Factors Affecting Solubility: Solubility, Factors Affecting Solubility, Hydration Enthalpy and Lattice Enthalpy, Factors affecting Hydration Enthalpy and Lattice Enthalpy, Solubility Product, Practice Problems & Frequently Asked Questions

You might have prepared salt and sugar solutions. But have ever examined the maximum amount of salt or sugar soluble in the same quantity of water? Do you think that will be the same or different? You might have heard about crystallisation. Do you know the principle of this crystallisation process?

Each substance has a unique capacity of dissolution in a solvent and shows a maximum limit. The nature of the solvent and the solute, as well as their level of aggregation, temperature, and pressure, all affect how soluble they are. Let's examine one of the most well-known chemistry solubility examples.

A chemical compound is composed of numerous identical molecules that are created from atoms of different chemically bonded elements. All connections, however, are not created equal. When ionic compounds dissolve in water, different things take place.

Let’s see the example of salt which is ionic compound added to water

Ionic compounds are molecules made up of oppositely charged ions with both a negative and positive charges. When ionic compounds dissolve in water, they dissociate into the ions from which they are formed. The ions are hydrated by the water molecules.

Table of Contents:

• Solubility
• Factors Affecting Solubility
• Hydration Enthalpy and Lattice Enthalpy
• Factors affecting Hydration Enthalpy and Lattice Enthalpy
• Solubility product
• Practice Problems
• Frequently Asked Questions

Solubility:

With a few exceptions, water is a universal solvent that can dissolve almost any solute. Solubility is described quantitatively as the greatest concentration of a solute that may be dissolved in a known concentration of a solvent at a specific temperature.

Solutes are categorised as being insoluble, sparingly soluble, or highly soluble depending on how much of the solute is dissolved in the solvent.

It is stated to be soluble if 0.1 g or more of the solute can be dissolved in 100 mL of the solvent.
It is stated to be sparingly soluble if a concentration of the solute dissolves in the solvent at less than 0.1 g.
Solutions are divided into three types based on their solubility: saturated, unsaturated, and supersaturated.
A saturated solution is one that is totally dissolved in a solvent at a given temperature.
An unsaturated solution is one that can dissolve more solutes at the same temperature.
A supersaturated solution is one in which the solute begins to precipitate once a specified quantity is dissolved at the same temperature.

Factors Affecting Solubility:

1. Nature of the solute and solvent

It is evident that a solute is soluble in nature if its type closely resembles that of the solvent. For instance, it is simple to dissolve ethanol, salt or sugar in water, since all of these substances are polar solutes and solvents. Naphthalene is a nonpolar solute, therefore it will be difficult to dissolve it in water and it can be dissolved only in non polar solvents like carbon tetrachloride, benzene etc.

Generally ionic solutes are soluble in water and polar solvents. And covalent compounds in nonpolar solvents. Still, covalent cmopunds like carbohydrates can dissolve in water if they can form hydrogen bonds with water.

2. Concentration of the Solute in solvent:

The solid particles are dissolved in the solution when the solute added is less or equal to their solubility.. Dissolution is the name given to this process.

At high concentrations, when solute particles and those in solution collide, some of the solute particles separate from the solution and crystallisation occurs.

When the amount of solute molecules entering a solution equals the number of molecules leaving it, a dynamic equilibrium state will subsequently be reached.

3. Temperature:

Temperature effects on solid solubility change depending on whether the process is endothermic or exothermic. The effects of temperature in both circumstances can be determined using Le Chatelier's principle.

• Think about a reaction that is endothermic (Δ H_solvation, < 0). According to Le-Chatelier's principle, the system will shift towards the side of the products to ease the stress when the temperature rises due to the added heat because there is stress in the reactant side. More of the solid is dissociated and the solubility increases when equilibrium is once more established by shifting towards the product side.

