13.1 Solution Formation and Solubility
PRE-HEALTH PREP
Like Dissolves Like
The Science of Solubility
General Chemistry 13.1
Lesson Overview
Solubility describes how and why substances dissolve, governed by the principle that "like dissolves like"—polar solutes dissolve in polar solvents and nonpolar solutes dissolve in nonpolar solvents. The 3-step energetics of solution formation, the role of entropy, and the distinctions among saturated, unsaturated, and supersaturated solutions establish the thermodynamic framework for predicting dissolving behavior. Additional topics include miscible and immiscible liquids, colloids and the Tyndall effect, and the factors—temperature and pressure—that govern the solubility of gases and solids, including Henry's Law.
The 3-Step Process of Solution Formation
Forming a solution is not as simple as tossing a solute into a solvent and hoping for the best. At the molecular level, solution formation follows a 3-step process that determines whether a given solute will actually dissolve in a given solvent.
- Separate the solvent molecules. The intermolecular forces holding solvent molecules together must be overcome so that gaps open up for incoming solute particles.
- Separate the solute particles. The intermolecular forces within the solute must also be broken so that individual solute particles can disperse.
- Mix solute and solvent together. The separated solute particles slip into the gaps between solvent molecules, forming new intermolecular forces in the process.
Energetics of Each Step
Steps 1 and 2 both require energy to break apart intermolecular forces, making them endothermic (ΔH is positive). Step 3 releases energy as new intermolecular forces form between solute and solvent, making it exothermic (ΔH is negative).
The overall ΔH for the solution process is simply the sum of all three steps. In the best-case scenario, the new intermolecular forces formed in Step 3 are roughly equivalent in strength to the combined forces broken in Steps 1 and 2, giving an overall ΔH of approximately zero.
Why ΔH ≈ 0 Is the Best Case
A negative ΔH favors spontaneity, while a positive ΔH works against it. For dissolving, the best realistic outcome is ΔH near zero—the new forces formed rarely exceed the old ones broken. The key is to avoid a significantly positive ΔH, which would oppose the dissolving process.
| Step | What Happens | Energy Change |
|---|---|---|
| 1. Separate solvent | Break solvent–solvent intermolecular forces | Endothermic (+ΔH) |
| 2. Separate solute | Break solute–solute intermolecular forces | Endothermic (+ΔH) |
| 3. Mix together | Form new solute–solvent intermolecular forces | Exothermic (−ΔH) |
How Entropy Drives the Dissolving Process
If ΔH is zero at best, what actually makes dissolving happen? The answer lies in entropy—a measure of randomness or disorder in a system. An increase in entropy favors spontaneity, and mixing a solute into a solvent always increases randomness compared to having pure solute and pure solvent sitting separately.
Two factors govern whether dissolving is spontaneous:
- Enthalpy (ΔH): A negative value helps; a positive value hinders.
- Entropy of mixing: Mixing always increases randomness, which favors spontaneity.
In most dissolving processes, the entropy of mixing is the primary driving force. Starting with pure solvent and pure solute gives lower randomness than a well-mixed solution. This entropy increase nudges the process forward—as long as ΔH does not work too strongly against it.
The Bottom Line
For dissolving to occur spontaneously, ΔH must stay close to zero so that it does not oppose the favorable entropy of mixing. A large positive ΔH will overpower the entropy benefit, and the solute simply will not dissolve.
What Does "Like Dissolves Like" Mean?
For ΔH to remain near zero, the intermolecular forces formed between solute and solvent in Step 3 must be comparable in strength to the forces broken in Steps 1 and 2. This happens when the solute and solvent share similar polarity—the core principle of "like dissolves like" in chemistry.
In practical terms, the like dissolves like rule means:
- Polar solutes dissolve well in polar solvents (both have strong dipole–dipole forces or hydrogen bonding).
- Nonpolar solutes dissolve well in nonpolar solvents (both rely primarily on London dispersion forces).
- Trying to dissolve a polar solute in a nonpolar solvent—or vice versa—typically fails because the new forces formed cannot compensate for the old ones broken.
Example: Which Is More Soluble in Water?
