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Desolvation removes some or all of the solvent environment of a species. Removing a hydration shell does not remove an ion's electric charge. Complete reversal of a specified solvation transfer reverses its Gibbs free energy and enthalpy changes when the same states and conventions are used. Partial desolvation, exchange of one solvent for another, and creation of a gaseous species have different endpoints; their costs need not match.

Dissolution transfers material into a chemical solution. It can include separation from an ionic crystal and solvation, or dispersal of neutral molecules. For sucrose and sodium chloride, respectively,Here $(aq)$ specifies water as the solvent; it does not specify a fixed number of attached water molecules. The neutral sucrose remains molecular. Dissociation of an ionic compound separates pre-existing ions in the second example; solvation stabilizes their new surroundings.^[ref][ref]

Ionization instead creates ions from neutral species. Hydrogen chloride reacting with water illustrates protonation of water rather than mere rearrangement of a hydration shell:The hydronium and negative ion together preserve total electric charge. Solvation can influence this chemical reaction, but an ion's solvation does not itself change its net electric charge.^[ref]

A hydration shell is a solvation shell in water. The existing hydration of ions discussion describes preferential orientations around cations and anions.

The hydrophobic core of a lipid bilayer makes transfer of a hydrated ion into bare lipid energetically costly. An ion channel supplies a water-accessible route across that barrier. Its narrow selectivity filter presents coordinating chemical groups that compensate for loss or rearrangement of the ion's hydration shell. Selectivity depends on the free-energy balance of hydration, coordination, electrostatics and geometry, rather than a sieve that passes all smaller ions.

A potassium-channel filter presents backbone carbonyl oxygens in a geometry favourable for potassium; smaller sodium does not automatically gain the same energetic compensation. Sodium and calcium channels use different filter architectures: charged side chains and their coordination environments favour particular ions, with ion occupancy and competing species also affecting permeation. In canonical vertebrate voltage-gated sodium channels, the DEKA selectivity ring illustrates chemical discrimination distinct from potassium-channel carbonyl coordination. The precise mechanism therefore depends on the channel family.

Rapid flux requires low successive free-energy barriers, not permanent binding of the preferred ion. Several occupied sites, nearby water and ion–ion interactions can facilitate passage without requiring the ion to cross the lipid. For scale, a $100\mathrm{pS}$ open channel driven by $100\mathrm{mV}$ carries $10\mathrm{pA}$, approximately $6\times 10^7$ monovalent ions per second. A nanometre-scale route with favourable coordination can support such rapid, thermally driven motion. An ion channel is not an ATP-driven pump: net flux follows the electrochemical potential difference,which vanishes at the species' Nernst equilibrium potential. ATP-consuming pumps maintain the gradients that make sustained channel currents possible.

Permeation selectivity and opening probability are separate properties. Voltage-sensor movement, ligand binding or mechanical tension can regulate access to the pore, while inactivation can terminate conduction during a maintained stimulus. Patch-clamp unitary currents establish discrete conducting states; reversal potentials under changed ionic compositions quantify selectivity. Structure determination locates the filter and gates, and mutations or specific blockers test their functional roles. Agreement between these measurements is stronger evidence than structure or current shape alone.

A chemical solution is a homogeneous mixture in which designated solutes are distributed in a solvent. Representative equal-volume samples have the same composition after mixing. Individual molecules and ions remain distinguishable. The solute/solvent labels describe their roles in this mixture.^[ref]

A dissolved solute is dispersed as molecules, chemical atoms, or ions, rather than as the larger solid particles of a particle suspension. For example, sucrose retains its molecular formula during dissolution in water; a chemical reaction is not required. Dissolved sodium chloride supplies sodium cations and chloride anions, surrounded by water molecules through solvation. Dissolved does not mean chemically destroyed or converted into water.^[ref]

This makes a useful connection between microscopic identity and laboratory sampling. If a chemical solution is well mixed, the amount of substance in a measured portion follows from its concentration and volume. A scoop from a poorly mixed particle suspension need not represent the whole vessel. Visual clarity alone cannot establish composition or exclude a fine dispersion; the quantitative claim needs representative sampling.