• Think of a exothermic reaction (ΔH_solvation > 0). According to Le-Chatelier's principle, the system will shift towards the side of the reactants to ease the stress when the temperature rises because there is stress in the product side. Less of the solid is dissociated, resulting in a reduced solubility, when equilibrium is once more established by shifting towards the reactant side.
• It is critical to remember that temperature is a measure of average kinetic energy. When studying the impact of temperature on gas solubility, kinetic energy rises as temperature rises. More the kinetic energy, the greater the molecular mobility of the gas particles. As a result, dissolved gas particles are more likely to escape to the gas phase, whereas existing gas particles are less likely to be dissolved. The inverse is also true. Thus, increased temperatures result in reduced solubility, while decreased temperatures result in greater solubility.

• The Le-Chatelier's approach enables a more precise comprehension of these trends. Remember that the process of dissolving gas in liquid is often exothermic. Consequently, raising the temperature stresses the product (because heat is on the product side). Le-Chatelier's principle predicts that the system will shift to the reactant side in order to reduce this increased stress. As a result, the equilibrium concentration of the gas particles in the gaseous phase decreases solubility.

• On the other hand, lowering the temperature stresses the reactant side (because heat is on the product side). Le-Chatelier's concept predicts that the system would move toward the product side to make up for this new stress. As a result, the equilibrium concentration of gas particles drops in the gaseous phase, increasing solubility.

4. Pressure:

i) Pressure has negligible effect on the solubility of solids.

ii) Solubility of Gas in Liquid:

The effect of pressure on the solubility of gases in liquids can be best explained by Henry’s Law.

Henry’s law:

It states that “the partial pressure of the gas in vapor phase(P) is directly proportional to the mole fraction of the gas (x) in the solution” and is expressed as

$$\begin{gathered} P \alpha x \\ P = K_H x \end{gathered}$$

Where, K_H = Henry’s law constant.

According to this formula, when the partial pressure lowers in a liquid (at a fixed temperature), the concentration of gas in the liquid also goes down, which therefore also affects solubility. In contrast, as the partial pressure rises in this case, the amount of gas in the liquid will also rise, increasing its solubility. Le-Chatelier's principle is more helpful in predicting the effects of pressure on the solubility of gases by extending the implications from Henry's law.

Consider a scenario in which a gas is partially dissolved in a liquid. A rise in pressure would result in a rise in partial pressure (because the gas is being further compressed). Because of the increasing partial pressure, more gas particles will enter the liquid (there is thus less gas above the liquid, therefore the partial pressure lowers) to relieve the stress caused by the pressure rise, resulting in enhanced solubility. In such a system, the opposite is also true, since a reduction in pressure results in more gas particles escaping the liquid to compensate.

5. Solubility of Liquid in Liquid:

Temperature: Increase in temperature increases the solubility of the solute.

Pressure: The effect of pressure on the solubility of liquids is negligible.

6. Surface Area and Its Effect on the Rate of Solubility:

Another thing to consider is how the surface area factor influences the rate of solubility. If the surface area of a solid is to be increased, the solid must be chopped into smaller portions, which aids in the dissolution of the solute in the solution. When a large amount of sugar is dissolved in water, the sugar cube crystal dissolves more slowly than an equal number of microscopic sugar crystals. Because the surface area of all the sugar crystals present is significantly larger than the surface area of a single sugar cube, the sugar crystals will have more contact with the water molecules. This is the primary reason why sugar crystals disintegrate so quickly.

When working in a lab, you may be asked to prepare a copper(II) sulphate solution. Copper(II) sulphate is available in a range of shapes and sizes, including large and little blue crystals. When equal volumes of both forms are placed in test tubes containing 10 mL of water, it is discovered that after 5 minutes, the afine crystals have dissolved more (and the solution has darkened blue) than the huge crystals contained in another test tube. By taking two samples, these can be broken into tiny crystals and placed in two different test tubes. Place the cork in one of the test tubes and gently shake it while the other test tube remains still.

Shaking enhances the amount of fine crystals in contact with the water, resulting in an increase or increment in surface area. The result is the same as before, and the test tube with the larger surface area dissolves more faster. Although increasing the surface area provides for faster achievement of maximum solubility, the solute concentration at maximum solubility stays unchanged.

7. Hydration Enthalpy and Lattice enthalpy:

When sonic solids are dissolved in water two things happen.

