Chapter 9: Solutions

Before you can understand colligative properties or solubility rules, you need the vocabulary. A solution is a homogeneous mixture of two or more substances - meaning the composition is uniform throughout. Unlike a salad (heterogeneous), a solution looks the same no matter which part you sample.

Solute and Solvent

Every solution has two components:

• Solvent - the substance present in the greater amount (the “dissolver”). It determines the phase of the solution.
• Solute - the substance present in the lesser amount (the “dissolved stuff”).

When water is the solvent, the solution is called an aqueous solution - by far the most common type on the MCAT. But solutions are not limited to liquids dissolving solids.

Types of Solutions

Solutions can exist in any combination of phases:

Solid-in-solid solutions are called alloys. Gas-in-liquid solutions are critical for biology - dissolved O₂ and CO₂ in blood are the reason you can breathe.

Concentration - How Much is Dissolved?

Concentration describes the amount of solute per amount of solution (or solvent). You already learned the main concentration units in Chapter 4 (molarity, molality, mole fraction), but here is a quick review since they show up constantly in solutions problems:

Saturation

When you dissolve sugar in water, at some point the water cannot hold any more. That limit is defined by three levels:

• Unsaturated - less solute is dissolved than the maximum. More can dissolve.
• Saturated - the maximum amount of solute is dissolved at that temperature. If you add more, it will not dissolve - it just sits at the bottom as undissolved solid (precipitate). A dynamic equilibrium exists between dissolving and precipitating.
• Supersaturated - MORE solute is dissolved than the maximum. This is an unstable, metastable state. It is created by dissolving solute at a high temperature (where solubility is greater), then slowly cooling without disturbing. The tiniest disturbance - a scratch, a seed crystal, even a vibration - causes the excess solute to crash out of solution dramatically.

Electrolytes vs. Nonelectrolytes

This distinction matters enormously for colligative properties:

• Strong electrolytes dissociate completely into ions in water. They conduct electricity well. Examples: NaCl, HCl, KOH, CaCl₂.
• Weak electrolytes dissociate partially. They conduct electricity weakly. Examples: CH₃COOH (acetic acid), NH₃, HF.
• Nonelectrolytes do not dissociate at all. They dissolve as intact molecules. No ions, no conductivity. Examples: glucose (C₆H₁₂O₆), sucrose, urea.

Why does this matter? Because NaCl dissolves into 2 particles (Na⁺ + Cl⁻), but glucose dissolves into 1 particle. Two particles have twice the effect on colligative properties. This is exactly what the van ‘t Hoff factor (i) captures - but we will get to that in Section 11.

The most important rule in solubility is also the simplest: like dissolves like. Polar solvents dissolve polar solutes. Nonpolar solvents dissolve nonpolar solutes. Trying to dissolve a nonpolar substance in a polar solvent (or vice versa) is like trying to mix oil and water - it does not work.

Why Does “Like Dissolves Like” Work?

Dissolution happens when the solute-solvent interactions are strong enough to compensate for breaking apart solute-solute and solvent-solvent interactions. Think of it as a three-step process:

1. Break apart the solute - separate solute molecules from each other (costs energy, endothermic)
2. Break apart the solvent - make room in the solvent for the solute (costs energy, endothermic)
3. Form solute-solvent interactions - the new attractions between solute and solvent (releases energy, exothermic)

If the energy released in step 3 roughly matches or exceeds the energy costs of steps 1 and 2, the substance dissolves. If step 3 is much weaker than steps 1 + 2, the substance is insoluble.

Polar dissolves polar because polar molecules form strong dipole-dipole and hydrogen bonding interactions with each other. When a polar solute enters a polar solvent, the new solute-solvent interactions (step 3) are comparable to the old solute-solute and solvent-solvent interactions.

Nonpolar fails in polar solvent because nonpolar solutes can only form weak London dispersion forces with polar solvent molecules. These weak new interactions (step 3) cannot compensate for the strong hydrogen bonds between water molecules that must be disrupted (step 2).

Water as a Solvent

Water is called the “universal solvent” because its high polarity and ability to hydrogen bond let it dissolve an enormous range of substances. Water dissolves ionic compounds by surrounding each ion with a hydration shell - water molecules orient their partial charges toward the ion (oxygen faces cations, hydrogen faces anions).

