The main type of bonding between two metals is so-called metal bond. The metal atoms give off all their valence electrons and thus reach the noble gas configuration.

The metal atoms become positively charged cations upon release of the electrons. Between these positively charged cations, the released electrons form the so-called electron gas, since the electrons can move freely in the atomic structure as in a gas so to speak. The cohesion of the atoms is due to the electrostatic attraction between the positively charged cations and the negatively charged electron gas. As the name implies, this type of binding has a special significance, especially in the case of metals.

The free mobility of the electrons in the electron gas is ultimately the cause of the generally good electrical and thermal conductivity of metals (the exception to this property is the group of so-called metalloids). The mutual repulsive forces of the metal cations and the simultaneous attracting force of the electron gas lead to a regular lattice structure.

In contrast to the lattice structure of an ionic bond, which consist of anions or cations, the atomic structure of the metal bond is completely identical. When single atoms or entire atomic series are displaced, there are basically no changes in the atomic structure in a metal.

In contrast to this, opposed charged ions encounter one another when shifting the ionic lattice. The repulsive forces between the identical ions finally “shatter” the material. This is the reason why ceramics are much more brittle due to their ionic bonding and can not be deformed like metals.

The ionic bond is the predominant type of bonding between a metal and a nonmetal. The metal atoms involved in the binding release their valence electrons, which are taken up by the nonmetal atoms. In both cases, the noble gas configuration for the respective atoms is achieved.

The metal atom becomes a positively charged ion (cation) after the release of the electrons. The non-metal atom becomes a negatively charged ion (anion) after the electrons are taken up. The cohesion between the metal and non-metal atoms is due to the electrostatic attraction of the resulting ions. The ionic bond has special significance for ceramics.

In an ionic bond, the metal atoms release their valence electrons, which are taken up the nonmetal atoms in order to reach the noble gas configuration for each atom!

Such crystalline solid compounds of anions and cations are often referred to as salts. A typical example of an ionic compound is therefore common salt, also known as table salt (NaCl). In this compound, the sodium atoms (Na) as alkali metals give off their only valence electrons. The metal atoms thus lose their third M shell, so that on the underlying L shell, the noble gas configuration with eight outer electrons is obtained. The emitted electrons of the sodium atoms are taken up by the nonmetallic chlorine atoms. The chlorine atoms with their seven outer electrons thus now bind eight outer electrons around each other, thus achieving the noble gas configuration. As a result, a crystal structure (ionic lattice structure) ist obtained by the attractive forces between the positive sodium atoms and the negative chlorine atoms.

Salts are ionic compounds consisting of anions and cations!

Basically, the tendency of an atom to bind electrons to itself is particularly great when only a few external electrons are missing to achieve the noble gas configuration. This applies in particular to the elements of the group of halogens with seven valence electrons (e.g. chlorine). Conversely, the tendency to take up electrons is low for those atoms which have only a small number of valence electrons. For these atoms it is usually energetically cheaper to donate the few electrons instead of take up so many. This is especially true for the group of alkali metals whose atoms have only one external electron each (e.g. Na).

Such a more or less pronounced tendency of atoms to attract additional electrons in the binding case is referred to electronegativity. For the above reasons, the electronegativity increases from left to right within a period of the periodic table. This can also be explained by the fact that the number of protons and thus the positive charge of the nucleus increases with increasing atomic number. This increases the atom’s ability to bind electrons as well. The values ​​for the electronegativity of the chemical elements are shown in the figure below (the darker the red the higher the electronegativity).

On the other hand the electronegativity usually decreases from top to bottom within a group. The reason for this is the greater distance of the outermost shell from the atomic nucleus, since with each period a new shell is added. Due to the greater distance, the attractive force between the nucleus and the valence electrons is reduced. Accordingly, the ability of the atoms to attract more electrons decreases. Note that the noble gases can not be assigned electronegativity because they do not bind or tend to donate electrons. The artificially generated elements can as well not be assigned electronegativity for practical reasons, since they are not stable and thus cannot be examined.

The tendency of atoms to attract electrons in the binding case is referred to as electronegativity. In the periodc table the electronegativity usually increases from left to right but decreases from top to bottom!

Whether an ionic bond wil be occur or not depends on the property of the atoms involved to release or take up electrons. The non-metal atoms should have the highest possible tendency to take up electrons (high electronegativity), while the metal atoms should tend to electron donation (low electronegativity). For this reason, it can be concluded of their ionic character from the difference of the electronegativity values of the two elements.

If the difference in the electronegativities of two chemical elements is greater than 1.8, then mainly an ionic bond will be present. On the other hand, if the difference is less than 1.8, a covalent bond is predominantly to be expected. In principle, however, there is always a certain amount of covalent behavior in ionic bonding. For example, sodium chloride has an ionic binding character of about 75%. The remaining 25% is accounted for a covalent part.

If the difference in the electronegativities of two chemical elements is greater than 1.8, then mainly an ionic bond will be present!

The covalent bond mainly occurs between two non-metals. The atoms involved in the bond, share together (“co”) valence electrons (“valent”) to reach the noble gas configuration. For that reason, this type of binding is also called covalent bond. Somewhat imprecisely this is often referred to as molecular bond. In addition to gases, this type of bonding has a special significance for plastics and ceramics.

In a covalent bond, the involved atoms use common valence electrons to reach the noble gas configuration!

With the understanding of covalent bonding, the reason why hydrogen does not occur in nature as a single H atom but always as an H₂ molecule (elementary hydrogen) is finally revealed. Two hydrogen atoms can share their outer electrons. This achieves the noble gas configuration of an helium atom with its two valence electrons.

