Polar bonds and electronegativity

Polar bonds and electronegativity

You are probably familiar with ionic bonds and covalent bonds from gcse chemistry. Before looking at polar covalent bonding it might be helpful to quickly re-cap ionic and covalent bonding.

Ionic bonding

Ionic bonding involves the transfer of an electron from a metal to a non-metal. This results in the formation of ions with positive and negative charges. In the example shown below the metal sodium reacts with the non-metal chlorine to form sodium chloride. Here the metal sodium loses its outer shell 3s1 electron to chlorine. This means that the sodium atom forms a sodium ion with a charge of 1+. The non-metal chlorine gains an electron from the sodium and forms a chloride ion with an electron arrangement of 2,8,8 or 1s22s22p62s23p6.

The formation of the ionic compound sodium chloride from sodium metal and chlorine by elecron transfer from the sodium atom to the chlorine atom.

Covalent bonding

Covalent bonding involves the sharing of a pair of electrons, usually between non-metal atoms present in small molecules. However these ideas of ionic and covalent bonding which were introduced in gcse chemistry should really be thought of as the two extremes when it comes to describing the bonding between atoms. If a compound has pure ionic bonds then the ions must be perfectly spherical and the electron must be transferred completely from the metal to the non-metal. In most of the reactions between metals and non-metals these two conditions are rarely met.

Similarly for pure covalent bonding to occur the 2 electrons in the bond must be shared EQUALLY. In a covalent bond between identical elements with identical electronegativity values then the electrons present in any covalent bonds will be shared equally. However in covalent compounds formed between elements with different electronegativities the electrons are likely to be shared unequally.
3D model showing how the electrons are shared in a covalent bond in a chlorine molecule.

Polar covalent bonds

Most covalent compounds involve bonding between elements with different electronegativities. This means that the electrons in any bonds formed will NOT be shared EQUALLY. In an ionic compound the ions formed have positive and negative charges because there is complete transfer of an electron from the metal to the non-metal atom. In a polar covalent bond the electron is shared unequally between the two atoms or we could simply say one atom in the bond has a much larger share of the electron. This means that the atoms will have partial charges and not full positive and negative charges.

The Greek symbol delta (lowercase) δ is used to show the resultant partial positive (δ+) and negative (δ-) and charges that result from the unequal sharing of the electrons in a polar covalent bond. A polar covalent bond or simply a polar bond can be thought as a sort of halfway house between covalent and ionic bonding in that the electrons in the bond are shared but not equally as is the case in a covalent bond and the polar covalent bond contains charges, similar to those found in an ionic compound. However the charges are only partial charges and not full charges as is the case with ionic compounds. The charges on the atoms are only partial charges simply because the electrons are not completely transferred, just shared unequally.

3d model showing how the electrons are shared unequally in a polar covalent bond in a hydrogen chloride molecule.

The larger the difference in the electronegativity between the two atoms in the covalent bond the more polar (polar covalent) the bond will be. If the bond between 2 different atoms has a very large difference in electronegativity values of say 1.7-2.0 or more it will be ionic. If the differences in the electronegativity is around 0.5 or less then the bond is likely to be covalent. Click on the link for bond character above or here for more information.

Hydrogen chloride (HCl)

Consider a molecule of hydrogen chloride (H-Cl). Chlorine has an electronegativity value of 3.0 and hydrogen has a value of 2.1; this means that the bond is polarised in such a way that the chlorine atom is electron rich -) and the hydrogen atom is electron deficient +). Since the H-Cl bond has 2 charged ends it is described as having a bond dipole or just a dipole (a dipole is simply a small charge difference across a bond due to differences in electronegativities, which results in one end of the bond containing an atom with a partial positive charge (δ+ ) and at the other end of the bond will be an atom with a partial negative charge -). The charges on both end of the bond will be the same but different in sign, one negative and one positive. So H-Cl is a non-symmetrical molecule with polar covalent bonds. We would also say that this is a polar molecule (that is one end of the molecule has a partial positive charge (δ+ ) and the other a partial negative charge (δ- ), this is shown in the image below:

3d model showing polar covalent bonding in a molecule of hydrogen chloride (HCl)

Polar molecules

Are all molecules which have polar covalent bonds polar molecules? The answer to this question is NO. To decide if a molecule is polar you not only have to decide if it contains polar covalent bonds but you MUST also consider the shape of the molecule. Bond dipoles are vector quantities and have both a size and a direction. If the molecule has polar bonds but is highly symmetrical the chances are that the molecule will be non-polar even though it contains polar covalent bonds. As an example consider a molecule of carbon tetrachloride (CCl4) which is shown below:

How to find the centres of positive and negative charge in a molecule.  Here the molecule carbon tetrachloride is used as an example to decide if it is a polar molecule or not. 3d model of carbon 
tetrachloride showing how the bond dipoles cancel to give a non-polar molecule

You may also see the carbon tetrachloride molecule drawn out using arrows to represent the bond dipoles (see the image opposite). This is often helpful in deciding if the molecule overall will be a polar one with a dipole moment. The arrows representing the bond dipoles give an indication of the size of the dipole present in any particular bond. I hope you can see that all the bond dipole arrows will effectively cancel each other out due to the shape of this molecule. However if you study the examples below you will see that this is not always the case.

Polar and non-polar molecules

Consider each of the examples below which show both polar and non-polar molecules. In each case you need to consider not only if the molecule has polar covalent bonds but also the shape of the molecule. Molecules which are highly symmetrical may have the centres of positive and negative charge directly on top of each other. If this is the case then the molecule will be non-polar despite having polar covalent bonds.

3d models to show a selection of molecules which are polar and non-polar bases on their bonding and also their shape.

Key Points

Practice questions

Check your understanding - Questions on polar bonds and electronegativity