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13.3: Solutions of Solids Dissolved in Water- How to Make Rock Candy

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  • Learning Objectives

    • Define electrolytes and non electrolytes
    • Explain why solutions form.
    • Discuss the idea of water as the "universal solvent".
    • Explain how water molecules attract ionic solids when they dissolve in water.

    We have learned that solutions can be formed in a variety of combinations using solids, liquids, and gases. We also know that solutions have constant composition, and that this composition can be varied up to a point to maintain the homogeneous nature of the solution. But how exactly do solutions form? Why is it that oil and water will not form a solution, and yet vinegar and water will? Why could we dissolve table salt in water, but not in vegetable oil? The reasons why solutions will form will be explored in this section, along with a discussion of why water is used most frequently to dissolve substances of various types.

    Solubility and Saturation

    Table salt \(\left( \ce{NaCl} \right)\) readily dissolves in water. In most cases, only a certain maximum amount of solute can be dissolved in a given amount of solvent. This maximum amount is specified as the solubility of the solute. It is usually expressed in terms of the amount of solute that can dissolve in 100 g of the solvent at a given temperature. Table \(\PageIndex{1}\) lists the solubilities of some simple ionic compounds. These solubilities vary widely. NaCl can dissolve up to 31.6 g per 100 g of H2O, while AgCl can dissolve only 0.00019 g per 100 g of H2O.

    Table \(\PageIndex{1}\): Solubilities of Some Ionic Compounds
    Solute Solubility (g per 100 g of H2O at 25°C)
    AgCl 0.00019
    CaCO3 0.0006
    KBr 70.7
    NaCl 36.1
    NaNO3 94.6

    When the maximum amount of solute has been dissolved in a given amount of solvent, we say that the solution is saturated with solute. When less than the maximum amount of solute is dissolved in a given amount of solute, the solution is unsaturated. These terms are also qualitative terms because each solute has its own solubility. A solution of 0.00019 g of AgCl per 100 g of H2O may be saturated, but with so little solute dissolved, it is also rather dilute. A solution of 36.1 g of NaCl in 100 g of H2O is also saturated, but rather concentrated. In some circumstances, it is possible to dissolve more than the maximum amount of a solute in a solution. Usually, this happens by heating the solvent, dissolving more solute than would normally dissolve at regular temperatures, and letting the solution cool down slowly and carefully. Such solutions are called supersaturated solutions and are not stable; given an opportunity (such as dropping a crystal of solute in the solution), the excess solute will precipitate from the solution.The figure below illustrates the above process and shows the distinction between unsaturated and saturated.

    Figure \(\PageIndex{1}\): When \(30.0 \: \text{g}\) of \(\ce{NaCl}\) is added to \(100 \: \text{mL}\), it all dissolves, forming an unsaturated solution. When \(40.0 \: \text{g}\) is added, \(36.0 \: \text{g}\) dissolves and \(4.0 \: \text{g}\) remains undissolved, forming a saturated solution.

    How can you tell if a solution is saturated or unsaturated? If more solute is added and it does not dissolve, then the original solution was saturated. If the added solute dissolves, then the original solution was unsaturated. A solution that has been allowed to reach equilibrium, but which has extra undissolved solute at the bottom of the container, must be saturated.

    Electrolyte Solutions: Dissolved Ionic Solids

    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 (good conductor). If only a relatively small fraction of the dissolved substance undergoes the ion-producing process, the substance is a weak electrolyte (does not conduct electricity as well).

    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.

    Water and other polar molecules are 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.

    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.

    Example \(\PageIndex{1}\): Identifying Ionic Compounds

    Which compound(s) will dissolve in solution to separate into ions?

    1. \(\ce{LiF}\)
    2. \(\ce{P_2F_5}\)
    3. \(\ce{C_2H_5OH}\)


    \(\ce{LiF}\) will separate into ions when dissolved in solution, because it is an ionic compound. \(\ce{P_2F_5}\) and \(\ce{C_2H_5OH}\) are both covalent and will stay as molecules in a solution.

