Polar solvents dissolve substances primarily because of the attraction of the opposite charges on the solvent and solute particles. This fundamental interaction, often summarized by the principle "like dissolves like," explains why polar solvents are effective at dissolving polar and ionic compounds.
Understanding Polarity
Molecules are considered polar when they have an uneven distribution of electron density, creating a positive end (partial positive charge) and a negative end (partial negative charge), forming a dipole. Water (H₂O) is a classic example of a polar solvent. The oxygen atom is more electronegative than the hydrogen atoms, pulling the shared electrons closer to itself and resulting in a partial negative charge on the oxygen and partial positive charges on the hydrogens.
Non-polar solvents, like hexane (C₆H₁₄), have an even distribution of electron density and no significant net dipole moment.
The Dissolution Process
When a polar solvent comes into contact with a polar or ionic solute, the charged ends of the solvent molecules are attracted to the oppositely charged parts of the solute.
- For Ionic Solutes (like salt, NaCl): The positive ends of the polar solvent molecules are attracted to the negatively charged ions (like Cl⁻), while the negative ends of the solvent molecules are attracted to the positively charged ions (like Na⁺). These attractions overcome the forces holding the ionic solute together, causing the ions to separate and become surrounded by solvent molecules. This process is called solvation (or hydration when the solvent is water).
- For Polar Solutes (like sugar, C₁₂H₂₂O₁₁): Polar solute molecules also have partial positive and negative charges. The partial charges on the solvent molecules are attracted to the opposite partial charges on the solute molecules. These dipole-dipole attractions pull the solute molecules apart and disperse them within the solvent.
The reference explicitly states: "Polar solvents will dissolve polar and ionic solutes because of the attraction of the opposite charges on the solvent and solute particles." This attraction is the driving force behind the dissolution process for polar solvents.
Why "Like Dissolves Like"?
This principle highlights the importance of similar intermolecular forces between the solvent and solute:
| Solvent Type | Solute Type | Interaction Forces involved | Dissolution? |
|---|---|---|---|
| Polar | Polar | Dipole-dipole attractions, Hydrogen bonding (if applicable) | Yes |
| Polar | Ionic | Ion-dipole attractions | Yes |
| Polar | Non-polar | Weak London dispersion forces | No (generally) |
| Non-polar | Non-polar | London dispersion forces | Yes |
| Non-polar | Polar or Ionic | Weak interactions, cannot overcome strong solute forces | No |
As the reference notes, "Non-polar solvents will only dissolve non-polar solutes because they cannot attract the dipoles or the ions" of polar or ionic substances.
Practical Examples
- Salt dissolving in water: Water (polar solvent) attracts the positive sodium ions and negative chloride ions of salt (ionic solute), pulling them apart.
- Sugar dissolving in water: Water (polar solvent) forms hydrogen bonds and dipole-dipole attractions with sugar (polar solute), breaking its crystal structure.
- Oil and water not mixing: Oil is non-polar, while water is polar. The strong attractions between water molecules are much greater than the weak attractions between water and oil molecules, so they separate.
In essence, polar solvents dissolve polar and ionic substances by forming favorable electrostatic attractions that overcome the forces holding the solute particles together, dispersing them throughout the solution.
