Solutions are mixtures in which a solute is evenly distributed in a solvent. This cheat sheet helps students connect concentration units, dilution calculations, and solubility ideas in one quick reference. It is especially useful for lab work, chemical equations, and multi-step stoichiometry problems.
Students need these tools to describe how much substance is dissolved and how solutions behave.
The most important idea is that concentration compares the amount of solute to the amount of solution. Common formulas include molarity, , dilution, , and mass percent, . Solubility depends on temperature, pressure, and the identities of the solute and solvent.
Electrolytes form ions in solution, so they affect conductivity and many reaction calculations.
Key Facts
- Molarity is calculated with , where is moles of solute and is volume of solution in liters.
- Moles of solute in a solution can be found with , which is useful before doing stoichiometry.
- Dilution problems use because the moles of solute stay constant when only solvent is added.
- Mass percent is calculated with .
- For dilute aqueous solutions, parts per million can be estimated with .
- A saturated solution contains the maximum amount of dissolved solute at a given temperature and pressure.
- For most solid solutes, solubility increases as temperature increases, but gas solubility usually decreases as temperature increases.
- Strong electrolytes dissociate almost completely into ions, so represents ion formation in water.
Vocabulary
- Solution
- A homogeneous mixture in which a solute is evenly distributed throughout a solvent.
- Solute
- The substance that is dissolved in a solution, such as in salt water.
- Solvent
- The substance that dissolves the solute and is usually present in the larger amount.
- Molarity
- A concentration unit equal to moles of solute per liter of solution, written as .
- Dilution
- The process of lowering concentration by adding solvent while keeping the moles of solute constant.
- Electrolyte
- A substance that forms ions in solution and allows the solution to conduct electricity.
Common Mistakes to Avoid
- Using milliliters directly in is wrong because molarity requires volume in liters, so must be converted to .
- Confusing solute mass with solution mass in mass percent is wrong because .
- Changing moles during dilution is wrong because adding solvent changes and , but the amount of solute stays the same.
- Treating all dissolved substances as strong electrolytes is wrong because molecular compounds such as sugar dissolve without forming many ions.
- Ignoring stoichiometric coefficients in solution reactions is wrong because gives moles of one substance, but the balanced equation gives the mole ratio.
Practice Questions
- 1 What is the molarity of a solution made by dissolving of in enough water to make of solution?
- 2 How many milliliters of are needed to prepare of ?
- 3 A solution contains of glucose dissolved in of water. What is the mass percent of glucose in the solution?
- 4 Explain why a solution of conducts electricity but a solution of sugar does not, even though both substances dissolve in water.
Understanding Solutions & Concentration
At the particle level, dissolving is a competition between attractions. Water molecules have slightly positive and slightly negative regions. They can surround many ionic compounds and pull their ions away from a crystal.
This process is called hydration. Sugar dissolves differently because whole sugar molecules spread between water molecules without forming ions. A substance can disappear from view yet still be present as particles throughout the liquid.
Stirring makes dissolving faster because it brings fresh solvent into contact with the solute. Crushing a solid increases surface area, which has a similar effect. Neither action necessarily increases the final amount that can dissolve.
Temperature changes the motion and energy of particles. A hot drink can often dissolve more sugar than a cold drink, but this pattern does not apply equally to every substance. Gases provide an important contrast.
Warm water holds less dissolved oxygen and carbon dioxide than cold water. This helps explain why a warm fizzy drink loses bubbles quickly. It also matters in lakes during hot weather, when lower oxygen levels can stress fish.
Pressure has little effect on most solids in liquids. It has a large effect on gases.
Carbon dioxide remains dissolved in a sealed soft drink because the gas above the liquid is under pressure. Opening the container lowers that pressure, so gas escapes.
Concentration calculations depend on careful measurement and unit choices. A volume in milliliters must be converted to liters before it is used in molarity work. The final volume means the volume after solute and solvent have been combined.
It is not always correct to add separate volumes from a recipe, especially for concentrated liquids. In a lab, a volumetric flask gives a much more reliable final volume than a beaker. A student adds some water, dissolves the measured solute completely, then fills to the calibration mark.
Reading the bottom of the curved liquid surface at eye level reduces volume error. Small errors can matter when a solution is later used to determine an unknown concentration.
Solution stoichiometry connects a measured volume to the chemical change in a reaction. First, use the solution concentration and volume to find the amount of dissolved substance in moles. Next, use the balanced reaction to compare mole amounts.
Finally, convert to the requested mass, volume, or concentration. Ionic equations require extra attention because dissolved strong electrolytes separate into ions. In a reaction between silver nitrate solution and sodium chloride solution, silver ions join chloride ions to form an insoluble solid.
Sodium ions and nitrate ions stay dissolved and are spectator ions. Counting the ions correctly is essential for conductivity, precipitation reactions, acid base reactions, and titrations. The most common mistakes are using the wrong volume units, forgetting a coefficient from the balanced equation, and treating a weak electrolyte as if it formed ions completely.