This process is called hydration (or solvation if the solvent is not water). The ion-dipole interactions in the hydration shell are strong enough to overcome the ionic bonds in the crystal lattice - which is why NaCl dissolves in water.

Key Vocabulary

Ion-Dipole Interactions - The Force That Dissolves Salts

When ionic compounds dissolve in water, the driving force is ion-dipole interactions - the attraction between an ion’s full charge and water’s partial charges. This is the strongest intermolecular force (stronger than hydrogen bonding, dipole-dipole, or London dispersion).

The strength of ion-dipole interactions depends on:

• Charge density of the ion - smaller ions with higher charges have stronger interactions (Li⁺ > Na⁺ > K⁺)
• Polarity of the solvent - more polar solvents interact more strongly

This is why ionic compounds dissolve well in water (very polar) but not in hexane (nonpolar). The ion-dipole forces with water compensate for the lattice energy holding the crystal together.

We know that “like dissolves like,” but why do some polar substances dissolve easily while others barely dissolve at all? The answer is thermodynamics - specifically, the balance between the energy it takes to break apart the solute and the energy released when the solvent surrounds the solute particles.

The Three-Step Energy Model

Dissolving an ionic compound in water involves three energy changes:

Step 1: Break the crystal lattice (endothermic)

The lattice energy (U) is the energy required to completely separate all ions in a solid ionic compound into gaseous ions. This is always endothermic - you must overcome the electrostatic attractions holding the crystal together.

Lattice energy depends on:

• Charge - higher charges mean stronger attraction (MgO has much higher lattice energy than NaCl)
• Size - smaller ions pack closer together, increasing attraction (LiF > NaF > KF)

Step 2: Break solvent-solvent interactions (endothermic)

Water molecules must separate to make room for the incoming ions. This requires breaking some hydrogen bonds between water molecules.

Step 3: Form ion-dipole interactions (exothermic)

Water molecules surround each ion, forming a hydration shell. The energy released in this step is called the hydration energy (Δ$H_{\text{hydration}}$). This is always exothermic - forming new attractions releases energy.

Hydration energy depends on charge density:

• Higher charge = stronger ion-dipole interactions = more energy released
• Smaller radius = higher charge density = more energy released

Enthalpy of Solution

But Wait - Endothermic Dissolution Still Happens?

If Δ$H_{\text{soln}}$ is positive (endothermic), the process is energetically uphill. Why would it happen at all?

Because entropy also drives dissolution. Dissolving a solid increases disorder - the organized crystal lattice breaks into randomly dispersed ions. The entropy increase (ΔS > 0) can make the overall free energy change favorable:

ΔG = ΔH - TΔS

Even if ΔH is positive (endothermic), a large positive TΔS can make ΔG negative (spontaneous). This is why NaCl dissolves in water even though the process absorbs a tiny amount of heat - the entropy increase from dispersing Na⁺ and Cl⁻ ions throughout the solution more than compensates.

Hydration Shells

When ions dissolve, water molecules arrange themselves around each ion with specific orientations:

• Around cations (like Na⁺): the oxygen end of water (partial negative) faces the positive ion
• Around anions (like Cl⁻): the hydrogen end of water (partial positive) faces the negative ion

The first layer of water molecules is tightly organized - this is called the primary hydration shell. Beyond it, water molecules become progressively less ordered. Smaller, more highly charged ions (like Li⁺ or Mg²⁺) have larger, more tightly bound hydration shells than larger, less charged ions (like K⁺ or Cs⁺).

How much of a substance dissolves depends on more than just “like dissolves like.” Two physical conditions dramatically affect solubility: temperature and pressure. Understanding these factors is essential for interpreting solubility curves and predicting when gases come out of solution - both high-yield MCAT topics.

Effect of Temperature on Solid Solubility

For most solid solutes, solubility increases as temperature increases. Heating provides energy to break apart the crystal lattice, allowing more solute to dissolve.

But there are exceptions. A few compounds, like Ce₂(SO₄)₃ and CaSO₄, have solubility that decreases with increasing temperature. Their dissolution is exothermic, so by Le Chatelier’s principle, adding heat (increasing temperature) shifts the equilibrium toward the undissolved solid.

The key to remembering which direction:

• Endothermic dissolution (most solids): heat is a “reactant,” so adding heat shifts equilibrium toward more dissolving. Solubility increases with T.
• Exothermic dissolution (some solids): heat is a “product,” so adding heat shifts equilibrium toward less dissolving. Solubility decreases with T.