At fixed temperature and pressure, compare fully specified initial and final states usingThe entropy change is for the whole selected system, including both solute and solvent. Dispersing a solute can increase accessible arrangements while orienting nearby solvent molecules restricts others. One picture of an ordered hydration shell cannot determine the sign of the total entropy change. A negative Gibbs free energy change favors the specified process; it does not determine its reaction rate.^[ref]

An illustrative enthalpy accounting path separates solute constituents, reorganizes the solvent, and establishes their interactions:This is a chosen decomposition between compatible states, not a measured sequence. If a gaseous-ion-to-chemical solution solvation term already includes solvent reorganization, do not add that reorganization a second time. Lattice enthalpy from solution calorimetry gives the alternative Hess's law route that combines lattice separation with the complete hydration of ions term.^[ref]

The invented inputs are $+100$ for separation and $+20$ for solvent reorganization, in $\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$. Interaction terms of $-150$ in solvent A and $-90$ in solvent B give $\Delta H_A=-30$ and $\Delta H_B=+30$. Stipulated total entropy changes of $-50$ and $+150\mathrm{J}\mathrm{mo}{\mathrm{l}}^{-1}{\mathrm{K}}^{-1}$ give $-T\Delta S_A=+15$ and $-T\Delta S_B=-45\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$ at $300\mathrm{K}$. Both selected processes therefore have $\Delta G=-15\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$. This deliberately constructed comparison shows why the signs of the enthalpy and Gibbs free energy changes need not agree; it does not predict either solvent's solubility.

For incremental dissolution at a particular composition, the relevant chemical potential difference also depends on thermodynamic activities. At chemical equilibrium that driving difference vanishes. A quoted standard Gibbs free energy change and the finite change between arbitrary mixtures answer different questions; use compatible standard states and the reaction quotient when discussing solubility.

The motivating calorimetry experiment is especially instructive. A weighed potassium chloride sample dissolves in a known amount of water in a platinum-lined, shielded calorimeter; a quartz-oscillator thermometer and measured electrical heating provide the readouts. Electrical calibration before and after the reaction, comparison with a resistance thermometer, and checks on residual water and particle size constrain errors. The inspected authored article reports positive enthalpy of solution at $298.15\mathrm{K}$ for the specified dilution, despite complete dissolution. This is a combined enthalpy balance, not an isolated solvation free energy or direct single-ion hydration of ions measurement. Shield tracking and final concentration matter; the existing lattice enthalpy from solution calorimetry discussion gives the numerical result and the conversion limits.^[ref]

What separates a hydration shell from the surrounding water experimentally? Electrical conductivity addresses a different question. The documented lamp demonstration immerses metal electrodes in a chemical solution, rinses them between samples, and records brightness; adding sodium chloride to distilled water makes it conductive. A water blank and dissolved sucrose comparison test whether dissolved material alone explains the signal. Fixed geometry, temperature, and stated concentrations are additional proposed controls. This reveals mobile electric charge, with limits from contamination, electrode interfaces and detection threshold; it supplies no shell positions. Conductivity experiment on salt solutions explains the apparatus and controls in detail.^[ref]

In the authored study Influence of ions on the hydrogen-bond structure in liquid water by Anne Willem Omta and others, pump-probe spectroscopy uses tunable mid-infrared pulses traversing a sample. The probe's wave polarization is set at $45^{\circ }$ to the pump; after the sample, splitting and polarizers select parallel and perpendicular transmission. Measuring each transmission with and without the pump gives $D_{\parallel }$ and $D_{\perp }$, the respective logarithms of the transmission ratios. Their dimensionless anisotropy isIts decay probes reorientation rather than merely loss of vibrational excitation. Dilute semiheavy water in heavy water separates the O–H and O–D bands. Varying the semiheavy water concentration checked that transfer of excitation between nearby O–H groups did not determine the decay. These apparatus and control details are from the inspected author-uploaded article, not a new experiment.^[ref]

The related thesis chapter documents two O–H absorption bands in a sodium perchlorate chemical solution: approximately $3400\mathrm{c}{\mathrm{m}}^{-1}$ for groups forming hydrogen bonds to water, and $3575\mathrm{c}{\mathrm{m}}^{-1}$ for groups forming hydrogen bonds to the anion. Selecting pump and probe bands, then a later delay window, separated their responses. The anion-associated anisotropy decayed more slowly; the water-bonded component showed no resolved change in decay time across the tested sodium perchlorate concentrations. The chapter also used O–D excitation to extend observation beyond the shorter O–H vibrational lifetime. These results support distinct local dynamics rather than a permanently fixed hydration shell. Spectral overlap, finite observation windows, and the selected salts limit the inference: this does not prove that every ion leaves all distant water properties unchanged.^[ref]