The Nature of Water
Water’s role as the “universal solvent” is due to its molecular structure. The unequal sharing of electrons makes it a polar molecule, allowing it to dissolve a vast range of ionic and polar substances.
Water’s Polarity
The oxygen atom is more electronegative, pulling electrons closer and creating a partial negative charge. The hydrogen atoms have a partial positive charge. This separation of charge makes water a polar molecule.
Dissolving Ionic Compounds
The positive ends ($\delta+$) of water molecules attract anions, while the negative ends ($\delta-$) attract cations. These **ion-dipole attractions** pull the ionic lattice apart.
Dissolving Polar Molecules
Polar molecules like ethanol form **hydrogen bonds** with water molecules, allowing them to mix freely without breaking into ions.
Interactive Solubility Curves
Solubility often changes with temperature. This chart visualizes this relationship for several common salts. Hover over the lines to see specific data points. Notice how most solids become more soluble as temperature increases, while gases become less soluble.
Precipitation Reactions
When two soluble ionic solutions are mixed, they can form an insoluble solid called a precipitate. Use the VCE solubility rules (including the SNAPE rule) to predict the outcome below.
Precipitate Predictor
Select ions to see the result.
Remember SNAPE: Salts containing Sodium ($Na^+$), Nitrate ($NO_3^-$), Ammonium ($NH_4^+$), Potassium ($K^+$), or Ethanoate ($CH_3COO^-$) are always soluble.
Measuring Concentration
Concentration is the language of solutions. These interactive calculators will help you master the essential formulas for converting units and performing dilutions.
Concentration Unit Converter
Select a substance, enter a known concentration, and see the equivalent values instantly.
Dilution Calculator (C₁V₁ = C₂V₂)
Fill in three of the four fields to find the unknown value for diluting a stock solution.
Lab Essentials: Standard Solutions
The accuracy of many analytical techniques depends on a **standard solution**—a solution of precisely known concentration. They are prepared from a highly pure and stable **primary standard**.
Properties of a Primary Standard
- High Purity: Free from impurities that would affect its mass.
- Known Formula: Allows for accurate molar mass calculation.
- Stable: Not hygroscopic (absorb water) or reactive with air.
- High Molar Mass: Minimises weighing errors.
- Readily Soluble: Dissolves completely in the solvent.
Procedure for Preparation
- Calculate & Weigh: Accurately weigh the required mass of primary standard.
- Dissolve: Dissolve the solid completely in a small amount of deionised water in a beaker.
- Quantitative Transfer: Carefully transfer the solution and all rinsings into a volumetric flask.
- Dilute to the Mark: Add water until the bottom of the meniscus rests exactly on the calibration mark.
- Homogenise: Stopper and invert the flask repeatedly to ensure uniform concentration.
Gravimetric Analysis Workflow
This technique determines an analyte’s mass by converting it into a solid precipitate. This interactive diagram breaks the process into its key stages. Click through the steps to see how the analysis is performed.
Click a step to begin.
UV-Vis Spectroscopy
This technique is used for coloured solutions. It works on the principle that concentration is proportional to the amount of light a solution absorbs (Beer-Lambert Law). A calibration curve is created using standard solutions to determine the concentration of an unknown sample.
Interactive Determination
Enter the absorbance of your “unknown” sample to see its concentration determined from the graph.
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