Solubility Curves

A solubility curve plots the maximum amount of solute (in grams per 100 g water) that dissolves at each temperature. These graphs are MCAT favorites.

How to read them:

• Points ON the curve represent saturated solutions
• Points BELOW the curve represent unsaturated solutions (more could dissolve)
• Points ABOVE the curve represent supersaturated solutions (unstable - more dissolved than the limit)

Effect of Temperature on Gas Solubility

Gases are the opposite of solids: gas solubility decreases as temperature increases.

Think about it physically: gas molecules in solution are trapped by intermolecular forces with the solvent. Heating gives them more kinetic energy, making it easier to escape into the gas phase.

Effect of Pressure on Gas Solubility - Henry’s Law

Pressure has virtually no effect on the solubility of solids and liquids (they are nearly incompressible). But for gases, pressure has a dramatic, direct effect.

Double the pressure, double the amount of gas that dissolves. Halve the pressure, half the gas stays in solution - the rest escapes as bubbles.

Summary: What Affects Solubility?

Not all ionic compounds dissolve in water. The MCAT expects you to know which salts are soluble and which form precipitates. The good news: the rules are straightforward, and a few simple patterns cover nearly every compound you will encounter.

The Rules - Soluble Compounds

Always soluble (no exceptions):

Usually soluble (with exceptions):

The Rules - Insoluble Compounds

Usually insoluble (with exceptions):

How to Apply the Rules

When asked “Is this compound soluble?”:

1. Check the cation first. Is it Na⁺, K⁺, or NH₄⁺? If yes, it is soluble. Done.
2. Check the anion. Is it NO₃⁻ or CH₃COO⁻? If yes, it is soluble. Done.
3. Check for exceptions. Is it a halide with Ag⁺, Pb²⁺, or Hg₂²⁺? Then it is insoluble. Is it a sulfate with Ba²⁺ or Pb²⁺? Insoluble.
4. Default for carbonates, phosphates, hydroxides, sulfides: insoluble (unless paired with Group 1 or NH₄⁺).

Predicting Precipitation Reactions

When two aqueous solutions are mixed, a precipitate forms if any combination of the ions produces an insoluble compound:

Example: Mix AgNO₃(aq) and NaCl(aq)

• Possible products: AgCl and NaNO₃
• AgCl: Ag⁺ + Cl⁻ = insoluble (halide exception with Ag⁺)
• NaNO₃: Na⁺ + NO₃⁻ = soluble (Na⁺ is always soluble)
• Result: AgCl precipitates as a white solid

The net ionic equation shows only the ions that actually participate:

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Na⁺ and NO₃⁻ are spectator ions - they remain dissolved and do not participate in the reaction.

Here is the single most important idea in this chapter: colligative properties depend on how many particles are dissolved, not what those particles are. Dissolve 1 mole of glucose (1 particle per formula unit) or 1 mole of NaCl (2 particles per formula unit), and the colligative effects scale with particle count alone. Identity is irrelevant.

What Are Colligative Properties?

The word “colligative” comes from the Latin “colligare,” meaning “to bind together.” There are four colligative properties:

All four share the same underlying cause: dissolved particles interfere with the solvent’s ability to escape into the gas phase.

The Root Cause - Vapor Pressure Lowering

Every colligative property traces back to one phenomenon: a nonvolatile solute lowers the vapor pressure of the solvent.

Why? When solute particles dissolve, they occupy positions at the liquid surface. Fewer solvent molecules are at the surface, so fewer can escape into the gas phase per second. The rate of evaporation decreases, so the equilibrium vapor pressure drops.

This is purely a numbers game. Each solute particle blocks a surface site. More particles = more blocking = lower vapor pressure. It does not matter what the particles are - only how many are there.

How Vapor Pressure Lowering Creates the Other Three Properties

Once you understand that solute lowers vapor pressure, the other three properties follow logically:

Boiling point elevation: The liquid must reach a higher temperature before its vapor pressure equals atmospheric pressure (the definition of boiling). Lower vapor pressure at any given temperature means you need more heat to boil. Boiling point goes up.

Freezing point depression: At the normal freezing point, the solution’s vapor pressure is now lower than the solid’s vapor pressure. The solid melts to equalize. You must cool further to reach the new temperature where solid and liquid vapor pressures match. Freezing point goes down.