1. Ionic solid breaks into cations and anions against the Lattice energy that holds them together in solid
2. .Both cations and anions are attracted and surrounded by the polar water molecules releasing hydration energy..

Solubility of the ionic solid is determined by the relative energy difference between lattice energy ad hydration energy. Lower lattice energy assisit easy information f ions and higher hydration energy stabilises the hydrated ions.

When ions are mixed with water, a huge number of dipolar water molecules surround them, attract and hold them. Any positive ion in the mixture will draw a dipolar water molecule's negative (oxygen) side. Ion-dipole interactions cause water molecules to collect near positive ions. Similarly negative ions pull water molecules' positive (hydrogen) ends toward them. A positive or negative ion brings water molecules close to them through the process of hydration.

When water molecules move closer to the ions due to their mutual attraction, the potential energy of tiny particles is decreased. By reducing their attraction to one another, ions are less likely to separate from a crystal lattice, which would otherwise result in a rise in potential energy.

As a result, two imaginary processes can be used to separate the dissolving of an ionic solid. The constituent gaseous ions of the crystalline salt are first isolated. The heat energy absorbed when the ions are separated in this way is known as the lattice enthalpy (or sometimes the lattice energy). The next step is to put the ions in solution, allowing water molecules to envelop them. The enthalpy change for this process is known as the hydration enthalpy.

An ionic solid dissolves quickly in water if the lattice energy is less and higher hydration energy.

Factors Affecting Hydration Enthalpy and Lattice Enthalpy:

Charge on ion: The magnitude of the attractive force will increase with the size of the ion charge, and as a result, the lattice enthalpy and hydration enthalpy will increase. For instance, magnesium oxide (MgO) lattice enthalpy is significantly higher than sodium chloride (NaCl).

Size of ion: The value of lattice enthalpy and hydration enthalpy increases when ion size decreases, internuclear distance decreases, and interionic attraction increases.

Because their interionic distances are closer, smaller ions have stronger binding forces. The result of this is higher lattice enthalpy and hydration enthalpy. For instance, in group 16 of the periodic table, as we travel from fluoride to iodide, the lattice enthalpy and hydration enthalpy of their sodium counterparts diminishes.

NaF > NaCl > NaBr > NaI

In general higher charges and smaller ions faour more solubility.

Solubility Product:

A solid substance dissolving in an aqueous solution has an equilibrium constant called the solubility product constant, K_sp. It shows the rate of solute dissolution in a solution. The higher a substance's K_sp value, the more soluble it is.

Consider the following general dissolving reaction in aqueous solutions:

a A + b B ⇄ c C + d D

The molarities or concentrations of the products (cC and dD) must be multiplied in order to find the K_sp. Any product that has a coefficient in front of it must be raised to the power of that coefficient (and also multiply the concentration by that coefficient). As demonstrated below:

K_sp = [C]^c [D]^d

Keep in mind that the K_sp equation does not contain the reactant, aA. Solids are not taken into account when calculating equilibrium constant expressions since their concentrations have little impact on the expression and are therefore left out. So, higher, K_sp denotes the greatest degree to which a solid can dissolve in solution.

Practice Problems:

Q1. Which of the following compounds are soluble in water?

1. C₂H₅OH
2. Benzene
3. Chloroform
4. All of these

Answer: (A)

Solution: Ethanol (C₂H₅OH) has a hydroxyl group which is highly polar and is capable of forming hydrogen bonds with polar solvents like water. The rest, C₂H₅ -group is a bulky group that does not really like to be in water but is dragged along. Benzene and chloroform are non-polar solvents which are not soluble in polar solvent water.

Q2. Which of the following best depicts the difficulties of breathing at higher altitudes?

(A) Henry's law (C) Raoult’s law

(B) Osmotic pressure (D) None of the above

Answer: (A)

Solution: According to Henry's law, the partial pressure of a gas in the vapour phase varies in direct proportion to the mole fraction of the gas in the solution. As oxygen partial pressure drops at higher elevations, its concentration also drops, making breathing more difficult.

