Single and multiple covalent bonds

Octet rule - Matter always wants to be in the most stable form. For any atom, stability is achieved by following the octet rule, which is to say all atoms (with a few exceptions) want 8 electrons in their outermost electron shell (just like noble gases). The electrons present in the outermost shell of an atom are called valence electrons.

Exceptions to the octet rule include hydrogen (H) and helium (He) that follow the duet rule instead. They are the first two elements of the periodic table and have a single electron shell which accommodates only 2 electrons. Other exceptions include some group 3 elements like boron (B) that contain three valence electrons. Theoretically, boron can accommodate five more electrons according to the octet rule, but boron is a very small atom and five non-metal atoms (like hydrogen) cannot pack around the boron nucleus. Thus, boron commonly forms three bonds, BH$\text{}_{3}$start text, end text, start subscript, 3, end subscript, with a total of six electrons in the outermost shell. This also results in some anomalous properties for boron compounds because they are kind of “short of electrons”. It should be thus noted that covalent bonding between non-metals can occur to form compounds with less than an octet on each atom.

In general, achieving the octet configuration (i.e. 8 electrons in the outermost shell) is the driving force for chemical bonding between atoms. Take a look at the outer shell configuration (i.e. number of valence electrons) of three atoms – sodium (Na), chlorine (Cl) and neon (Ne):

Let’s look at the following two scenarios A and B. There are two kids, Emily and Sarah. They both are very good friends.

Scenario B:

Now let’s apply the above analogy to chemical bonding. Assume that Emily and Sarah represent two atoms, and the blanket symbolizes their valence electrons. In scenario A, atom Emily is willing to donate her electrons (blanket) to atom Sarah because by doing so both achieve an octet configuration of 8 electrons in their respective outer shells, making them both happy and stable. This donation of electrons is called ionic bonding.

Example of an ionic bond

In scenario B, both the atoms Emily and Sarah are equally electronegative. So, neither Emily nor Sarah is ready to part with her electrons (blanket), and they instead share their valence electrons with each other. This is called a covalent bond. Electronegativity is a measure of how strongly an atom attracts electrons from another atom in a chemical bond and this value is governed by where the particular atom is located in the periodic table (francium is the least electronegative element while fluorine is the most electronegative).

Let’s go back to Emily and Sarah:

Scenario C:

Scenario D:

Let’s apply the above analogy to a covalent bond formation. In scenario C, both Emily and Sarah are equally cold (in our analogy this translates to them having the same electronegativity). Because they have the same electronegativity, they will share their valence electrons equally with each other. This type of a covalent bond where electrons are shared equally between two atoms is called a non-polar covalent bond.

Example of a non-polar covalent bond

In scenario D, Emily is cold but Sarah is much colder (no doubt mild hypothermia from playing outside in the rain too long)! Together they share the blanket, but Sarah has a tendency to keep pulling the blanket from Emily in order to warm up more. In the atomic world, one atom (Sarah) is more electronegative than another atom (Emily), and naturally pulls the shared electrons towards itself. This pulling of electrons creates slight polarity in the bond. Covalent bonds where electrons are not shared equally between two atoms are called polar covalent bond.

Example of a polar covalent bond

As shown above, the electrons in a covalent bond between two different atoms (H and Cl in this case) are not equally shared by the atoms. This is due to the electronegativity difference between the two atoms. The more electronegative atom (Cl) has greater share of the electrons than the less electronegative atom (H). Consequently, the atom that has the greater share of the bonding electrons bears a partial negative charge (δ-) and the other atom automatically bears a partial positive charge (δ+) of equal magnitude.

Methane (CH$\text{}_{4}$start text, end text, start subscript, 4, end subscript) is an example of a compound where non-polar covalent bonds are formed between two different atoms. One carbon atom forms four covalent bonds with four hydrogen atoms by sharing a pair of electrons between itself and each hydrogen (H) atom. The electronegativity value for carbon (C) and hydrogen (H) is 2.55 and 2.1 respectively, so the difference in their electronegativity values is only 0.45 (<0.5 criteria); the electrons are thus equally shared between carbon and hydrogen. So we can conveniently say that a molecule of methane has a total of four non-polar covalent bonds.

Single and Multiple Covalent Bonds

The number of pairs of electrons shared between two atoms determines the type of the covalent bond formed between them.

Now let’s move on to a couple of examples and try to determine the type of covalent bonds formed

Nitrogen atom can attain an octet configuration by sharing three electrons with another nitrogen atom, forming a triple bond (three pairs of electrons shared)

Consider the molecule carbon dioxide (CO$\text{}_{2}$start text, end text, start subscript, 2, end subscript). Let’s determine the type of covalent bonds it forms.