    Exercise \(\PageIndex{1}\)

    Which compounds will dissolve in solution to separate into ions?

    1. C6H12O11, glucose
    2. CCl4
    3. CaCl2
    4. AgNO3
    c & d

    How Temperature Influences Solubility

    The solubility of a substance is the amount of that substance that is required to form a saturated solution in a given amount of solvent at a specified temperature. Solubility is often measured as the grams of solute per \(100 \: \text{g}\) of solvent. The solubility of sodium chloride in water is \(36.0 \: \text{g}\) per \(100 \: \text{g}\) water at \(20^\text{o} \text{C}\). The temperature must be specified because solubility varies with temperature. For gases, the pressure must also be specified. Solubility is specific for a particular solvent. We will consider solubility of material in water as solvent.

    The solubility of the majority of solid substances increases as the temperature increases. However, the effect is difficult to predict and varies widely from one solute to another. The temperature dependence of solubility can be visualized with the help of a solubility curve, a graph of the solubility vs. temperature (Figure \(\PageIndex{4}\)).

    Figure \(\PageIndex{4}\): Solubility curves for several compounds.

    Notice how the temperature dependence of \(\ce{NaCl}\) is fairly flat, meaning that an increase in temperature has relatively little effect on the solubility of \(\ce{NaCl}\). The curve for \(\ce{KNO_3}\), on the other hand, is very steep and so an increase in temperature dramatically increases the solubility of \(\ce{KNO_3}\).

    Several substances—\(\ce{HCl}\), \(\ce{NH_3}\), and \(\ce{SO_2}\)—have solubility that decreases as temperature increases. They are all gases at standard pressure. When a solvent with a gas dissolved in it is heated, the kinetic energy of both the solvent and solute increase. As the kinetic energy of the gaseous solute increases, its molecules have a greater tendency to escape the attraction of the solvent molecules and return to the gas phase. Therefore, the solubility of a gas decreases as the temperature increases.

    Solubility curves can be used to determine if a given solution is saturated or unsaturated. Suppose that \(80 \: \text{g}\) of \(\ce{KNO_3}\) is added to \(100 \: \text{g}\) of water at \(30^\text{o} \text{C}\). According to the solubility curve, approximately \(48 \: \text{g}\) of \(\ce{KNO_3}\) will dissolve at \(30^\text{o} \text{C}\). This means that the solution will be saturated since \(48 \: \text{g}\) is less than \(80 \: \text{g}\). We can also determine that there will be \(80 - 48 = 32 \: \text{g}\) of undissolved \(\ce{KNO_3}\) remaining at the bottom of the container. Now suppose that this saturated solution is heated to \(60^\text{o} \text{C}\). According to the curve, the solubility of \(\ce{KNO_3}\) at \(60^\text{o} \text{C}\) is about \(107 \: \text{g}\). Now the solution is unsaturated since it contains only the original \(80 \: \text{g}\) of dissolved solute. Now suppose the solution is cooled all the way down to \(0^\text{o} \text{C}\). The solubility at \(0^\text{o} \text{C}\) is about \(14 \: \text{g}\), meaning that \(80 - 14 = 66 \: \text{g}\) of the \(\ce{KNO_3}\) will re-crystallize.


    • Solubility is the specific amount of solute that can dissolve in a given amount of solvent.
    • Saturated and unsaturated solutions are defined.
    • Ionic compounds dissolve in polar solvents, especially water. This occurs when the positive cation from the ionic solid is attracted to the negative end of the water molecule (oxygen) and the negative anion of the ionic solid is attracted to the positive end of the water molecule (hydrogen).
    • Water is considered the universal solvent since it can dissolve both ionic and polar solutes, as well as some nonpolar solutes (in very limited amounts).
    • The solubility of a solid in water increases with an increase in temperature.


    • Miscible - Liquids that have the ability to dissolve in each other.
    • Immiscible - Liquids that do not have the ability to dissolve in each other.
    • Electrostatic attraction - The attraction of oppositely charged particles.

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