This unit is about what happens when one substance dissolves in another, how much of it can dissolve, and how dissolved particles change a liquid's behavior. By the end, you should be able to predict solubility trends, calculate concentration in several units, and use molality to find new boiling and freezing points.
Read the overview for the big picture, then work through each topic note and try its checkpoint question before you look at the answer.
Solutions are everywhere in chemistry, from the HCl in a titration flask to the salt on an icy road. This unit gives you the vocabulary and the math to describe them. Most of the calculations build directly on the mole conversions you already practiced in stoichiometry.
Honors Chemistry Unit 7 covers solutions and solubility. You learn what solutes and solvents are, why polar substances tend to dissolve in polar solvents, and how temperature and pressure change solubility for solids and gases. You then calculate concentration as molarity (mol/L of solution), molality (mol/kg of solvent), mass percent, and parts per million, and you use for dilutions. The unit ends with colligative properties, where you use to predict how much a solute raises a boiling point or lowers a freezing point.
Polar and many ionic solutes tend to dissolve in polar solvents like water, while nonpolar solutes tend to dissolve in nonpolar solvents like hexane. Salt dissolves in water, while oil forms a separate layer because it is nonpolar. Most oils float because they are less dense than water. Temperature can change solubility, and pressure has its largest effect on gases.
Molarity divides moles of solute by liters of solution. Molality divides moles of solute by kilograms of solvent. Mass percent divides grams of solute by grams of solution. Picking the right denominator is most of the work.
Adding a nonvolatile solute lowers vapor pressure, raises the boiling point, lowers the freezing point, and produces osmotic pressure. For solutions with the same molality, these changes depend mainly on the number of dissolved particles. Under the ideal assumption of complete dissociation, NaCl has about twice the effect of sugar because NaCl produces two ions per formula unit while sugar remains as one molecule.
Solute-solvent attractions help predict whether a substance dissolves, while solubility sets the maximum amount that can dissolve under given conditions. Concentration units measure the amount that actually dissolves. For colligative properties, the number of dissolved particles determines the size of the boiling point, freezing point, vapor pressure, and osmotic effects.
A solution is a homogeneous mixture, meaning it looks the same all the way through, even under a microscope.
Molarity is moles of solute per liter of solution, M = \frac{n}{V}.
Colligative properties depend mainly on how many solute particles are dissolved, not on their chemical identity.
A solution is a homogeneous mixture, meaning it looks the same all the way through, even under a microscope. The solute is what gets dissolved and the solvent does the dissolving. Solutions can exist in any phase. Air is a gas dissolved in a gas, soda is CO₂ gas dissolved in water, and brass is zinc dissolved in solid copper.
Whether something dissolves often depends on polarity. "Like dissolves like" is a useful guideline, so polar substances tend to dissolve in polar solvents and nonpolar substances tend to dissolve in nonpolar solvents. Many ionic compounds dissolve in water, but some are only slightly soluble or insoluble because their ionic lattices are too difficult to separate. When NaCl dissolves, water's partially negative oxygen end points toward Na⁺ and its partially positive hydrogen end points toward Cl⁻, forming a hydration shell around each ion.
How much dissolves depends on temperature and pressure. Most solids get more soluble as temperature rises, but gases get less soluble, which is why warm soda goes flat faster. Pressure mostly matters for gases. At constant temperature, Henry's law says gas solubility rises with the partial pressure of that gas above the liquid, so a sealed soda holds more CO₂ than an open one.
Keep "how much" separate from "how fast." Stirring and grinding a solid make it dissolve faster, but they don't change its solubility.
Check yourself. The solubility of KNO₃ at 20 °C is about 32 g per 100 g of water. You have 25 g of KNO₃ fully dissolved in 100 g of water at 20 °C. Is the solution unsaturated, saturated, or supersaturated? The solution is unsaturated because 25 g is less than the 32 g limit, so about 7 g more could still dissolve.
| Change | Effect on solid solutes in water | Effect on gas solutes in water |
|---|---|---|
| Raise temperature | Solubility usually increases | Solubility decreases |
| Raise the solute gas's partial pressure | Almost no effect | Solubility increases according to Henry's law |
| Stir | A solid dissolves faster, with the same solubility | Gas escapes faster from an open container |
| Grind a solid into smaller pieces | The solid dissolves faster, with the same solubility | Not applicable |
Molarity is moles of solute per liter of solution, . The volume is the whole solution, not just the water you started with.
Worked example. You dissolve 5.85 g of NaCl and add water until the solution reaches 250 mL. Moles of NaCl . Volume . Molarity .
