While covalent bonds are often described as the simple sharing of a pair of electrons, this sharing is rarely completely equal. Different elements have different abilities to attract electrons, leading to polar bonds and dipole moments that dictate how molecules interact with one another.
🔑 Key Principle
Bond polarity is a spectrum. If two atoms have the same electronegativity, the bond is pure covalent (electrons are shared equally). If there is a moderate difference, the bond is polar covalent. If the difference is very large, electrons are transferred completely, resulting in an ionic bond.
What is Electronegativity?
Electronegativity is a measure of an atom's ability to attract the shared pair of electrons in a covalent bond. It is measured on the Pauling scale, where fluorine is the most electronegative element with a value of 4.0.
The power of an atom to attract the pair of electrons in a covalent bond towards itself.
Electronegativity depends on three main factors:
- Nuclear charge: The more protons in the nucleus, the stronger the attraction for the bonding electrons.
- Atomic radius: The closer the bonding electrons are to the nucleus, the stronger the electrostatic attraction.
- Shielding: Inner shells of electrons shield the outer bonding electrons from the attraction of the nucleus.
When defining electronegativity, you must mention that it refers to attracting electrons in a covalent bond. If you define it simply as the ability of an atom to attract electrons, you will lose the mark, as this could describe electron affinity.
Periodic Trends in Electronegativity
Electronegativity varies systematically across the periodic table:
- Across a Period (increases): The nuclear charge increases (more protons), while shielding remains constant because electrons are added to the same main energy shell. The atomic radius also decreases, bringing the bonding pair closer to the nucleus.
- Down a Group (decreases): Although nuclear charge increases, atomic radius increases significantly due to the addition of extra electron shells. Shielding also increases. This screens the bonding pair from the nucleus, resulting in a weaker attraction.
Bond Polarity
When two atoms with different electronegativities form a covalent bond, the bonding electrons are pulled closer to the more electronegative atom. This creates an unequal distribution of charge.
A covalent bond in which the electron density is distributed unsymmetrically, resulting in partial positive (\( \delta+ \)) and partial negative (\( \delta- \)) charges on the bonded atoms.
A separation of charge, consisting of equal but opposite charges separated by a small distance. In a bond, it is represented by an arrow pointing towards the more electronegative element.
Bond Polarity vs Molecular Polarity
It is vital to distinguish between a polar bond and a polar molecule. A molecule can contain polar bonds but still be entirely non-polar overall if the dipoles are arranged symmetrically.
Symmetrical Molecules (Non-polar)
If the bond dipoles are of equal strength and act in opposing directions, they cancel each other out, resulting in no overall dipole moment.
Examples: \( \text{CO}_2 \) (linear), \( \text{BF}_3 \) (trigonal planar), \( \text{CCl}_4 \) (tetrahedral).
Asymmetrical Molecules (Polar)
If the dipoles do not cancel (often due to asymmetric shapes or the presence of lone pairs), the molecule has a net dipole moment.
Examples: \( \text{H}_2\text{O} \) (bent), \( \text{NH}_3 \) (trigonal pyramidal), \( \text{CH}_3\text{Cl} \) (tetrahedral but asymmetric).
In exam questions asking why a molecule is non-polar despite containing polar bonds, you must explicitly state that the dipoles cancel out because of the symmetrical shape of the molecule. If you just say the shape is symmetrical, you will miss the explanation mark.
Worked Examples
Step 1: Compare electronegativities and identify polar bonds.
Oxygen is more electronegative than carbon and hydrogen. Both molecules contain polar bonds: \( \delta-\text{ O}=\text{C}^{\delta+}=\text{O}^{\delta-} \) and \( \text{H}^{\delta+}-\text{O}^{\delta-}-\text{H}^{\delta+} \).
Step 2: Examine the molecular geometry.
- \( \text{CO}_2 \) is linear. The dipoles from each \( \text{C}=\text{O} \) bond act in exactly opposite directions. Because the molecule is symmetrical, the dipoles cancel each other out, leaving no overall dipole moment.
- \( \text{H}_2\text{O} \) is bent (V-shaped) due to the two lone pairs on the oxygen atom. The dipoles do not cancel each other out. This results in a net dipole pointing towards the oxygen atom, making water a polar molecule.
Step 1: Check bond polarity.
Chlorine is more electronegative than carbon, so each \( \text{C}-\text{Cl} \) bond is polar covalent, with the dipole pointing towards the chlorine atom.
Step 2: Check molecular geometry.
\( \text{CCl}_4 \) has a tetrahedral shape. Since all four outer atoms are identical chlorine atoms, the dipoles are arranged symmetrically in three dimensions. As a result, they pull equally in all directions and cancel each other out, making \( \text{CCl}_4 \) a non-polar molecule.
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