Water is highly polar and capable of hydrogen bonding. Consider two potential solutes:
- Butane (CH3CH2CH2CH3): A nonpolar hydrocarbon with no electronegative atoms. It mixes poorly with water because polar and nonpolar do not match.
- Ethanol (CH3CH2OH): The O–H bond makes ethanol fairly polar and capable of hydrogen bonding. Ethanol dissolves readily in water.
Ethanol wins by a wide margin because its polarity closely matches water's—a textbook demonstration of like dissolves like.
Example: Which Is More Soluble in Cyclohexane?
Cyclohexane (C6H12) is a nonpolar hydrocarbon. Now the situation reverses:
- Butane dissolves readily in cyclohexane—both are nonpolar, so like dissolves like.
- Ethanol dissolves less well in cyclohexane because its polar O–H group does not match the nonpolar solvent.
Exam Strategy
When asked which solute is most soluble in a given solvent, first identify whether the solvent is polar or nonpolar. Then choose the solute whose polarity most closely matches. Polar solvent? Pick the most polar solute. Nonpolar solvent? Pick the least polar solute.
Miscible vs. Immiscible Liquids
When both the solute and solvent are liquids, chemists use special vocabulary to describe their solubility behavior. The key distinction—miscible vs. immiscible—follows directly from the like dissolves like principle.
| Term | Definition | Example |
|---|---|---|
| Miscible | Two liquids that mix in any proportion | Ethanol and water |
| Immiscible | Two liquids that largely do not mix | Oil and water |
Miscible liquids dissolve in each other completely regardless of ratio. Ethanol and water are a classic pair—you can combine them in any proportion and they will form a uniform solution every time.
Immiscible liquids refuse to mix to any significant extent. Oil and water demonstrate this perfectly: the oil settles into a separate layer on top with no appreciable mixing between the two. Oil is nonpolar; water is polar—exactly what the like dissolves like rule predicts.
Vocabulary Note
The terms "miscible" and "immiscible" are reserved exclusively for liquid–liquid combinations. You would not describe a solid dissolving in a liquid as "miscible."
3 Types of Solutions: Saturated, Unsaturated, and Supersaturated
For any given solute–solvent combination at a specific temperature, there is a maximum amount of solute that can dissolve. This limit defines the 3 types of solutions in chemistry.
| Solution Type | Solute Amount | Description |
|---|---|---|
| Unsaturated | Below the maximum | More solute could still dissolve |
| Saturated | At the maximum | No additional solute can dissolve |
| Supersaturated | Above the normal maximum | Unstable; contains more solute than normally possible |
Real-World Example: Making Rock Candy
Rock candy perfectly illustrates supersaturation. Here is how it works:
- Heat water to increase its ability to dissolve sugar.
- Dissolve sugar until the hot solution is saturated.
- Cool the solution very slowly and gently. Because sugar is more soluble at higher temperatures, the cooling solution now holds more sugar than it normally could at the lower temperature—it has become supersaturated.
- Drop in a string, stick, or seed crystal. The excess sugar rapidly crystallizes onto that surface, producing rock candy.
This rapid crystallization upon adding a seed crystal is the hallmark of a supersaturated solution's instability.
What Are Colloids?
Between fully soluble substances (which form true solutions) and insoluble substances (which settle out entirely) lies an intermediate category: colloids, also called colloidal suspensions or colloidal dispersions.
| Type | Particle Size | Behavior | Can Be Filtered? |
|---|---|---|---|
| True solution | Individual atoms or molecules | Completely homogeneous | No |
| Colloid | Clumps of molecules (larger particles) | Appears homogeneous; particles stay suspended | No (normal filtration) |
| Insoluble mixture | Large aggregates | Settles out over time | Yes |
In a colloid, particles are larger than individual molecules—they exist as clumps or globs—yet they remain permanently suspended rather than settling to the bottom. Unlike insoluble substances, colloids cannot be separated by ordinary filtration or centrifugation.
Example: Milk as a Colloid
Milk is a classic colloidal suspension. Fat molecules cluster together into globs that are larger than individual molecules, yet these globules remain evenly suspended throughout the liquid rather than settling to the bottom.
