Predicting whether substances mix is a major application of intermolecular forces. “Like dissolves like” connects molecular structure to solubility by comparing which attractions are present and whether new solute–solvent attractions can replace old ones.

Core idea: matching intermolecular attractions

Dissolving requires disrupting some solute–solute attractions and some solvent–solvent attractions, then forming new solute–solvent attractions. Mixing is favoured when the new attractions are comparable in type and strength to those that were broken.

“Like” refers to similarity in intermolecular forces (IMFs) (and therefore similar charge distribution and polarizability), not simply similar chemical formulas.

The key IMF categories used for “like dissolves like”

• Nonpolar substances (roughly symmetric, no net dipole): primarily London dispersion forces (LDF).

• Polar molecular substances (net dipole): dipole–dipole attractions (plus LDF).

• Hydrogen-bonding substances (H bonded to N, O, or F and a lone-pair acceptor present): strong, directional hydrogen bonding (plus other IMFs).

• Ionic substances: strong ion–ion attractions in the solute; dissolution requires strong ion–dipole interactions with the solvent.

Soluble vs miscible: language you must use correctly

A liquid solute in a liquid solvent is described as miscible or immiscible; a solid (or gas) in a liquid is described using solubility (soluble, slightly soluble, insoluble).

Applying “like dissolves like” to common cases

Nonpolar solutes in nonpolar solvents

Nonpolar solutes (e.g., many hydrocarbons) tend to dissolve in nonpolar solvents because:

• Breaking LDF in the solute and solvent is “paid for” by forming new LDF of similar magnitude.

• Similar polarizability and surface contact enable effective dispersion attractions.

Polar solutes in polar solvents

Polar solutes tend to dissolve in polar solvents because:

• Dipole–dipole attractions and alignment between partial charges can form readily.

• If hydrogen bonding is possible, matching H-bond donors/acceptors increases compatibility.

Hydrogen-bonding compatibility

Hydrogen bonding is most effective when the solute and solvent can both participate.

• Molecules with O–H or N–H groups often dissolve well in solvents that can accept H-bonds (lone pairs).

• Molecules with only H-bond acceptors (e.g., a carbonyl oxygen) may dissolve better in solvents that can donate H-bonds.

Ionic solutes in polar solvents

Ionic compounds dissolve best in solvents with strong molecular dipoles (often water) because ion–dipole attractions can compensate for the energy needed to separate ions.

Ion–dipole attraction is shown by water molecules orienting their partial charges toward an ion: the oxygen end ($\delta^-$) points toward a cation like Na$^+$, while the hydrogen ends ($\delta^+$) point toward an anion. This directional alignment illustrates how hydration shells stabilize separated ions in solution, helping offset the energetic cost of disrupting an ionic lattice. Source

• If ion–dipole attractions are too weak, the crystal lattice remains intact and solubility is low.

What to look for when predicting solubility from structure

Compare the “interaction inventory”

Ask which attractive forces are available for each substance:

• Does the solute have a net dipole?

• Can it donate an H-bond (N–H, O–H, F–H)?

• Can it accept an H-bond (lone pairs on N, O, F)?

• Is it largely nonpolar (hydrocarbon-like) with only LDF?

• Does it form ions (ionic compound) requiring ion–dipole stabilization?

Balance of polar and nonpolar regions

Many molecules contain both polar and nonpolar parts. Solubility depends on which parts dominate the intermolecular behaviour:

• Increasing nonpolar surface area often decreases solubility in very polar solvents.

• Adding polar functional groups generally increases solubility in polar solvents by adding stronger solute–solvent attractions.

Why “like dissolves like” works (a thermodynamic lens)

Mixing is favoured when it leads to a negative Gibbs free energy change.

A $\Delta G$ vs. $T$ plot makes the equation $\Delta G = \Delta H - T\Delta S$ visual: the intercept corresponds to $\Delta H$ and the slope corresponds to $-\Delta S$. The temperature where the line crosses $\Delta G = 0$ marks the threshold between nonspontaneous and spontaneous behavior, clarifying why entropy can dominate at higher $T$ when $\Delta S>0$. Source

When IMFs “match,” $\Delta H$ is often small (or favourable), and dispersal tends to make $\Delta S$ positive, so dissolution is more likely.

Common pitfalls and boundaries of the rule

• All molecules have LDF; the question is whether stronger, more specific attractions (dipole–dipole, H-bonding, ion–dipole) are needed to dissolve.

• “Polar” is not all-or-nothing; small structural changes can shift the dominant interactions and change solubility noticeably.

• Some mixtures form separate layers because solute–solvent attractions are too weak compared with solute–solute and solvent–solvent attractions, producing immiscibility even though each pure liquid is held together well.