Polar and Non-polar Covalent Bonds — Definition

Definition

Imagine two friends sharing a toy. If they are equally strong and pull with the same force, the toy stays exactly in the middle – this is like a non-polar covalent bond. Now, if one friend is much stronger and pulls harder, the toy moves closer to them, even though both are still 'sharing' it – this is like a polar covalent bond. In chemistry, atoms share electrons to form covalent bonds. The 'strength' with which an atom pulls on these shared electrons is called its electronegativity.

Non-polar Covalent Bonds: These bonds occur when two atoms share electrons *equally*. This happens primarily in two situations:

1
Between identical atoms: — For example, in a hydrogen molecule ($H_2$), both hydrogen atoms have the same electronegativity. They pull on the shared electron pair with equal strength, so the electrons are perfectly in the middle. Other examples include $O_2$, $N_2$, $Cl_2$, $F_2$, etc.
2
Between atoms with a very small or negligible electronegativity difference: — Sometimes, even if the atoms are different, their electronegativity values are so close that the sharing is considered practically equal. For instance, a carbon-hydrogen (C-H) bond is often considered non-polar because the electronegativity difference between carbon (2.55) and hydrogen (2.20) is quite small (0.35).

In a non-polar bond, there is no separation of charge; no part of the bond is significantly more positive or negative than the other.

Polar Covalent Bonds: These bonds occur when two atoms share electrons *unequally*. This happens when there is a *significant difference* in electronegativity between the two bonded atoms. The atom with higher electronegativity pulls the shared electron pair closer to itself, becoming slightly negative (denoted as $delta^-$), while the other atom becomes slightly positive (denoted as $delta^+$).

Think of a water molecule ($H_2O$). Oxygen is much more electronegative than hydrogen. So, in each O-H bond, oxygen pulls the shared electrons closer, making the oxygen atom partially negative ($delta^-$) and each hydrogen atom partially positive ($delta^+$). Other common examples include $HCl$, $HF$, $HBr$, $NH_3$, etc.

This separation of charge within a bond creates what we call a 'dipole' – essentially, two poles, one positive and one negative. The magnitude of this polarity is measured by something called the 'dipole moment'.

It's crucial to understand that while a bond can be polar, the overall molecule might still be non-polar if its shape causes these individual bond dipoles to cancel each other out, like in carbon dioxide ($CO_2$) or carbon tetrachloride ($CCl_4$).

So, we must distinguish between bond polarity and molecular polarity.