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7.7: Solubility Rules for Ionic Compounds

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    289397
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    ⚙️ Learning Objectives

    • Define and give examples of electrolytes.
    • Determine the solubility of common ionic compounds.


    When some substances are dissolved in water, they undergo either a physical or a chemical change that yields ions in solution. These substances constitute an important class of compounds called electrolytes. Substances that do not yield ions when dissolved are called nonelectrolytes. If the physical or chemical process that generates the ions is essentially 100% efficient (all of the dissolved compound yields ions), then the substance is known as a strong electrolyte. If only a relatively small fraction of the dissolved substance undergoes the ion-producing process, it is called a weak electrolyte.

    Substances may be identified as strong, weak, or nonelectrolytes by measuring the electrical conductance of an aqueous solution containing the substance. To conduct electricity, a substance must contain freely mobile, charged species. Most familiar is the conduction of electricity through metallic wires, in which case the mobile, charged entities are electrons. Solutions may also conduct electricity if they contain dissolved ions, with conductivity increasing as ion concentration increases. Applying a voltage to electrodes immersed in a solution permits assessment of the relative concentration of dissolved ions, either quantitatively, by measuring the electrical current flow, or qualitatively, by observing the brightness of a light bulb included in the circuit (Figure \(\PageIndex{1}\)).
     

    This diagram shows three separate beakers. Each has a wire plugged into a wall outlet. In each case, the wire leads from the wall to the beaker and is split resulting in two ends. One end leads to a light bulb and continues on to a rectangle labeled with a plus sign. The other end leads to a rectangle labeled with a minus sign. The rectangles are in a solution. In the first beaker, labeled “Ethanol No Conductivity,” four pairs of linked small green spheres suspended in the solution between the rectangles. In the second beaker, labeled “K C l Strong Conductivity,” six individual green spheres, three labeled plus and three labeled minus are suspended in the solution. Each of the six spheres has an arrow extending from it pointing to the rectangle labeled with the opposite sign. In the third beaker, labeled “Acetic acid solution Weak conductivity,” two pairs of joined green spheres and two individual spheres, one labeled plus and one labeled minus are shown suspended between the two rectangles. The plus labeled sphere has an arrow pointing to the rectangle labeled minus and the minus labeled sphere has an arrow pointing to the rectangle labeled plus.
    Figure \(\PageIndex{1}\): Solutions of nonelectrolytes, such as ethanol, do not contain dissolved ions and cannot conduct electricity. Solutions of electrolytes contain ions that permit the passage of electricity. The conductivity of an electrolyte solution is related to the strength of the electrolyte.


    Ionic Electrolytes

    Water and other polar molecules are characterized by a slightly positive region and a slightly negative region and are therefore attracted to ions, as shown in Figure \(\PageIndex{2}\). The electrostatic attraction between an ion and a molecule with a dipole is called an ion-dipole attraction. These attractions play an important role in the dissolution of ionic compounds in water, which will be later discussed in Chapter 14.
     

    The diagram shows eight purple spheres labeled K superscript plus and eight green spheres labeled C l superscript minus mixed and touching near the center of the diagram. Outside of this cluster of spheres are seventeen clusters of three spheres, which include one red and two white spheres. A red sphere in one of these clusters is labeled O. A white sphere is labeled H. Two of the green C l superscript minus spheres are surrounded by three of the red and white clusters, with the red spheres closer to the green spheres than the white spheres. One of the K superscript plus purple spheres is surrounded by four of the red and white clusters. The white spheres of these clusters are closest to the purple spheres.
    Figure \(\PageIndex{2}\): As potassium chloride (KCl) dissolves in water, the ions are hydrated. The polar water molecules are attracted by the charges on the K+ and Cl ions. Water molecules in front of and behind the ions are not shown.


    When ionic compounds dissolve in water, the ions in the solid separate and disperse uniformly throughout the solution because water molecules surround and solvate the ions, reducing the strong electrostatic forces between them. This process represents a physical change known as dissociation. Under most conditions, ionic compounds will dissociate nearly completely when dissolved, and so they are classified as strong electrolytes.

