Heat of precipitation

The heat of precipitation is the heat change when one mole of a precipitate is formed from its ions in solution. Malay: haba pemendakan · Chinese: 沉淀热.

When two solutions are mixed and an insoluble solid forms, ions that were free in solution join together into a solid lattice, and heat is usually released as they do. The value is quoted per mole of the precipitate and, for the reactions you meet at SPM, it is negative because the change is exothermic.

Example. Mixing lead(II) nitrate solution with potassium iodide solution forms a yellow precipitate of lead(II) iodide: Pb²⁺ + 2I⁻ → PbI₂. Only the ions that end up in the solid appear in this ionic equation; the spectator ions stay in solution. You measure the temperature rise, use the total volume of the mixed solutions as the mass in Q = mcθ, and divide by the moles of precipitate to obtain the heat of precipitation in kJ mol⁻¹.

Confusion to avoid. Write the ionic equation, not the full equation, when you reason about a heat of precipitation, and count moles of the precipitate, not of one of the starting solutions. Using the full equation or the wrong mole count is a frequent slip that gives a value per the wrong quantity.

The heat of precipitation completes the set of four heats in this chapter, all measured by the same calorimetry routine, so SPM Chemistry can test any one of them with the single method of Q = mcθ followed by division by the moles. In the experiment the two solutions are usually measured at the same starting temperature before mixing, and the temperature change is taken from that starting temperature to the highest value reached once the precipitate has formed. Because only the ions that end up in the solid take part, writing the correct ionic equation first makes it easy to identify the precipitate and count its moles accurately before any numbers go into the heat equation.