Molality is moles of solute per kilogram of solvent. Mass does not change with temperature the way volume does, so molality is used for boiling point elevation and freezing point depression calculations.
To dilute, use . Adding water doesn't change the moles of solute, only the volume they're spread through. To make 500 mL of 0.50 M HCl from 2.0 M stock, solve . Put some water in a 500 mL volumetric flask, carefully add the 125 mL of HCl stock to that water, and swirl to mix. Let the solution cool to room temperature before adding water to the 500 mL mark. Acid goes into water, not the other way around, because mixing releases heat.
Percent shows up two ways. Mass percent of a solution is grams of solute over grams of solution, times 100, so 60 g glycerin in a 150 g lotion is 40%. Percent composition of a compound is the mass of one element over the molar mass, times 100. Hydrogen in water is .
Check yourself. How many grams of glucose (C₆H₁₂O₆, 180.2 g/mol) do you need to make 0.500 L of a 0.200 M solution? You need 18.0 g. Moles = 0.200 mol/L × 0.500 L = 0.100 mol, and 0.100 mol × 180.2 g/mol = 18.0 g.
| Unit | Numerator | Denominator | Typical use |
|---|---|---|---|
| Molarity (M) | Moles of solute | Liters of solution | Lab prep, titrations, stoichiometry |
| Molality (m) | Moles of solute | Kilograms of solvent | Boiling point and freezing point problems |
| Mass percent | Grams of solute | Grams of solution | Product labels and everyday mixtures (multiply the ratio by 100) |
| ppm | Grams of solute | Grams of solution | Pollutants and trace amounts (multiply the ratio by 10⁶) |
Colligative properties depend mainly on how many solute particles are dissolved, not on their chemical identity. A nonvolatile solute lowers the solvent's mole fraction, which lowers the equilibrium vapor pressure. Because boiling occurs when vapor pressure equals external pressure, the solution must reach a higher temperature to boil. Solute particles also make formation of the solvent's solid structure less favorable, so the solution freezes at a lower temperature. Osmotic pressure is the pressure needed to stop the net flow of solvent through a semipermeable membrane into the more concentrated solution.
The formulas are and . For water, °C/m and °C/m. In ideal calculations, sugar has , NaCl has , and CaCl₂ has . Real electrolyte solutions can have smaller effective values because ions interact.
Worked example. Assume NaCl dissociates completely. Dissolve 1.00 mol of NaCl in 2.00 kg of water. Molality . Then °C, so the ideal solution freezes at °C. Similarly, °C, so it boils at about 100.51 °C at 1 atm.
Check yourself. Under the ideal assumption of complete dissociation, which lowers water's freezing point more, 0.10 m CaCl₂ or 0.20 m glucose? CaCl₂ lowers it more. It produces 3 ions per formula unit, giving an ideal particle molality of 0.30 m, while glucose remains as molecules at 0.20 m.
| Term | Definition |
|---|---|
| Solution | A homogeneous mixture of two or more substances in a single phase, like salt water or air. |
| Solute | The substance that gets dissolved, usually present in the smaller amount. |
| Solvent | The substance that does the dissolving, usually present in the larger amount. |
| Solubility | The maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature and pressure. |
| Like dissolves like | A guideline stating that polar solutes tend to dissolve in polar solvents and nonpolar solutes tend to dissolve in nonpolar solvents. |
| Electrolyte | A solute that forms ions in water, so its solution conducts electricity. |
| Miscible | Describes two liquids that mix in any proportion, like ethanol and water. |
| Colligative property | A solution property that depends on the number of dissolved particles rather than their identity. |
| Percent composition | The percent by mass of each element in a compound, found by dividing the element's mass by the molar mass and multiplying by 100. |
Molarity divides by liters of the final solution. If a problem says you dissolved solute and diluted to 250 mL, use 0.250 L, even if the amount of water you poured in was different. For molality, it's the opposite, so use only the kilograms of solvent.
In ideal calculations, NaCl separates into Na⁺ and Cl⁻, so use . Leaving out this factor makes the predicted freezing point change half as large. Real electrolyte solutions can have smaller effective values because ions interact. Molecular solutes such as sugar and ethylene glycol usually use .
Heat helps most solids dissolve, but it pushes gases out of solution. Warm water holds less dissolved oxygen, which is why fish have trouble in warm ponds. Check whether the solute is a solid or a gas before you predict the trend.
Stirring and smaller particle size speed up dissolving by bringing fresh solvent into contact with more solute surface. They do not change equilibrium solubility. Temperature can change solubility, pressure strongly affects gas solubility, and changing the solvent or chemical conditions can also matter.