Hydration of Ions
solubilityThe extent to which a solute dissolves in a solvent; often expressed as the mass of a substance that will dissolve in 100 mL of solvent. in water is a characteristic property of many ionic compounds, and Table 1 from Precipitation Reactions provides further confirmation of this fact. We have also presented experimental evidence that ions in solution are nearly independent of one another. This raises an important question, though, because we have also stated that attractive forces between oppositely charged ions in a crystal latticeAn orderly, repeating arrangement of points in 3-D space in which each p;oint has surroundings identical to every other point. A crystal's constituent atoms, molecules, and ions are arranged about each lattice point. are large. The high meltingThe process of a liquid forming from a solid. and boiling points of ionic compounds provide confirmation of the expected difficulty of separating oppositely charged ions. How, then, can ionic compounds dissolve at room temperatureA physical property that indicates whether one object can transfer thermal energy to another object.? Surely far more energyA system's capacity to do work. would be required for an ion to escape from the crystalA solid with a regular polyhedral shape; for example, in sodium chloride (table salt) the crystal faces are all at 90° angles. A solid in which the atoms, molecules, or ions are arranged in a regular, repeating lattice structure. lattice into solution than even the most energetic ions would possess.
The resolution of this apparent paradox lies in the interactions between ions and the molecules of water or other polarDescribes a molecule that has separated, equal positive and negative charges that consitute a positive and a negative pole; such a molecule tends to assume certain orientations more than others in an electric field. solvents. The negative (oxygen) side of a dipolar water molecule attracts and is attracted by any positive ion in solution. Because of this ion-dipole forceThe attraction between polar molecules as a result of the partially positively charged portion of one molecule being oriented toward the partially negatively charged portion of another molecule., water molecules cluster around positive ions, as shown in Figure 1a. Similarly, the positive (hydrogen) ends of water molecules are attracted to negative ions. This process, in which either a positive or a negative ion attracts water molecules to its immediate vicinity, is called hydration.
When water molecules move closer to ions under the influence of their mutual attraction, there is a net lowering of the potential energy of the microscopic particles. This counteracts the increase in potential energy which occurs when ions are separated from a crystal lattice against their attractions for other ions.
Thus the process of dissolving an ionic solidA state of matter having a specific shape and volume and in which the particles do not readily change their relative positions. may be divided into the two hypothetical steps shown in Fig. 2. First, the crystalline saltAn ionic compound that can be formed by replacing the hydrogen ion of an acid with a different cation. is separated into gaseous ions. The heatEnergy transferred as a result of a temperature difference; a form of energy stored in the movement of atomic-sized particles. energy absorbed when the ions are separated this way is called the lattice enthalpyA thermodynamic state function, symbol H, that equals internal energy plus pressure x volume; the change in enthalpy corresponds to the energy transferred as a result of a temperature difference (heat transfer) when a reaction occurs at constant pressure. (or sometimes the lattice energyThe heat energy tranfer into a system as gaseous ions come together to form an ionic compound. Different textbooks define the sign of this quantity differently.). Next, the separate ions are placed in solution; that is, water molecules are permitted to surround the ions. The enthalpy change for this process is called the hydration enthalpy.
Since there is a lowering of the potential energy of the ions and water molecules, heat energy is given off and hydration enthalpies are invariably negative.
The heat energy absorbed when a soluteThe substance added to a solvent to make a solution. dissolves (at a pressureForce per unit area; in gases arising from the force exerted by collisions of gas molecules with the wall of the container. of 1.00 atmAbbreviation for atmosphere, a unit of pressure equal to 101.325 kPa or 760 mmHg.) is called the enthalpy of solution. It can be calculated using Hess' law, provided the lattice enthalpy and hydration enthalpy are known.
EXAMPLE 1 Using data given in Fig. 2, calculate the enthalpy of solution for NaCl(s).
Solution According to the figure, the lattice enthalpy is 773 kJ mol–1. The hydration enthalpy is – 769 kJ mol–1. Thus we can write the thermo-chemical equations
NaCl(s) → Na+(g) + Cl–(g) ΔHl = 773 kJ mol–1
Na+(g) + Cl–(g) → Na+(aq) + Cl–(aq) ΔHh = –769 kJ mol–1
NaCl(s) → Na+(aq) + Cl–(aq) ΔHs = ΔHl + ΔHh
ΔHs = (773 – 769) kJ mol–1 = +4 kJ mol–1
When NaCl(s) dissolves, 773 kJ is required to pull apart a mole of Na+ ions from a mole of Cl– ions, but almost all of this requirement is provided by the 769 kJ released when the mole of Na+ and the mole of Cl– becomes surrounded by water dipoles. Only 4 kJ of heat energy is absorbed from the surroundings when a mole of NaCl(s) dissolves. You can verify the small size of this enthalpy change by putting a few grains of salt on your moist tongue. The quantity of heat energy absorbed as the salt dissolves is so small that you will feel no cooling, even though your tongue is quite a sensitive indicatorA substance for which a physical property (such as color) changes abruptly when the equivalence point is reached in a titration. of temperature changes.
Few molecules are both small enough and polar to cluster around positive and negative ions in solution as water does. Consequently water is one of the few liquids which readily dissolves many ionic solids. Hydration of Na+, Cl– and other ions in aqueous solution prevents them from attracting each other into a crystal lattice and precipitating.