    Let us consider what happens at the microscopic level when we add solid KCl to water. Ion-dipole forces attract the slightly positive (hydrogen) end of the polar water molecules to the negative chloride ions at the surface of the solid, and they attract the slightly negative (oxygen) end to the positive potassium ions. The water molecules penetrate between individual K+ and Cl ions and surround them, reducing the strong interionic forces that bind the ions together and letting them move off into solution as solvated ions, as Figure \(\PageIndex{2}\) shows. The reduction of the electrostatic attraction permits the independent motion of each hydrated ion in a dilute solution, resulting in an increase in the disorder of the system, as the ions change from their fixed and ordered positions in the crystal to mobile and much more disordered states in solution. This increased disorder is responsible for the dissolution of many ionic compounds, including KCl, which dissolve with absorption of heat.

    In other cases, the electrostatic attractions between the ions in a crystal are so large, or the ion-dipole attractive forces between the ions and water molecules are so weak, that the increase in disorder cannot compensate for the energy required to separate the ions, and the crystal is insoluble. Such is the case for compounds such as calcium carbonate (limestone), calcium phosphate (the inorganic component of bone), and iron oxide (rust).
     

    Solubility Rules

    Some combinations of aqueous reactants result in the formation of a solid precipitate as a product. However, some combinations will not produce such a product. If solutions of sodium nitrate and ammonium chloride are mixed, no reaction occurs. One could write an equation showing an exchange of ions; but both products, sodium chloride and ammonium nitrate, are soluble and remain in the solution as ions. It is useful to be able to predict when a precipitate will occur in a reaction. To do so, you can use a set of guidelines called solubility rules (Tables \(\PageIndex{1}\) and \(\PageIndex{2}\)).
     

    Table \(\PageIndex{1}\): Solubility Rules for Soluble Substances
    Soluble in Water Important Exceptions (Insoluble)
    All Group IA (alkali metals) and ammonium compounds
    (containing Li+, Na+, K+, Rb+, Cs+, NH4+)

    none
     
    All nitrates, chlorates, and acetates
    (containing NO3, ClO3, C2H3O2)

    none
     
    Most sulfates
    (containing SO42)

    Sulfates of Ca2+, Sr2+, Ba2+, Ag+, and Pb2+
     
    Most chlorides, bromides, and iodides
    (containing Cl, Br, I)

    Chlorides, bromides, and iodides of Ag+, Pb2+, and Hg22+
     

     

    Table \(\PageIndex{2}\): Solubility Rules for Insoluble Substances
    Insoluble in Water Important Exceptions (Soluble)
    Most carbonates, oxalates, and phosphates
    (containing CO32, C2O42, PO43)

    Group IA (alkali metals) and ammonium compounds
     

    Most hydroxides
    (containing OH​​​​)
     
    Group IA (alkali metals) and ammonium compounds

    Hydroxides of Ca2+, Sr2+, Ba2+ sparingly soluble – though usually considered insoluble

    Most sulfides
    (containing S2)
     
    Group IA (alkali metals) and ammonium compounds

    Sulfides of  Ca2+, Sr2+, Ba2+

     

    ✅ Example \(\PageIndex{1}\): Solubility

    Classify each compound as soluble or insoluble

    1. Zn(NO3)2
    2. PbBr2
    3. Sr3(PO4)2

    Solution

    1. Nitrates are soluble in water with no exceptions, so Zn(NO3)2 is soluble.
    2. Most bromides are soluble in water. However, combinations with Pb2+ are an exception, so PbBr2 is insoluble.
    3. Most phosphates are insoluble and there is no exception when combined with Sr2+, so Sr3(PO4)2 is insoluble.

     

    ✏️ Exercise \(\PageIndex{1}\): Solubility

    Classify each compound as soluble or insoluble.

    1. Mg(OH)2
    2. KBr
    3. Pb(NO3)2
    Answer A
    insoluble
    Answer B
    soluble
    Answer C
    soluble


    Summary

    • Substances that dissolve in water to yield ions are called electrolytes.
    • Nonelectrolytes are substances that do not produce ions when dissolved in water.
    • Solubility rules allow prediction of what products will be insoluble in water.

     

     


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