Water hydrated ionic compounds

Naming Hydrates (Water - Containing Ionic Compounds)... 

Many ionic compounds are prepared by crystallization from a water solution, and water molecules become a part of the crystal. A compound in which there is a specific ratio of water to ionic compound is called a hydrate. In a hydrate, the water molecules are chemically bonded to the ionic compound. 

Epsom salts is a hydrated ionic compound with the foUowing formula MgS04 A 4.93 g sample of Epsom salts was heated to drive off the water of hydration. The mass of the sample after complete dehydration was 2.41 g. Find the number of waters of hydration (x) in Epsom salts. 

It is soluble in organic solvents (a characteristic of a covalent compound). but dissolves in water and can form hydrates (a characteristic of an ionic compound), hence the hydrated must be... 

Ionic compounds often separate from water solution with molecules of water incorporated into the solid. Such compounds are referred to as hydrates. An example is hydrated copper sulfate, which contains five moles of H20 for every mole of CuS04. Its formula is CuS04- 5H20 a dot is used to separate the formulas of the two compounds CuS04 and H20. A Greek prefix is used to show the number of moles of water the systematic name of CuS04- 5H20 is copper(ll) sulfate pentahydrate. 

Ionic compounds are named by starting with the name of the cation (with its oxidation number if more than one charge is possible), followed by the name of the anion hydrates are named by adding the word hydrate, preceded by a Greek prefix indicating the number of water molecules in the formula unit. 

Many ionic compounds can have water molecules incorporated into their solid structures. Such compounds are called hydrates. To emphasize the presence of discrete water molecules in the chemical structure, the formula of any hydrate shows the waters of hydration separated from the rest of the chemical formula by a dot. A coefficient before H2 O indicates the number of water molecules in the formula. Copper(II) sulfate pentahydrate is a good example. The formula of this beautiful deep blue solid is C11SO4 5 H2 O, indicating that five water molecules are associated with each CuSOq unit. Upon prolonged heating, CuSOq 5 H2 O loses its waters of hydration along with its color. Other examples of hydrates include aluminum nitrate nonahydrate, A1 (N03)3 9 H2 O,... 

Whereas we can deduce the ratios of anions to cations in an ionic compound from the eiectricai charges on the individuai ions, we cannot determine the number of water moiecuies in a hydrate from the nature of the anions and cations in the ionic compound. The number of water moiecuies, which can range from 0 to as high as 18, must be determined by doing experiments. In fact, some ionic substances exist in several different forms with different numbers of water moiecuies. 

In cases where the solvation energies are large, as for example when ionic compounds dissolve in water, these hydrophobic effects, based on adverse changes in entropy, are swamped. Dissolving such compounds can be readily accomplished due to the very large energies released when the ions become hydrated. 

Many ionic compounds contain what used to be referred to as water of crystallization . For example, magnesium chloride can exist as a fully hydrated salt which was formerly written MgCla.bHjO, but is more appropriately written Mg(OH2)eCl2, since the water molecules occupy coordination sites around the magnesium ions. This is typical. In most compounds that contain water of crystallization, the water molecules are bound to the cation in an aquo complex in the manner originally proposed by Alfred Werner (1866-1919) in 1893 (Kauffman, 1981). Such an arrangement has been confirmed in numerous cases by X-ray diffraction techniques. 

When an ionic compound dissolves in water, energy is needed to break the ionic bonds of the crystal. As the ions attach to the water molecules and become hydrated, energy is released. The process is endothermic if the energy needed to break the bonds is greater than the energy released when the ions attach to water. 

Water is the most common solvent used to dissolve ionic compounds. Principally, the reasons for dissolution of ionic crystals in water are two. Not stated in any order of sequence of importance, the first one maybe mentioned as the weakening of the electrostatic forces of attraction in an ionic crystal known, and the effect may be alternatively be expressed as the consequence of the presence of highly polar water molecules. The high dielectric constant of water implies that the attractive forces between the cations and anions in an ionic salt come down by a factor of 80 when water happens to be the leaching medium. The second responsible factor is the tendency of the ionic crystals to hydrate. 

When an ionic compound is dissolved in a solvent, the crystal lattice is broken apart. As the ions separate, they become strongly attached to solvent molecules by ion-dipole forces. The number of water molecules surrounding an ion is known as its hydration number. However, the water molecules clustered around an ion constitute a shell that is referred to as the primary solvation sphere. The water molecules are in motion and are also attracted to the bulk solvent that surrounds the cluster. Because of this, solvent molecules move into and out of the solvation sphere, giving a hydration number that does not always have a fixed value. Therefore, it is customary to speak of the average hydration number for an ion. 

The water that is trapped within the crystal structure of some ionic compounds (the water of hydration) can be removed easily by heating. The amount of the water present in a given sample can be determined by weighing the sample before and after this heating. The weight loss that occurs is the weight of the water in the sample. The percent of water in the hydrate is calculated as it is in a loss-on-drying experiment. 

Reactive ionic compounds are therefore useless to derive hydration enthalpies (or more generally, solvation enthalpies). Fortunately, there are many alternatives. Take lithium chloride, for example, and data from the NBS Tables [ 17]. The enthalpy of solution of this solid in water, at infinite dilution, is given by... 

Hydrates Some ionic compounds incorporate a fixed number of water molecules into their formula unit. The compound that contains the water is called a hydrate, and removal of the water affords the anhydrous salt. Compounds that have a strong tendency to absorb water are called hygroscopic. To name a hydrate, you simply name the ions and then add the appendage hydrate, along with a multiplier to indicate the number of water molecules in the formula. 

Some types of ionic compounds can absorb water so that each formula unit is attached to a specific number of water molecules. They are called hydrates. Ba0H2 8H20 is called barium hydroxide octahydrate. CaS04 2H20 is called calcium sulfate dihydrate. Can you see the pattern Try naming MgS04 7H20. Use Table 3.8 to help you. You will learn more about hydrates in Chapter 6. 

For example, many ionic compounds crystallize from a water solution with water molecules incorporated into their crystal structure, forming a hydrate. Hydrates have a specific number of water molecules chemically bonded to each formula unit. A chemist may know the formula of the ionic part of the hydrate but not how many water molecules are present for each formula unit. 

Epsom salts, for example, consist of crystals of magnesium sulfate heptahydrate, MgS04-7H20. Every formula unit of magnesium sulfate has seven molecules of water weakly bonded to it. A raised dot in a chemical formula, in front of one or more water molecules, denotes a hydrated compound. Note that the dot does not include multiplication, but rather a weak bond between an ionic compound and one or more water molecules. Some other examples of hydrates are shown in Table 6.4. 

As you have just discovered, calculations involving hydrates usually involve comparing the anhydrous form of the ionic compound to the hydrated form. Many chemicals are available in hydrated form. Usually chemists are only interested in how much of the ionic part of the hydrate they are working with. This is because, in most reactions involving hydrates, the water portion of the compound does not take part in the reaction. Only the ionic portion does. 

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