Osmotic pressure: When a solution and pure solvent are separated by a semipermeable membrane, solvent flows toward the solution (from high vapor pressure to low). The pressure needed to stop this flow is osmotic pressure. More particles = lower vapor pressure = greater driving force = higher osmotic pressure.

Nonvolatile vs. Volatile Solutes

Colligative property equations (ΔTb, ΔTf, π) assume the solute is nonvolatile - meaning the solute does not evaporate. Sugar, NaCl, and glucose are nonvolatile. They stay in the liquid phase and only the solvent contributes to vapor pressure.

If the solute IS volatile (like mixing ethanol and water), both components contribute to vapor pressure, and the simple colligative property formulas do not apply cleanly. The MCAT almost always uses nonvolatile solutes for these calculations.

Ranking Colligative Effects

To compare colligative effects between different solutions, count the total dissolved particles:

A 1 M CaCl₂ solution has a higher boiling point, lower freezing point, and higher osmotic pressure than a 1 M NaCl solution - because it produces more particles.

Raoult’s law is the quantitative version of “adding solute lowers vapor pressure.” It tells you exactly how much the vapor pressure drops, and the math is beautifully simple.

The Law

Since χ_solvent is always less than 1 (there is always some solute present), $P_{\text{solution}}$ is always less than P°_solvent. The vapor pressure drops.

The vapor pressure lowering (ΔP) is:

ΔP = P°_solvent - $P_{\text{solution}}$ = χ_solute × P°_solvent

This form is sometimes more convenient: the drop in vapor pressure equals the mole fraction of the solute times the pure solvent’s vapor pressure.

Worked Example

Problem: The vapor pressure of pure water at 25 C is 23.8 mmHg. What is the vapor pressure of a solution made by dissolving 0.50 mol of glucose in 2.0 mol of water?

Solution:

• χ_water = 2.0 / (2.0 + 0.50) = 2.0 / 2.5 = 0.80
• $P_{\text{solution}}$ = 0.80 × 23.8 mmHg = 19.0 mmHg
• Vapor pressure lowering: ΔP = 23.8 - 19.0 = 4.8 mmHg

Ideal vs. Non-Ideal Solutions

Raoult’s law describes ideal solutions perfectly - solutions where solute-solvent interactions are identical to solute-solute and solvent-solvent interactions. In an ideal solution, every molecule “feels” the same regardless of its neighbors.

Positive deviations from Raoult’s law: The actual vapor pressure is HIGHER than Raoult predicts. This happens when solute-solvent interactions are weaker than the original pure-component interactions. The molecules escape more easily than expected.

• Example: ethanol + hexane (breaking H-bonds in ethanol, replacing with weak LDF)
• Δ$H_{\text{mix}}$ > 0 (endothermic mixing)

Negative deviations from Raoult’s law: The actual vapor pressure is LOWER than Raoult predicts. This happens when solute-solvent interactions are stronger than the pure-component interactions. The molecules are held in more tightly.

• Example: acetone + chloroform (new H-bonding between them)
• Δ$H_{\text{mix}}$ < 0 (exothermic mixing)

Two Volatile Components

When both the solute and solvent are volatile (e.g., mixing two liquids that both evaporate), both contribute to the total vapor pressure:

$P_{\text{total}}$ = χ_A × P°_A + χ_B × P°_B

This is the extended form of Raoult’s law. The total vapor pressure is the sum of each component’s partial pressure.

Salt your pasta water and it boils at a slightly higher temperature. That tiny change demonstrates boiling point elevation - one of the four colligative properties and a straightforward MCAT calculation.

The Formula

Important details:

• ΔTb is always positive (boiling point goes UP)
• The new boiling point = normal boiling point + ΔTb
• Kb for water = 0.512 C/m (usually given on the MCAT)
• Use molality (m), not molarity (M)

Why Does This Happen?

At the normal boiling point of pure water (100 C), the vapor pressure equals 1 atm. But a solution has a lower vapor pressure at 100 C (thanks to Raoult’s law). So at 100 C, the solution’s vapor pressure is still below 1 atm - it is not boiling yet.

You need to heat it past 100 C to push the vapor pressure up to 1 atm. The extra temperature needed is ΔTb.

