As stated earlier, calcium is in Group 2. Therefore, to gain a full outer shell, it needs to lose two electrons. In doing this, it will have a 2+ charge.
Fluorine is in Group 7. Therefore, to gain a full outer shell, it needs to gain one electron. By doing this, it has a negative charge. To balance the two electrons given up by calcium, there need to be two fluorine atoms present that can each take one electron.
2.1.4 Covalent bonding
Covalent bonding occurs between two non‐metal atoms. In a covalent bond, each atom contributes one or more electrons to share with the other atom. The driving force behind covalent bonding is similar to that in ionic bonding: each element wants to achieve a full outer shell of electrons. In the case of covalent bonding, electrons are shared rather than lost or gained. The covalent bond itself is very strong; a large amount of energy is required to break a covalent bond.
Covalent bonding occurs between non‐metals. A covalent bond is a shared pair of electrons.
We will use the formation of the H—H bond in molecular hydrogen to explain covalent bonding. Hydrogen is an element that can undergo covalent bonding. In an atom of hydrogen, there is only one electron in a 1s orbital (the outermost shell). To achieve a full outer‐shell, hydrogen needs to have two electrons; therefore, one hydrogen atom shares its electron with that of another atom of hydrogen. The 1s orbitals containing the electrons overlap. This is shown in Figure 2.8a. In this figure, the electrons on each hydrogen are shown as a dot or a cross. When the electrons are shared, they are shown within the overlapping rings of the outermost shell. When two electrons are shared, they are said to form a bonding pair between the two atoms. One pair of electrons is one bond and can be represented by a single line, as in Figure 2.8b. In the case of hydrogen, H2, there is one pair of electrons that is located between the two nuclei. This is called a (sigma) σ bond. Figure 2.8c is a representation of the electron cloud showing where the two electrons in a σ bond are located. This will be discussed further in Chapter 12.
The symbol σ is the Greek letter sigma. A σ bond consists of a pair of electrons shared between two covalently bonded atoms.
Figure 2.8 (a) Formation of a single bond in hydrogen, H2; (b) an alternative representation of a hydrogen molecule with a single bond; (c) a representation of the location of the electrons in a σ bond.
Although the electrons in the hydrogen atoms are shown as dots and crosses, all the electrons are identical.
The previous example for hydrogen, H2 shows the formation of a single covalent bond where two electrons are shared between the two atoms. In many cases, atoms must share more than two electrons. In such cases, a double or triple bond is formed. Double and triple bonds are made up of sigma (σ) and pi (π) bonds, which will be discussed further in Chapter 12. An example of a molecule where a double bond is present is oxygen, O2. This is shown in Figure 2.9.
Figure 2.9 (a) Formation of a double bond in oxygen, O2; (b) an alternative representation of an oxygen molecule with a double bond.
The symbol π is the Greek letter pi. A π bond consists of a pair of electrons located between two covalently bonded atoms in a plane at 90° to the bond axis.
The oxygen atom has six outer electrons. The electron configuration is 1s22s22p4. An oxygen atom requires two additional electrons to fill its outer shell and attain a complete octet of electrons as in the atom neon, the closest noble gas. Thus, in the formation of an oxygen molecule, each oxygen atom shares two of its outer electrons with another oxygen atom. There is therefore a total of four electrons bonding the oxygen atoms together. When four electrons are shared between two atoms, a double bond is formed, which is represented by a double horizontal line between the atoms in the bond.
In covalent bonding, the first two electrons to be shared form a sigma bond. If more than two electrons are shared, the electrons occupy pi bonds.
Worked Example 2.4
Draw a dot‐and‐cross diagram to show the bonding present in carbon dioxide, CO2.
Solution
Carbon dioxide contains two elements: carbon and oxygen. Both are non‐metals; therefore, the type of bonding between them will be covalent, i.e. shared pairs of electrons. Carbon has four electrons in its outer shell and therefore needs to gain another four electrons to fill its outer shell. Oxygen has six electrons, so it needs to gain two more electrons. In addition, the name carbon dioxide gives a hint about the structure, with the ‘di’ showing that there are two oxygen atoms.
This information allows us to deduce that if carbon shares two electrons with each oxygen, each oxygen will have eight electrons in its outer shell and so will have a complete octet. Conversely, if oxygen shares two electrons each with carbon, carbon will have gained a share of four electrons in total, so it will also have a full outer shell. In the case of carbon dioxide, there are two pairs of electrons between each nucleus (a). Thus, there is a double bond between each carbon and oxygen atom. This is represented by a double line between the atoms (b).
For all structures containing covalent bonds, it is possible to work out the bonding and how the atoms are arranged. Some common molecules are shown in Table 2.1.
Table 2.1 The names, molecular formulae, dot‐and‐cross diagrams, and display formulae of some commonly encountered covalently bonded molecules.
| Name and molecular formula | Dot‐and‐cross diagram | Display formula |
|---|---|---|
| Fluorine, F2 |
|
|