Predict molecular shape from electron-pair repulsion, use electronegativity difference to identify polar bonds, and decide whether a molecule is polar or nonpolar from its shape (MA STE HS-PS1-3 support, structure and polarity).

What this topic is asking

Standard HS-PS1-3 is about the structure of substances and the forces between particles, and that begins with the shape of a molecule and whether it is polar. Massachusetts high school chemistry expects you to predict a molecule's shape from the repulsion between its electron pairs, to spot a polar bond from an electronegativity difference, and to decide whether a whole molecule is polar by combining the bond polarities with the shape. Polarity then explains solubility, boiling points, and the intermolecular forces in the next topic.

Shape from electron-pair repulsion

Counting only the groups around the central atom gives the basic shapes:

• Two electron groups spread to opposite sides: linear (180 degrees), as in $\text{CO}_2$.
• Three groups spread into a flat triangle: trigonal planar (120 degrees), as in $\text{BF}_3$.
• Four groups spread into three dimensions to the corners of a tetrahedron (about 109.5 degrees), as in $\text{CH}_4$.

Lone pairs count toward the repulsion but are not drawn as bonds, and they repel slightly more strongly than bonding pairs, so they squeeze the bond angles. With four electron groups, replacing one bond with a lone pair gives a pyramidal molecule (ammonia, $\text{NH}_3$), and replacing two gives a bent molecule (water, $\text{H}_2\text{O}$). Recognizing the bent shape of water is essential, because it explains water's polarity and many of its special properties.

Bond polarity from electronegativity

A covalent bond shares electrons, but not always equally. Electronegativity (from the periodic table and periodic trends) measures how strongly an atom pulls shared electrons.

• If the two atoms have the same electronegativity (as in $\text{O}_2$), the sharing is equal and the bond is nonpolar.
• If they differ, the more electronegative atom pulls the electrons closer, gaining a partial negative charge ($\delta-$) while the other atom is left partial positive ($\delta+$). This is a polar covalent bond.
• A very large difference means the electrons are essentially transferred, which is an ionic bond.

So bond character runs along a scale from nonpolar covalent, through polar covalent, to ionic, set by the electronegativity difference.

Whether the molecule is polar

A molecule can contain polar bonds yet be nonpolar overall. What decides it is whether the bond polarities cancel, which depends on the shape:

• In a symmetrical molecule, equal bond polarities point in directions that cancel, so the molecule is nonpolar. Carbon dioxide is linear and symmetrical, so its two polar bonds cancel.
• In an unsymmetrical molecule (often because of lone pairs or different atoms), the bond polarities do not cancel and add up to a net direction, so the molecule is polar. Water is bent, so its two polar bonds combine to give a positive end and a negative end.

Polar molecules attract one another more strongly and dissolve in polar solvents like water; nonpolar molecules do not. This "like dissolves like" rule is used in solutions, solubility and concentration.

Try this

Q1. State the shape of a molecule with a central atom surrounded by three bonding pairs and no lone pairs. [1]

• Cue. Trigonal planar (bond angles about 120 degrees).

Q2. Why is carbon dioxide nonpolar even though it has polar bonds? [2]

• Cue. It is linear and symmetrical, so the two equal bond polarities point in opposite directions and cancel.