Worked Example

Problem: Calculate the boiling point of a solution containing 58.5 g of NaCl dissolved in 1.00 kg of water. (Kb for water = 0.512 C/m)

Solution:

1. Moles of NaCl: 58.5 g / 58.5 g/mol = 1.00 mol
2. Molality: 1.00 mol / 1.00 kg = 1.00 m
3. Van ‘t Hoff factor: NaCl → Na⁺ + Cl⁻, so i = 2
4. ΔTb = i × Kb × m = 2 × 0.512 × 1.00 = 1.024 C
5. New boiling point = 100.0 + 1.024 = 101.0 C

Phase Diagram View

On a phase diagram, adding solute shifts the liquid-gas boundary to the right (higher temperature). The liquid phase region expands because the solution remains liquid at temperatures where pure water would have started boiling.

This shift is directly visible on the classic phase diagram that shows both boiling point elevation and freezing point depression - a favorite MCAT figure.

Freezing point depression is the mirror image of boiling point elevation: adding solute makes it harder for the solvent to freeze. This is the reason cities dump salt on icy roads and why you add antifreeze to your car’s radiator.

The Formula

Important details:

• ΔTf is the amount the freezing point drops (always positive as a magnitude)
• New freezing point = normal freezing point minus ΔTf
• Kf for water = 1.86 C/m (usually given)
• Kf is larger than Kb (1.86 vs. 0.512), so freezing point changes are more dramatic than boiling point changes for the same concentration

Why Does This Happen?

Freezing requires solvent molecules to arrange into an ordered crystal lattice. Dissolved solute particles disrupt this ordering process - they physically occupy positions in the liquid and prevent the regular crystal structure from forming.

At the normal freezing point (0 C for water), a solution is still liquid because the solute particles keep interfering with crystallization. You must cool further to provide enough driving force for the crystal to form despite the solute interference.

Why Is Kf Larger Than Kb?

This is a subtle but testable point. Kf (1.86 C/m) is about 3.6 times larger than Kb (0.512 C/m) for water. The freezing point is more sensitive to dissolved particles than the boiling point.

The reason: freezing involves ordering molecules into a crystal (huge entropy decrease), so even small amounts of solute have a big disrupting effect. Boiling is less structurally dependent.

Worked Example

Problem: What is the freezing point of a solution containing 0.50 mol of CaCl₂ in 500 g of water? (Kf = 1.86 C/m)

Solution:

1. Molality: 0.50 mol / 0.500 kg = 1.00 m
2. CaCl₂ → Ca²⁺ + 2 Cl⁻, so i = 3
3. ΔTf = i × Kf × m = 3 × 1.86 × 1.00 = 5.58 C
4. New freezing point = 0.0 - 5.58 = -5.58 C

Real-World Applications

Osmotic pressure is where solution chemistry meets biology. Every cell in your body is surrounded by a semipermeable membrane, and osmotic pressure determines whether that cell swells, shrinks, or stays the same size. This is one of the highest-yield topics in the entire solutions chapter.

Osmosis - The Driving Force

Osmosis is the net movement of solvent (usually water) through a semipermeable membrane from a region of lower solute concentration to higher solute concentration.

The membrane allows solvent molecules to pass but blocks solute particles. Solvent flows toward the more concentrated solution to equalize concentrations on both sides.

The osmotic pressure (π) is the minimum pressure that must be applied to the more concentrated side to prevent this solvent flow.

The Formula

Tonicity - The Biological Connection

This is where chemistry meets biology. The concepts below are explored in depth in the cell transport section - master them here, and you have already learned the biology.

When comparing two solutions separated by a membrane, we use three terms:

Why Osmotic Pressure Matters in Biology

Osmotic pressure explains:

• IV fluids must be isotonic (0.9% NaCl) - hypertonic IV would shrink red blood cells, hypotonic IV would lyse them
• Kidney function - the countercurrent multiplier creates an osmotic gradient in the medulla to concentrate urine
• Plant turgor pressure - water entering plant cells by osmosis creates internal pressure that keeps the plant rigid
• Edema - low blood protein (albumin) reduces osmotic pressure in capillaries, causing fluid to leak into tissues

Worked Example

Problem: What is the osmotic pressure of a 0.10 M NaCl solution at 37 C (body temperature)?

Solution:

1. i = 2 (NaCl → Na⁺ + Cl⁻)
2. M = 0.10 mol/L
3. R = 0.0821 L·atm/mol·K
4. T = 37 + 273 = 310 K
5. π = iMRT = 2 × 0.10 × 0.0821 × 310 = 5.09 atm

That is a substantial pressure - about 5 atmospheres from a relatively dilute solution. Osmotic pressure is highly sensitive to particle concentration.