The Tyndall Effect: Identifying Colloids
Colloids scatter light in a way that true solutions do not. Shine a beam of light through a colloid and the beam becomes visible as it scatters off the suspended particles. This phenomenon, known as the Tyndall effect, provides a reliable way to distinguish a colloidal suspension from a true solution.
2 Factors That Affect Gas Solubility: Temperature and Pressure
The solubility of a gas in a liquid depends on two key factors: temperature and pressure. Understanding both is essential for general chemistry and is directly tested in exams.
- All gases are more soluble at lower temperatures.
- All gases are more soluble at higher pressures.
Temperature and Gas Solubility
As temperature increases, gas solubility decreases. This is true universally for gases—O2, CO2, N2, and every other gas all become less soluble as the solution warms up.
Everyday Example: Boiling Water for Spaghetti
When heating a pot of water for spaghetti, small bubbles appear long before the water reaches a rolling boil. These are not early signs of boiling—they are dissolved gases (primarily O2 and CO2) escaping the water as the temperature rises and their solubility decreases.
What Is Henry's Law?
The relationship between gas solubility and pressure is quantified by Henry's Law, one of the most important equations in general chemistry solubility:
S = kH × P
where:
- S = solubility of the gas (typically in mol/L)
- kH = Henry's Law constant (specific to each gas–solvent pair at a given temperature)
- P = partial pressure of the gas above the solution (in atm)
The relationship is directly proportional: double the partial pressure of a gas, and you double its solubility in the solution.
Henry's Law Constants Are Specific
Each Henry's Law constant applies to a specific gas in a specific solvent at a specific temperature. For example, the constant for O2 in water at 25°C is 1.3 × 10−3 M/atm. Change any one of those conditions and the constant changes too.
Worked Example: Solubility of O2 at Sea Level
At sea level, the partial pressure of O2 in the atmosphere is approximately 0.21 atm. Using the Henry's Law constant for O2 in water at 25°C:
S = (1.3 × 10−3 M/atm)(0.21 atm) = 2.7 × 10−4 M
This result makes sense: since we have considerably less than 1 atm of O2, the solubility should be less than the Henry's Law constant itself.
How Temperature Affects Solid Solubility
Unlike gases, the vast majority of solid solutes become more soluble at higher temperatures. As temperature increases, molecules gain kinetic energy, making it easier to break apart the solute–solute intermolecular forces and disperse solute particles throughout the solvent.
Example: Ionic and Molecular Solids
Ionic compounds like lead(II) nitrate, potassium dichromate, and potassium bromide all show increasing solubility with temperature—each displays a positive slope on a solubility-vs.-temperature graph. The same pattern holds for molecular solids like sugar, which is noticeably easier to dissolve in hot water than cold.
Exceptions Exist
While most solids follow this trend, there are rare exceptions where solubility decreases with increasing temperature. These cases are few and far between, but they do appear on exams. The general rule of thumb: assume most solids are more soluble at higher temperatures unless told otherwise.
| Solute Type | Effect of Increasing Temperature | Exceptions? |
|---|---|---|
| Gases | Solubility decreases | No (universal) |
| Most solids | Solubility increases | Rare exceptions |
Key Takeaways
- Solution formation follows 3 steps: separate solvent molecules, separate solute particles, then mix them together. Steps 1 and 2 are endothermic; Step 3 is exothermic.
- Entropy of mixing is the primary driving force for dissolving, as long as ΔH remains near zero.
- "Like dissolves like" means polar solutes dissolve in polar solvents and nonpolar solutes dissolve in nonpolar solvents, because matching polarity keeps ΔH close to zero.
- Miscible liquids mix in any proportion (e.g., ethanol and water); immiscible liquids do not mix (e.g., oil and water).
- The 3 types of solutions—saturated, unsaturated, and supersaturated—describe how much solute is dissolved relative to the maximum possible at a given temperature.
- Colloids are suspensions of larger-than-molecular particles that remain suspended and scatter light (Tyndall effect).
- Gas solubility increases with higher pressure (Henry's Law: S = kH × P) and decreases with higher temperature.
- Solid solubility generally increases with higher temperature, the opposite trend from gases.