Compare the strengths of intermolecular forces (dispersion, dipole-dipole, hydrogen bonding) and the bonds in ionic and network solids, and use them to explain bulk properties (MA STE HS-PS1-3, structure and forces between particles).

What this topic is asking

Standard HS-PS1-3 asks you to compare the structure of substances at the bulk scale to infer the strength of the electrical forces between particles. In practice that means linking what you can measure (melting point, boiling point, hardness, conductivity, solubility) to the type and strength of the forces holding the particles together. Massachusetts high school chemistry wants you to rank the forces, recognize hydrogen bonding, and reason from a property to a structure and back.

Intermolecular forces, weakest to strongest

• London dispersion forces exist between all molecules, caused by temporary, fluctuating charge imbalances. They are the only intermolecular force in nonpolar substances, and they get stronger as molecules get larger (more electrons), which is why larger nonpolar molecules have higher boiling points.
• Dipole-dipole forces act between polar molecules (those with a permanent positive and negative end, from molecular geometry and polarity). The positive end of one molecule attracts the negative end of another, an attraction stronger than dispersion alone.
• Hydrogen bonding is an especially strong dipole-dipole force that occurs when hydrogen is bonded to a small, highly electronegative atom (nitrogen, oxygen, or fluorine). It is responsible for water's unusually high boiling point and many of its special properties.

How forces set melting and boiling points

To melt or boil a substance you must give the particles enough energy to overcome the forces holding them together. So the stronger the forces, the higher the melting and boiling points and the less volatile the substance. This is the key reasoning the standard wants:

• Small nonpolar molecules (only dispersion forces) have very low melting and boiling points (oxygen, methane).
• Polar molecules (dipole-dipole) are higher.
• Hydrogen-bonded substances (water, ammonia) are higher still for their size.
• Ionic solids and covalent networks are far higher again, because melting them breaks strong bonds, not weak intermolecular forces.

Comparing the four structure types

Bringing Module 2 together, four structure types have characteristic properties:

• Ionic (giant lattice of ions): high melting point, hard but brittle, conducts only when molten or dissolved.
• Covalent molecular (small molecules): low melting point, soft, does not conduct, held by weak intermolecular forces.
• Covalent network (giant covalent, like diamond or quartz): very high melting point, very hard, usually does not conduct, because every atom is locked by strong covalent bonds.
• Metallic (cations in an electron sea, see metallic bonding and material properties): conducts, malleable, mostly high melting point.

Given an unknown's properties, you can infer which type it is, which is exactly the inference HS-PS1-3 asks for.

Solubility and "like dissolves like"

A substance dissolves best in a solvent with similar forces. Polar and ionic substances dissolve in polar solvents such as water, because water's polar molecules can surround and attract the solute particles. Nonpolar substances (oils, fats) dissolve in nonpolar solvents but not in water, because water molecules attract each other more strongly than they attract a nonpolar molecule. This is why oil and water do not mix.

Try this

Q1. Which has the higher boiling point, a small nonpolar molecule or a similarly sized hydrogen-bonded molecule? Explain. [2]

• Cue. The hydrogen-bonded molecule; its stronger intermolecular forces need more energy to overcome, so it boils at a higher temperature.

Q2. Will a nonpolar substance dissolve well in water? Explain. [2]

• Cue. No; water is polar and its molecules attract each other more strongly than they attract a nonpolar molecule, so the nonpolar substance does not dissolve (like dissolves like).