What Does The Phrase Like Dissolves Like Mean

The principle of "like dissolves like" is a guiding rule in chemistry that explains why certain substances mix well, while others don't. It's a simple, yet powerful concept that dictates the solubility of one substance (solute) in another (solvent). This article delves deep into the meaning of this phrase, exploring the underlying principles, practical applications, and the exceptions that make chemistry a fascinating field of study.

Understanding "Like Dissolves Like": A Deep Dive

At its core, "like dissolves like" suggests that substances with similar intermolecular forces are more likely to dissolve in each other. This similarity refers primarily to the polarity of the molecules involved. Polar solvents tend to dissolve polar solutes, while nonpolar solvents tend to dissolve nonpolar solutes. To fully grasp this concept, we must first understand what polarity and intermolecular forces entail.

Polarity: The Uneven Distribution of Electrons

Polarity arises when there is an unequal sharing of electrons within a molecule. This uneven distribution creates a partial positive charge (δ+) on one part of the molecule and a partial negative charge (δ-) on another. This occurs when atoms with differing electronegativities form a bond. Electronegativity is the ability of an atom to attract electrons in a chemical bond.

For instance, in a water molecule (H2O), oxygen is more electronegative than hydrogen. This means oxygen pulls the shared electrons closer to itself, resulting in a partial negative charge on the oxygen atom and partial positive charges on the hydrogen atoms. This makes water a polar molecule.

Intermolecular Forces: The Glue That Holds It Together

Intermolecular forces (IMFs) are attractive or repulsive forces that exist between molecules. These forces are weaker than the intramolecular forces (e.g., covalent bonds) that hold atoms together within a molecule, but they play a crucial role in determining the physical properties of substances, including their solubility. There are several types of IMFs, each with varying strengths:

• Hydrogen Bonding: This is the strongest type of IMF and occurs when a hydrogen atom is bonded to a highly electronegative atom like oxygen (O), nitrogen (N), or fluorine (F). The partially positive hydrogen is attracted to the lone pair of electrons on another electronegative atom. Water exhibits strong hydrogen bonding.

• Dipole-Dipole Interactions: These forces occur between polar molecules. The positive end of one polar molecule is attracted to the negative end of another. These interactions are stronger than London dispersion forces but weaker than hydrogen bonds.

• London Dispersion Forces (LDF): These are the weakest type of IMF and are present in all molecules, both polar and nonpolar. LDF arise from temporary, instantaneous fluctuations in electron distribution, creating temporary dipoles. Larger molecules with more electrons tend to have stronger LDF.

The Dissolution Process: A Molecular Dance

When a solute dissolves in a solvent, the process involves breaking the intermolecular forces holding the solute together and the intermolecular forces holding the solvent together. New intermolecular forces then form between the solute and solvent molecules.

For dissolution to occur favorably, the new solute-solvent interactions must be comparable in strength to the solute-solute and solvent-solvent interactions that were broken. This is where the "like dissolves like" principle comes into play.

• Polar Solvents and Polar Solutes: Polar solvents, like water, have strong dipole-dipole interactions and/or hydrogen bonding. They readily dissolve polar solutes because the solute-solvent interactions (e.g., dipole-dipole, hydrogen bonding) are strong enough to overcome the solute-solute and solvent-solvent interactions. A classic example is sugar dissolving in water. Sugar molecules are polar and can form hydrogen bonds with water molecules, facilitating the dissolution process.

• Nonpolar Solvents and Nonpolar Solutes: Nonpolar solvents, like hexane or benzene, primarily exhibit London dispersion forces. They readily dissolve nonpolar solutes because the solute-solvent interactions (LDF) are comparable in strength to the solute-solute and solvent-solvent interactions. For example, oil dissolves well in hexane. Both oil and hexane are nonpolar and interact through LDF.

• Polar Solvents and Nonpolar Solutes (and vice versa): When trying to dissolve a polar solute in a nonpolar solvent (or vice versa), the interactions are mismatched. The strong intermolecular forces in the polar substance cannot be adequately compensated by the weak intermolecular forces in the nonpolar substance. As a result, the solute-solvent interactions are too weak to overcome the solute-solute and solvent-solvent interactions, and dissolution does not occur to a significant extent. For example, oil and water do not mix. The strong hydrogen bonds between water molecules are far stronger than the weak London dispersion forces that could form between water and oil molecules.

