Lyophilic and Lyophobic Colloids

Every colloid disperses matter of size roughly 1 nm to 1000 nm through a continuous medium, as introduced in the classification of colloids by dispersed phase and dispersion medium. What separates lyophilic from lyophobic behavior is the sign of the Gibbs energy change on forming that dispersion from bulk solute and pure solvent. A lyophilic colloid — starch in water, proteins in water, rubber in a nonpolar oil — disperses because the particle surface interacts favorably with the solvent. Polar or charged groups on a macromolecule hydrogen-bond with water; long flexible chains gain conformational entropy as they unfurl into solution. The dispersion process resembles ordinary dissolution: given time, the two components mix spontaneously, with $\Delta G_{\text{disp}} < 0$. Because the system sits at (or near) a genuine free-energy minimum, it is thermodynamically stable. Removing solvent by evaporation and adding fresh solvent later reverses the process cleanly — the sol reforms. This reversibility is the operational fingerprint of a lyophilic system. A lyophobic colloid — gold, silver iodide, or arsenic sulfide dispersed in water — has essentially no favorable chemical interaction between particle surface and solvent. The interfacial (surface) tension between particle and medium is high, so creating more interface by subdividing a bulk solid into colloidal particles costs Gibbs energy: $\Delta G_{\text{disp}} > 0$. Such a dispersion never forms spontaneously; it must be made by mechanical grinding (dispersion methods) or by controlled precipitation (condensation methods, e.g., reducing AuCl₄⁻ slowly). Once made, the particles are thermodynamically driven to aggregate and lower total surface area, but they may persist for years because charged double layers or adsorbed ions create an energy barrier to close approach. This is kinetic, not thermodynamic, stability — remove the barrier (add salt) and the sol coagulates irreversibly. The connection to wetting is direct: a lyophilic particle surface is wet by the solvent with a contact angle near 0°, while a lyophobic particle surface, if it were flat, would show a large contact angle with poor wetting. Colloidal affinity is the microscopic analogue of macroscopic wettability.

where: lyophilic sols have $\Delta G_{\text{disp}} < 0$ (spontaneous, reversible); lyophobic sols have $\Delta G_{\text{disp}} > 0$ (particles held apart only kinetically).

where: $\gamma_{\text{SL}}$ is the solid–liquid interfacial tension and $\Delta A$ is the new interfacial area created on subdividing bulk solid into colloidal particles.

where: $r$ is interparticle separation; the DLVO-type balance of electrostatic repulsion and van der Waals attraction sets the height of the kinetic barrier for lyophobic sols (developed further under colloid stability and coagulation).

where: $c_{\text{ccc}}$ is the critical coagulation concentration of electrolyte and $z$ is the charge number of the counter-ion (Schulze–Hardy rule), applicable specifically to lyophobic sols.

Example 1 — foundational: comparing viscosity signatures. A 1.0% (by mass) aqueous solution of gelatin (lyophilic) has a viscosity of $\eta = 1.5\times10^{-3}\ \text{Pa s}$ at 298 K, compared with pure water's $\eta_0 = 0.890\times10^{-3}\ \text{Pa s}$. A 1.0% aqueous gold sol (lyophobic, particles 10–20 nm) shows $\eta = 0.895\times10^{-3}\ \text{Pa s}$, barely different from water. Relative viscosity increase:

Sanity check: gelatin chains are long, solvated, and entangled, so they raise viscosity strongly even at low mass fraction — a hallmark of lyophilic character. Rigid, compact, dilute gold particles barely perturb bulk flow, consistent with lyophobic dispersion where solvent and particle interact weakly. This 100-fold difference in fractional viscosity rise is a standard qualitative test distinguishing the two colloid types in the laboratory.

Example 2 — applied: predicting coagulation from the Schulze–Hardy rule. A negatively charged lyophobic As₂S₃ sol coagulates at a critical concentration of $c_1 = 51\ \text{mmol L}^{-1}$ for NaCl (counter-ion Na⁺, $z_1=1$). Estimate the critical coagulation concentration for MgCl₂ ($z_2 = 2$) and for AlCl₃ ($z_3 = 3$), using $c_{\text{ccc}} \propto z^{-6}$. Step 1 — set up the ratio using the $z^{-6}$ scaling:

Step 2 — repeat for Al³⁺:

Sanity check: these predicted values (0.80 and 0.070 mmol L⁻¹) fall within the correct order of magnitude of classic Schulze–Hardy measurements on As₂S₃ sols (historically reported ratios near 1 : 700 : 500,000 for the ccc of mono-, di-, and trivalent cations against negative sols, though real ratios deviate from an exact $z^{-6}$ law because of ion specificity). The qualitative conclusion — trivalent counter-ions coagulate a lyophobic sol at roughly two orders of magnitude lower concentration than monovalent ones — is the physically important, exam-relevant result. A comparable lyophilic sol (gelatin) would show no sharp coagulation at all across this concentration range; it would simply "salt out" at much higher, roughly additive electrolyte concentrations, reflecting its different, hydration-driven stabilization mechanism.

Michael Faraday's 1857 preparation of ruby-red gold sols by reducing chloroauric acid, and his careful observation that these sols scattered light strongly (a phenomenon later named for John Tyndall) yet remained suspended for months, gave the first rigorous experimental foothold for what would become the lyophobic/lyophilic classification. Faraday recognized that these tiny gold particles were not truly dissolved in the chemical sense, yet persisted far longer than gravity alone would allow — a puzzle that later stability theory would resolve through the electrical double layer and kinetic barriers described above.

Ask the tutor: How does the DLVO balance of van der Waals attraction and double-layer repulsion quantitatively set the coagulation threshold for a lyophobic sol?