Q3. What quality of water is responsible for its special solvent properties?

(A) Polar
(B) Flexible
(C) Cohesive
(D) Low viscosity

Answer: (A)

Solution: A universal solvent is water. More than any other solvent, it can dissolve the widest range of compounds. This unique property is explained by the water molecule's highly polar nature, which contains two highly negatively charged oxygen ions in addition to a slightly positively charged hydrogen ion. It is referred to as the "universal solvent" because it attracts a wide variety of compounds.

Q4. Toluene is more soluble in :

1. Benzene
2. Water
3. Ethyl alcohol
4. none of the above

Answer: (A)

Solution: Toluene is a nonpolar compound. So, it will be more soluble in non-polar solvents. Water and ethyl alcohol are polar compounds and benzene is a non-polar compound. So, Toluene is more soluble in benzene.

Frequently asked questions:

Q1. In water, ammonia is more soluble than oxygen. Why?

Answer: Ammonia forms hydrogen bonds with water molecules; these intermolecular bonds are extremely strong, making ammonia more soluble in water. Ammonia reacts strongly with water to form ammonium hydroxide. However, because oxygen is more electronegative, it cannot interact as much with water. As a result, ammonia (NH₃) is more soluble in water than oxygen (O₂).

Q2. What is the rate of dissolution?

Answer: A solute dissolves into a solvent during the process of dissolution, creating a solution. We know that the collisions between the solvent molecules and the particles in the solid crystal determine the dissolution of a solid by water.The rate of dissolution will rise if anything can be done to make such collisions happen more frequently or with more energy. Consider attempting to dissolve some sugar in a cup of tea. A cube of sugar would dissolve more slowly than a packet of powdered sugar. Stirring or agitating the solution would speed up the rate of dissolution. Finally, hot tea would allow the sugar to dissolve more quickly than cold tea.

Q3. How is the solubility of ionic compounds affected by the common ion effect?

Answer: The common-ion effect defines the decrease in solubility of an ionic compound caused by the addition of a salt containing an ion that already exists in chemical equilibrium to the mixture. Le-Chatelier's principle best explains this impact. Consider adding the somewhat soluble ionic substance calcium sulphate to water. The resulting chemical equilibrium's net ionic equation is as follows:

${CaSO}_{4\left(s\right)} \rightleftharpoons {Ca}_{\left(aq\right)}^{+2} + {SO}_{4\left(aq\right)}^{2-}$

Calcium sulphate is moderately soluble; at equilibrium, the majority of the calcium and sulphate exist as calcium sulphate.

Assume the solution contains the soluble ionic component copper sulphate (CuSO₄). Because copper sulphate is soluble, its main effect on the net ionic equation is the addition of additional sulphate (${SO}_4^{2-}$) ions.

${CuSO}_{4\left(s\right)} \rightleftharpoons {Cu}_{\left(aq\right)}^{2+} + {SO}_{4\left(aq\right)}^{2-}$

Because of the moderate dissociation of calcium sulphate, the sulphate ions dissociated from copper sulphate are already present (common to) in the combination. As a result, the addition of sulphate ions stresses the already established equilibrium. According to Le Chatelier's principle, more stress on the product side of the equilibrium causes the equilibrium to move to the reactants side in order to alleviate the new stress. Because of the shift toward the reactant side, the somewhat soluble calcium sulphate's solubility is reduced even further.

Q4. In terms of Henry's law, why is opening a can of soda in a low-pressure atmosphere a bad idea?

Answer: Soda's fizz is caused by dissolved CO₂, which is partially in the form of carbonic acid. The amount of CO₂ dissolved in the soda is determined by the amount of ambient pressure exerted on the liquid. As a result, the soda can will be under pressure to maintain the desired CO₂ concentration. When the can is opened to a lower pressure environment (e.g., the ambient atmosphere), the soda will swiftly "outgas" (CO₂ will come out of solution) at a rate dependent on the surrounding atmospheric pressure. If a can of soda is opened in a low pressure environment, the outgassing will be faster and thus more explosive (and dangerous) than in a high pressure environment.