Examples Illustrating "Like Dissolves Like"

To solidify the understanding of this principle, let's explore some specific examples:

• Salt (NaCl) in Water (H2O): Salt is an ionic compound and is highly polar. Water, as mentioned earlier, is also highly polar due to its bent shape and the electronegativity difference between oxygen and hydrogen. When salt is added to water, the partially negative oxygen atoms in water are attracted to the positive sodium ions (Na+), and the partially positive hydrogen atoms are attracted to the negative chloride ions (Cl-). These ion-dipole interactions are strong enough to break the ionic bonds in the salt crystal lattice, allowing the ions to disperse throughout the water.

• Iodine (I2) in Carbon Tetrachloride (CCl4): Iodine is a nonpolar molecule, as the two iodine atoms share electrons equally. Carbon tetrachloride is also nonpolar because, although the C-Cl bonds are polar, the symmetrical tetrahedral shape of the molecule cancels out the individual bond dipoles, resulting in a zero net dipole moment. Iodine dissolves readily in carbon tetrachloride because both substances interact through London dispersion forces.

• Ethanol (C2H5OH) in Water (H2O): Ethanol is an alcohol that contains both a polar hydroxyl (-OH) group and a nonpolar ethyl (C2H5) group. The -OH group allows ethanol to form hydrogen bonds with water, while the ethyl group can interact through London dispersion forces. This dual nature makes ethanol miscible (mixable in all proportions) with water.

• Sugar (C12H22O11) in Hexane (C6H14): Sugar is a polar molecule with numerous -OH groups that can form hydrogen bonds. Hexane, on the other hand, is a nonpolar solvent with only London dispersion forces. Sugar does not dissolve in hexane because the strong hydrogen bonds between sugar molecules cannot be overcome by the weak London dispersion forces between sugar and hexane molecules.

Applications of "Like Dissolves Like"

The "like dissolves like" principle has numerous practical applications in various fields:

• Cleaning: Understanding solubility helps in selecting appropriate cleaning agents. For example, greasy stains (nonpolar) are best removed with nonpolar solvents like turpentine or mineral spirits, while sugary spills (polar) are easily cleaned with water (polar). Soaps and detergents work because they have both polar and nonpolar regions, allowing them to emulsify nonpolar grease in polar water.

• Pharmaceuticals: Drug delivery often depends on the solubility of the drug in bodily fluids. Drugs need to be formulated in a way that ensures they dissolve in the appropriate environment (e.g., the stomach, bloodstream) to be effectively absorbed. The "like dissolves like" principle guides the selection of suitable solvents and excipients (inactive ingredients) in drug formulations.

• Chemical Reactions: Solubility plays a crucial role in chemical reactions. Reactants must be soluble in a given solvent for the reaction to occur efficiently. Choosing the right solvent based on the polarity of the reactants is essential for optimizing reaction rates and yields.

• Environmental Science: Understanding solubility is important for predicting the fate and transport of pollutants in the environment. For example, nonpolar pollutants like oil spills tend to persist in the environment because they do not dissolve readily in water and can contaminate soil and groundwater.

• Chromatography: Chromatography is a separation technique that relies on the differential solubility of compounds between two phases: a stationary phase and a mobile phase. The "like dissolves like" principle is used to select the appropriate stationary and mobile phases to achieve effective separation of the components in a mixture. For example, in reverse-phase chromatography, a nonpolar stationary phase is used to retain nonpolar compounds, which are then eluted using a more polar mobile phase.

Exceptions and Limitations

While the "like dissolves like" principle is a valuable guideline, it's not without its exceptions and limitations:

• Amphipathic Molecules: Molecules with both polar and nonpolar regions, like soaps and detergents, can dissolve in both polar and nonpolar solvents to some extent. This is because the polar region interacts with polar solvents, while the nonpolar region interacts with nonpolar solvents.

• Temperature: Temperature can significantly affect solubility. In general, the solubility of solids in liquids increases with increasing temperature, while the solubility of gases in liquids decreases with increasing temperature. This is because increasing temperature provides more energy to overcome the intermolecular forces holding the solute together.

• Pressure: Pressure has a significant effect on the solubility of gases in liquids. According to Henry's Law, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid. This means that increasing the pressure of a gas will increase its solubility in a liquid.