Every colligative property formula has the factor “i” in it - the van ‘t Hoff factor. This single number is what makes the difference between a 1 m NaCl solution (i = 2, producing 2 mol particles) and a 1 m glucose solution (i = 1, producing 1 mol particles). Getting “i” right is essential for any colligative property calculation.

Determining i

The van ‘t Hoff factor equals the number of particles produced per formula unit of solute when dissolved:

Weak Electrolytes - Partial Dissociation

Weak electrolytes like acetic acid (CH₃COOH) partially dissociate in water. If the degree of dissociation is α (where 0 < α < 1):

i = 1 + α(n - 1)

where n = the number of ions the compound would produce if it fully dissociated.

For acetic acid (n = 2) with α = 0.05 (5% dissociation):

i = 1 + 0.05(2 - 1) = 1 + 0.05 = 1.05

The van ‘t Hoff factor is just barely above 1, because very few molecules actually dissociate.

Expected vs. Measured i Values

Here is one of the most commonly tested subtleties in this chapter:

Notice that MgSO₄ has an especially low measured i (1.21 vs. expected 2.00). This is because Mg²⁺ and SO₄²⁻ are both doubly charged - the electrostatic attraction between them is very strong, leading to extensive ion pairing.

Ion Pairing and Concentration

Ion pairing increases at higher concentrations because ions are closer together and more likely to encounter each other. This means:

• At low concentrations: measured i is close to expected i
• At high concentrations: measured i is significantly below expected i

Putting It All Together

To solve any colligative property problem:

1. Identify the solute - strong electrolyte, weak electrolyte, or nonelectrolyte?
2. Determine i - count the ions (strong electrolyte), use 1 (nonelectrolyte), or calculate from α (weak electrolyte)
3. Plug into the formula - ΔTb = iKbm, ΔTf = iKfm, or π = iMRT
4. Check for reasonableness - more particles always means a bigger effect

Not every mixture is a true solution. Some mixtures have particles too large to dissolve at the molecular level but too small to settle out immediately. These intermediate mixtures - colloids and suspensions - have unique properties that show up on the MCAT.

The Three Types of Mixtures

The Tyndall Effect

The Tyndall effect is the scattering of light by colloidal particles. Shine a flashlight through a true solution and the beam is invisible (particles too small to scatter light). Shine it through a colloid and the beam becomes visible as a glowing path through the mixture.

This is why:

• Car headlights are visible in fog (fog is a colloid of water droplets in air)
• A flashlight beam is visible in a dusty room (dust is a colloid of solid particles in air)
• Saltwater appears clear (true solution, no scattering)
• Milk appears white (colloid, scatters all wavelengths)

Brownian Motion

Brownian motion is the random, zigzag movement of colloidal particles suspended in a fluid. It is caused by constant bombardment from surrounding solvent molecules. The particles are small enough to be jostled by individual molecular collisions but large enough to be observed under a microscope.

Brownian motion:

• Is evidence for the kinetic molecular theory (molecules are in constant random motion)
• Keeps colloid particles suspended (prevents settling)
• Increases with temperature (faster-moving solvent molecules hit harder)

Types of Colloids

Colloids are classified by the phases of the dispersed substance and the medium:

Emulsifying Agents

Oil and water form a suspension (they separate on standing). But add an emulsifier - a molecule with both hydrophilic and hydrophobic parts - and the oil is dispersed into tiny droplets that stay suspended as a colloid (an emulsion).

Examples of emulsifiers:

• Soap and detergent - clean grease by emulsifying it in water
• Bile salts - emulsify dietary fats in the small intestine so lipase can digest them
• Lecithin (in egg yolk) - why mayonnaise stays mixed (oil emulsified in vinegar)

The emulsifier coats the oil droplets, with its polar head facing water and nonpolar tail embedded in the oil. This prevents the droplets from merging back together.

Coagulation and Flocculation

Colloidal particles often carry surface charges that create electrostatic repulsion, keeping them dispersed. Adding electrolytes (like salts) neutralizes these charges and causes the particles to clump together and settle out. This process is called coagulation or flocculation.

This is why:

• Adding alum to murky water causes impurities to clump and settle (water treatment)
• River water clears when it meets salty ocean water (salt neutralizes colloidal clay charges)