electrons, a full octet.

Figure 2. Bohr diagrams for hydrogen, helium, lithium, carbon, fluorine, neon, sodium, silicon, chlorine, and argon.

 

Bohr diagrams indicate how many electrons fill each principal shell. Group 18 elements (helium, neon, and argon are shown in Figure 2) have a full outer, or valence, shell. A full valence shell is the most stable electron configuration. Elements in other groups have partially filled valence shells and gain or lose electrons to achieve a stable electron configuration.

An atom may give, take, or share electrons with another atom to achieve a full valence shell, the most stable electron configuration. Looking at this figure, how many electrons do elements in group 1 need to lose in order to achieve a stable electron configuration? How many electrons do elements in groups 14 and 17 need to gain to achieve a stable configuration?

Understanding that the organization of the periodic table is based on the total number of protons (and electrons) helps us know how electrons are distributed among the outer shell. The periodic table is arranged in columns and rows based on the number of electrons and where these electrons are located. Take a closer look at the some of the elements in the periodic table’s far right column in Figure 3.

The periodic table consists of eighteen groups and seven periods. Two additional rows of elements, known as the lanthanides and actinides, are placed beneath the main table. The lanthanides include elements 57 through 71 and belong in period seven between groups three and four. The actinides include elements 89 through 98 and belong in period eight between the same groups. These elements are placed separately to make the table more compact. For each element, the name, atomic symbol, atomic number, and atomic mass are provided. The atomic number is a whole number that represents the number of protons. The atomic mass, which is the average mass of different isotopes, is estimated to two decimal places. For example, hydrogen has the atomic symbol H, the atomic number 1, and an atomic mass of 1.01. The atomic mass is always larger that the atomic number. For most small elements, the atomic mass is approximately double the atomic number as the number of protons and neutrons is about equal. The elements are divided into three categories: metals, nonmetals and metalloids. These form a diagonal line from period two, group thirteen to period seven, group sixteen. All elements to the left of the metalloids are metals, and all elements to the right are nonmetals.

Figure 3. The periodic table shows the atomic mass and atomic number of each element. The atomic number appears above the symbol for the element and the approximate atomic mass appears below it.

The group 18 atoms helium (He), neon (Ne), and argon (Ar) all have filled outer electron shells, making it unnecessary for them to share electrons with other atoms to attain stability; they are highly stable as single atoms. Their non-reactivity has resulted in their being named the inert gases (or noble gases). Compare this to the group 1 elements in the left-hand column. These elements, including hydrogen (H), lithium (Li), and sodium (Na), all have one electron in their outermost shells. That means that they can achieve a stable configuration and a filled outer shell by donating or sharing one electron with another atom or a molecule such as water. Hydrogen will donate or share its electron to achieve this configuration, while lithium and sodium will donate their electron to become stable. As a result of losing a negatively charged electron, they become positively charged ions. Group 17 elements, including fluorine and chlorine, have seven electrons in their outmost shells, so they tend to fill this shell with an electron from other atoms or molecules, making them negatively charged ions. Group 14 elements, of which carbon is the most important to living systems, have four electrons in their outer shell allowing them to make several covalent bonds (discussed below) with other atoms. Thus, the columns of the periodic table represent the potential shared state of these elements’ outer electron shells that is responsible for their similar chemical characteristics.

Video Review

This video gives another overview of the electron:

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