• Complex Formation: In some cases, substances that are normally insoluble can dissolve in a solvent if they form a complex with another substance that is soluble. For example, iodine (I2) is only sparingly soluble in water, but it dissolves readily in a solution of potassium iodide (KI) due to the formation of the triiodide ion (I3-), which is soluble.

• Salting Out: Adding a salt to a solution of a polar organic compound in water can sometimes decrease the solubility of the organic compound. This phenomenon, known as "salting out," occurs because the salt ions compete with the organic compound for interactions with water molecules, effectively reducing the number of water molecules available to solvate the organic compound.

The Scientific Basis: Thermodynamics and Entropy

The dissolution process is governed by thermodynamic principles, particularly the Gibbs free energy change (ΔG), which is defined as:

ΔG = ΔH - TΔS

where:

• ΔG is the change in Gibbs free energy
• ΔH is the change in enthalpy (heat absorbed or released during the process)
• T is the absolute temperature
• ΔS is the change in entropy (a measure of disorder or randomness)

For a dissolution process to be spontaneous (i.e., favorable), ΔG must be negative.

• Enthalpy (ΔH): The enthalpy change is related to the energy required to break the solute-solute and solvent-solvent interactions and the energy released when new solute-solvent interactions are formed. If the solute-solvent interactions are stronger than the solute-solute and solvent-solvent interactions, ΔH will be negative (exothermic), favoring dissolution. If the solute-solvent interactions are weaker, ΔH will be positive (endothermic), disfavoring dissolution.

• Entropy (ΔS): The entropy change is related to the increase in disorder when the solute and solvent molecules mix. In general, the dissolution process leads to an increase in entropy (ΔS is positive) because the solute and solvent molecules are more randomly distributed in the solution than in the pure substances. The increase in entropy favors dissolution.

The "like dissolves like" principle can be understood in terms of these thermodynamic parameters. When the solute and solvent have similar intermolecular forces, the enthalpy change (ΔH) is relatively small because the energy required to break the solute-solute and solvent-solvent interactions is roughly equal to the energy released when new solute-solvent interactions are formed. In this case, the entropy change (ΔS) becomes the dominant factor, and the dissolution process is favored because it leads to an increase in disorder.

However, when the solute and solvent have very different intermolecular forces, the enthalpy change (ΔH) can be large and positive, disfavoring dissolution. For example, when trying to dissolve a nonpolar solute in water, the strong hydrogen bonds between water molecules must be broken, requiring a significant amount of energy. The weak London dispersion forces that could form between the nonpolar solute and water molecules are not strong enough to compensate for the energy required to break the hydrogen bonds. In this case, the enthalpy change (ΔH) outweighs the entropy change (ΔS), and the dissolution process is not favored.

FAQ: "Like Dissolves Like"

• Q: Is "like dissolves like" always true?

  • A: No, it's a general guideline with exceptions. Factors like temperature, pressure, and the presence of other substances can influence solubility. Amphipathic molecules also exhibit behavior that deviates from the simple rule.

• Q: What happens if I try to mix a polar and a nonpolar substance?

  • A: They will likely not mix well. They may form separate layers, with the denser substance settling at the bottom.

• Q: Why is water considered the "universal solvent"?

  • A: Water is an excellent solvent for many polar and ionic compounds due to its polarity and ability to form hydrogen bonds. However, it is not a universal solvent as it does not dissolve nonpolar substances well.

• Q: How does temperature affect solubility?

  • A: Generally, the solubility of solids in liquids increases with increasing temperature, while the solubility of gases in liquids decreases with increasing temperature.

• Q: Can I use "like dissolves like" to predict the solubility of complex molecules?

  • A: Yes, but it requires careful consideration of the different functional groups within the molecule and their relative polarities.

• Q: What are some real-world examples of "like dissolves like"?

  • A: Removing grease with oil-based solvents, dissolving sugar in water, and formulating drugs for effective absorption are all examples of this principle in action.

Conclusion

The "like dissolves like" principle is a cornerstone of understanding solubility in chemistry. It highlights the importance of intermolecular forces and polarity in determining whether one substance will dissolve in another. By understanding this principle, we can predict the solubility of substances, select appropriate solvents for various applications, and gain insights into the behavior of molecules in solutions. While there are exceptions and limitations, the "like dissolves like" principle remains a valuable tool for chemists and scientists in various fields. From cleaning greasy stains to formulating life-saving drugs, this simple phrase encapsulates a fundamental concept that governs the world around us.