Polarity in Molecules (VCE SSCE Chemistry): Revision Notes

Polarity in Molecules

Introduction

Have you ever noticed how water behaves strangely around static electricity? When you bring a statically charged balloon near a stream of water, the water actually bends towards the balloon. This fascinating phenomenon occurs because water molecules have an uneven distribution of electrons, making them partially charged. In this note, you'll learn how the shape of molecules and the electronegativity of atoms create these charge distributions.

What is electronegativity?

Electronegativity describes how strongly an atom attracts electrons when it forms a chemical bond. Think of it as a tug-of-war between atoms—the more electronegative an atom is, the harder it pulls on the shared electrons.

Trends in electronegativity

Electronegativity follows clear patterns across the periodic table:

• Increases from left to right across each period
• Decreases down each group

This means fluorine (in the top right of the periodic table) is the most electronegative element, whilst elements like caesium (bottom left) have very low electronegativity values.

Understanding polar and non-polar bonds

Non-polar bonds

When two identical atoms form a covalent bond, they have equal electronegativity. This means the bonding electrons are shared equally between them. We call these non-polar bonds because there's no charge difference across the bond.

Examples of non-polar bonds occur in diatomic molecules like:

• Chlorine ($\text{Cl}_2$)
• Oxygen ($\text{O}_2$)
• Hydrogen ($\text{H}_2$)
• Nitrogen ($\text{N}_2$)

Electron density refers to the probability of finding an electron at a particular location. In non-polar molecules like fluorine ($\text{F}_2$), the electron density is distributed symmetrically between the two atoms.

Polar bonds

When two different atoms form a covalent bond, they usually have different electronegativities. The more electronegative atom pulls the shared electrons closer to itself, creating an unequal distribution. These are called polar bonds.

Consider hydrogen fluoride (HF). Fluorine is much more electronegative than hydrogen, so the bonding electrons spend more time near the fluorine atom. This creates:

• A partial negative charge ($\delta^-$) on the fluorine atom
• A partial positive charge ($\delta^+$) on the hydrogen atom

Comparing bond polarity

The greater the electronegativity difference between two bonded atoms, the more polar the bond becomes. You can calculate and compare bond polarity using electronegativity values.

The spectrum of bond types

Chemical bonds exist on a continuum from non-polar covalent through polar covalent to ionic bonding. The electronegativity difference between atoms determines where a bond falls on this spectrum.

Non-polar covalent bonds

• Electronegativity difference: Zero
• Electron behaviour: Shared equally
• Examples: $\text{F}_2$, $\text{H}_2$, $\text{CS}_2$ (carbon and sulfur both have electronegativity 2.6)

Polar covalent bonds

• Electronegativity difference: Between 0 and 1.7
• Electron behaviour: Attracted more strongly to the more electronegative atom
• Examples: $\text{NH}_3$, $\text{H}_2\text{O}$, HCl

Ionic bonds

• Electronegativity difference: Greater than 1.7
• Electron behaviour: Completely transferred from one atom to another
• Examples: NaCl, $\text{CaF}_2$

Polarity in polyatomic molecules

For molecules with more than two atoms, determining polarity becomes more complex. The overall polarity depends on both the polarity of individual bonds AND the shape of the molecule.

Key principles

• Molecules containing only non-polar bonds are always non-polar
• Symmetrical molecules with polar bonds are non-polar because the individual dipoles cancel out
• Asymmetrical molecules with polar bonds are polar because they have a net dipole

Non-polar polyatomic molecules

Methane ($\text{CH}_4$) provides an excellent example. Carbon is more electronegative than hydrogen, so each C-H bond is polar with $\delta^-$ on carbon and $\delta^+$ on hydrogen. However, methane has a tetrahedral shape that's perfectly symmetrical.

When you represent each bond dipole as an arrow (pointing from positive to negative), these arrows point symmetrically in all directions. They cancel each other out completely, resulting in no net dipole. Therefore, methane is non-polar despite having polar bonds.

Other examples of symmetrical non-polar molecules include:

• Carbon dioxide ($\text{CO}_2$) - linear shape
• Tetrafluoromethane ($\text{CF}_4$) - tetrahedral shape

Polar polyatomic molecules

In asymmetrical molecules, the individual bond dipoles don't cancel out. This creates a net dipole, making the molecule polar overall.

Chloromethane ($\text{CH}_3\text{Cl}$) demonstrates this well. The chlorine atom is more electronegative than carbon, which is more electronegative than hydrogen. The molecule has a tetrahedral shape, but it's not symmetrical because one position has chlorine instead of hydrogen.

When you add up all the individual bond dipoles, they don't cancel completely. The result is a net dipole pointing from the hydrogen end towards the chlorine end, making chloromethane a polar molecule.

Examples of polar and non-polar molecules

Water ($\text{H}_2\text{O}$) and ammonia ($\text{NH}_3$) are both asymmetrical and polar:

Real-world application: microwave ovens

Microwave ovens use the polarity of molecules to heat food. The microwaves create an electric field that interacts with polar molecules, especially water. This electric field causes the polar molecules to rotate billions of times per second, increasing their kinetic energy. As kinetic energy increases, temperature rises, heating and cooking the food.

Remember!