For the same reason, chlorine particles always occur at a microscopic scale not as a single atom but as a Cl₂ molecule (elementary chlorine). In addition to these examples, the figure above shows the covalent bonding of four hydrogen atoms (4 H) and one carbon atom (C) to one methane molecule (CH₄). Also shown is the covalent bond of a water molecule (H₂O) consisting of one oxygen atom (O) and two hydrogen atoms (2 H). Note that in the figure, only the valence electrons of the atoms are shown, as only these are responsible for the chemical bond.

The representation of covalent bonds shown in the figure above is indeed very clear, but in many cases too costly or especially for double bonds not even possible. For this reason, one uses the so-called structural formula. The valence electrons of the individual atoms are represented by a dot next to the element symbol. For a hydrogen atom with only one valence electron, there is thus a point to the right of the H symbol. In the case of four external electrons, such as the carbon atom, on the other hand, there is one point to the right, left, above and below the element symbol. For atoms with more than four valence electrons, one electron is added to each side. Such a notation is also referred to as Lewis dot structure.

The Lewis dot structure is a structural formula which discribes the covalent bond. Binding electrons are represented by a dash between the atoms.

If atoms now enter into covalent bonds, the electrons involved in the binding are combined with a dash. These electrons are referred to as bonding electrons (bonding pairs or shared pairs). In the same way, the remaining pairs of electrons of an atom are also connected by a dash. However, these are no longer bonding electrons, since they have nothing to do with the bond itself. With the help of the Lewis dot structure one obtains a very clear information about the binding of the molecule.

The figure above shows very clear the noble gas configuration, which is formed by the surrounding electron pairs around an element. Thus, in all cases shown, either the noble gas configuration of helium with two electrons (or an electron pair) or the noble gas configuration with eight valence electrons (or four electron pairs). With the Lewis dot structure, double bonds can also be illustrated very easily, as is the case with carbon dioxide (CO₂) or ethylene (C₂H₄).

The notation describing only the type and number of atoms occurring in a molecule (for example H₂, CO₂, C₂H₄), etc.) is called the molecular formula!

In nature, substances rarely appear as pure elements. Much more, different elements bond for energetic reasons with each other and form chemical compounds. A typical example of this is water (H₂O). In this case, two hydrogen atoms (2H) combine with one oxygen atom (O) to form a stable water molecule:

\begin{align}
\label{wasser}
& 2 H ~+~ O ~\rightarrow ~ H_2O \\[5px]
\end{align}

On the other hand, bringing the two hydrogen atoms into contact with an argon atom (Ar) will not result in a stable bonding between these elements. Rather, the argon atom remains for itself and the two hydrogen atoms combine to form elementary hydrogen (H₂):

\begin{align}
\label{argon}
& 2 H ~+~ Ar ~\rightarrow ~ H_2 ~+~ Ar \\[5px]
\end{align}

If one considers the chemical elements with regard to their bonding behavior, it is noticeable that the elements of the 8th main group in the periodic table, are particularly stable. They take virtually no chemical reactions with other atoms and therefore do not form molecules. In nature they occur only monatomic (that means as single atoms).

For this reason, the argon atom does not bind chemically with the two hydrogen atoms. The elements of the 8th main group are all gaseous at room temperature, which gives this group the name noble gases. Based on their “sluggish” chemical behavior, these gases are also referred to as inert gases. Since these noble gases practically do not react with other substances, some of them are used as shielding gases against unwanted oxidation during welding.

Since the number of outermost electrons decisively influence the chemical behavior of an atom, the number of eight so-called valence electrons (or two in case of helium) means a particularly stable electron occupation. This electron configuration is obviously very favorable in terms of energy. The experimental investigations of the chemical bonding behavior of different atoms also confirm this assumption. This shows that atoms always try to form chemical bonds in such a way that eight or two valence electrons form around the atoms involved.

This state of an atom within a chemical bond with eight or two external electrons is also referred to as an noble gas configuration. Thus, an important rule for the chemical bonding behavior can be derived:

Each atom strives to reach the closest noble gas configuration in the periodic table (octet rule).

Based on the noble gases with their eight valence electrons (exception: helium), the effort to achieve the noble gas configuration is also called octet rule. The noble gas state is achieved by the fact that the atoms form chemical bonds and thereby

• absorb or release electrons (ionic bond, metal bond), or
• use it together with other atoms (covalent bond).

The respective chapters will briefly explain these most important types of bonding. It should always be noted that in reality, bonds can not be sharply limited to a certain type of bond. Rather, chemical compounds have features of different types of bonds.

Important note

in the context of Bohr’s atomic model, it is often incorrectly claimed that the noble gas configuration means a fully occupied outermost shell. This statement is wrong! For in the Bohr model, the maximum number of electrons on the \(n\)^th shell results from the following equation:

\begin{equation}
N_{max} = 2 \cdot n^2
\end{equation}

With \(n\) = 3 argon offers space on its outermost third shell for a maximum of \(N\) = 18 electrons. However, argon only has 8 valence electrons on this shell. The outermost shell is therefore far from being fully occupied! At this point it is argued with the wrong atomic model. Rather, the statement of the fully occupied (sub)shell is to be seen in connection with the orbitals that Sommerfeld introduced as an enhancement of the Bohr model (Bohr-Sommerfeld model):

The noble gas configuration means a fully occupied orbital (subshell) in the Bohr-Sommerfeld model!

More specifically, the noble gas configuration means a fully occupied p orbital. The exception is helium with a fully occupied s orbital.