= Lattice enthalpy from solution calorimetry
A second indirect route makes the experimental distinction especially clear: a <calorimeter> can measure dissolution of a weighed <salt> into a known amount of <water>, but its final <ions> are hydrated, not separated gaseous <ions>. At compatible <standard states> and the same <temperature>,
$$
\Delta H_{\rm sol}^{\circ}=L_{\rm diss}+\Delta H_{\rm hyd,pair}^{\circ},\qquad
L_{\rm diss}=\Delta H_{\rm sol}^{\circ}-\Delta H_{\rm hyd,pair}^{\circ}.
$$
Here the combined <hydration of ions> term takes one <mole> each of gaseous $\mathrm{Na}^+$ and $\mathrm{Cl}^-$ to the specified aqueous reference state. Its sum suffices: dissolution <calorimetry> alone cannot independently assign single-ion <hydration of ions> <enthalpies>, or determine $L_{\rm diss}$ without another input. A finite-<concentration> measurement also requires dilution corrections before it can supply $\Delta H_{\rm sol}^{\circ}$.
In a simple constant-<pressure> solution <calorimeter>, an insulated vessel, stirrer and <thermometer> follow the <temperature> change after mixing. A calibrated effective <heat capacity> $C_{\rm eff}$ converts the corrected <temperature> change to <heat>. With no applied heating, negligible external <heat transfer>, and only pressure-volume <work>, the approximate molar result is
$$
\Delta H_{\rm sol}\simeq-\frac{C_{\rm eff}\Delta T}{n}.
$$
Here $n$ is the <amount of substance> dissolved and $C_{\rm eff}$ includes the <chemical solution> and the vessel. Cooling means positive dissolution <enthalpy>, because the dissolving sample takes <heat> from its surroundings.<Chemistry 2e>[§5.2, Figures 5.11–5.13 and discussion of including solution and calorimeter heat capacities]{ref}
An actual reference measurement used <potassium chloride>, not <sodium chloride>: a platinum-lined <calorimeter> with an adiabatic shield and a quartz-oscillator <thermometer>. Because dissolution absorbed <heat>, measured electrical heating prevented loss of shield control; the readouts were electrical input and the vessel/shield <temperature> record. Stirring was checked visually in a dummy vessel. The <thermometer> was compared with a calibrated resistance <thermometer>, and electrical calibration bracketed the reaction. The reported result was $+(17.584\pm0.017)\,\mathrm{kJ\,mol^{-1}}$ at $298.15\,\mathrm K$, for one <mole> of <potassium chloride> in $500$ <moles> of <water>.<The Enthalpy of Solution of SRM 1655 (KCl) in H2O by Marthada V. Kilday>[§3 and §4: apparatus, calibration, readouts and specified dilution; inspected through indexed article text]{ref}
Checks on residual <water>, particle size and final <concentration>, and agreement with a second <calorimeter>, constrained systematic errors. Imperfect shield tracking required <heat transfer> corrections. This experiment measures an <enthalpy of solution>, not a direct <lattice energy>; the hydration and dilution bookkeeping still has to be supplied. The full article could not be opened here; the inspected indexed text supports the stated apparatus and checks, without reconstructing its complete uncertainty budget.<The Enthalpy of Solution of SRM 1655 (KCl) in H2O by Marthada V. Kilday>[§§2–4 and conclusion: moisture, particle-size and dilution studies; confirmation with an isoperibol calorimeter; full-page access blocked]{ref}
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