8.2 - Factors That Affect Solubility and Rate of Dissolving

8.2 Factors That Affect Solubility and Rate of Dissolving

Solubility and Forces Between Particles

• Three Categories of Forces: The formation of solutions depends on the relative strength of three types of forces:

  1. Forces that attract solute particles to each other.

  2. Forces that attract solute particles to solvent particles.

  3. Forces that attract solvent particles to each other.

• Solution Formation: When a solution forms, the intermolecular forces between solute particles are broken, as are some of the intermolecular forces between solvent particles.

• Gaseous Solutions: Gases mix readily to form solutions because the forces of attraction between gas particles are negligible.

Solubility in Water

• Water as a Solvent: Water is a good solvent due to its polar nature, allowing it to dissolve a wide range of solutes.

• Polarity of Water Molecules:

  • Oxygen-Hydrogen Bonds: Water molecules have polar oxygen-hydrogen bonds and a V-shape (Creating a dipole).

  • Electronegativity: Oxygen is more electronegative than hydrogen, causing electrons to be displaced toward the oxygen atom.

  • Partial Charges: This displacement gives oxygen a small net negative charge and hydrogen atoms a net positive charge.

  • Dipoles: Both O-H bonds are polar and act as dipoles.

• Hydrogen Bonding: There is a strong attraction between the oxygen atom of one water molecule and the hydrogen atoms of adjacent water molecules, which is an example of hydrogen bonding.

• Definition of Hydrogen Bonding: Hydrogen bonding is a type of dipole-dipole attraction between molecules.

• Hydrogen Bond Formation: A hydrogen bond can form between a hydrogen atom on one molecule and a highly electronegative atom (oxygen, fluorine, or nitrogen) on another molecule.

Solubility of Ionic Compounds in Water

• Solubility of Ionic Compounds: Most ionic compounds are soluble in water because the attraction between the ions and the dipoles on water molecules is strong enough to pull ions away from the ionic compound.

• Dissolving Process:

  • Hydration: Water molecules surround each ion in a process called hydration.

  • Separation and Dispersion: This process causes ions to separate and disperse throughout the water.

• Insoluble Ionic Compounds: Some ionic compounds do not dissolve in water because the attraction between the ions in the compound is greater than the attraction between the ions and water dipoles.

• Chemical Formula Representation: The chemical formula of an ionic compound in solution represents the overall composition of the solute, not solute ions bonded to each other. For example, $NaCl(aq)$ indicates equal numbers of sodium and chlorine ions that are separated and evenly distributed.

Solubility of Molecular Compounds in Water

• Dipole-Dipole Attraction: Dipole-dipole attraction can occur between molecules of any polar compound, but it is weaker than the attraction between ions in an ionic compound.

• Crystal Formation: Dipole-dipole attraction can hold polar molecules together as a solid crystal, such as sucrose.

Solubility of Polar Compounds in Water

• Solubility of Polar Compounds: Most polar compounds dissolve in water because the dipole-dipole attraction between polar solute molecules is generally weaker than the hydrogen bonds between solute and water molecules.

• Sucrose Solubility: Sucrose ($C<em>{12}H</em>{22}O_{11(s)}$) is very soluble in water because it has eight polar -OH groups that can form hydrogen bonds with water molecules. The attraction between sucrose and water is stronger than the attraction between sucrose molecules themselves.

• Non-polar Molecules: Non-polar molecules do not dissolve in water because they are only weakly attracted to water molecules, and this attraction is too weak to break the hydrogen bonds between water molecules.

Conductivity of Aqueous Solutions

• Conductivity Tests: Conductivity tests can indicate whether a compound dissolved in water is ionic or molecular.

• Ionic Compounds: Ionic compounds dissociate into separate ions when dissolved in water, and these ions can carry charge to electrodes, conducting electric current.

• Molecular Compounds: Most molecular compounds remain intact and neutral when dissolved, so they do not conduct electricity well.

Predicting Whether an Ionic Compound Is Soluble in Water

• Factors Affecting Solubility: The solubility of an ionic compound depends on how easily water can separate the solute's particles. A stronger attraction between cations and anions makes the substance less soluble.

• Two Main Factors:

  1. Amount of charge on each ion.

  2. Size of each ion.

• Effect of Ion Charge: A Greater charge on each ion results in a less soluble compound due to a stronger force of attraction. For example, oxide ions ($O^{2-}$ ) have twice the charge of fluoride ions ($F^{-}$$), making oxides less soluble.

• Effect of Ion Size: Larger ion size generally increases solubility because the force of attraction between opposite charges decreases as the distance between charges increases. For example, silver nitrate ($AgNO<em>3(s)$) is very soluble, while silver chloride ($AgCl(s)$) is insoluble because the $NO</em>3^-$ ion is much larger than the $Cl^-$ ion.

Solubility Guidelines for Ionic Compounds

• General Rule: An ion with a small charge and a large radius typically forms a soluble ionic compound.

• Solubility Guidelines Summary:

  1. Hydrogen, ammonium, and Group 1 (alkali metal) ions form soluble compounds with nearly all anions.

  2. Nitrate and acetate ions form soluble compounds with nearly all cations.

  3. Chloride, bromide, and iodide ions form compounds with low solubility when combined with silver, lead(II), mercury(I), copper(I), and thallium cations.

  4. Fluoride ions form compounds with low solubility when combined with magnesium, calcium, barium, and lead(II) cations.

  5. Sulfate ions form compounds with low solubility when combined with calcium, strontium, barium, and lead(II) cations.

  6. Sulfide ions form soluble compounds only with the ions listed in guideline 1 and the Group 2 cations.

  7. Hydroxide ions form soluble compounds only with the cations listed in guideline 1, as well as strontium, barium, and thallium cations.

  8. Phosphate, carbonate, and sulfite ions form compounds with low solubility with all cations except those listed in guideline 1.

Predicting Whether a Molecular Compound Is Soluble in Water

• Size and Solubility: The size of a molecular compound affects its solubility. Methanol ($CH<em>3OH(l)$) and ethanol ($CH</em>3CH_2OH(l)$) are completely soluble in water because of hydrogen bonds between water and the -OH group.

• Non-polar Portion Effect: The non-polar part of the molecule affects solubility, with smaller non-polar portions resulting in greater solubility.

"Like Dissolves Like"

• Polar vs. Non-polar: Non-polar molecules are only weakly attracted to water molecules and are generally insoluble in water. However, non-polar solvents can dissolve non-polar molecules.

• Generalization: Polar solvents tend to dissolve ionic and polar molecules, while non-polar solvents tend to dissolve non-polar molecules. This is summarized as "like dissolves like."

"Like Dissolves Like" in the Solubility of Gases

• Gas Solubility: Hydrogen chloride ($HCl(g)$) and ammonia ($NH<em>3(g)$) are soluble in water because they are polar molecules. Carbon dioxide ($CO</em>2(g)$) is relatively insoluble because it is non-polar.

Molecules That Have Both Polar and Non-Polar Components

• Amphipathic Molecules: Many molecules contain both polar and non-polar components, such as acetic acid ($CH<em>3COOH(l)$), which has a polar O-H bond and a non-polar -$CH</em>3$ group.

• Solubility in Different Solvents: Acetic acid is soluble in both water and non-polar solvents like benzene ($C<em>6H</em>6(l)$) and carbon tetrachloride ($CCl_4(l)$).

• Soaps and Detergents: Soap and detergent molecules also have both polar and non-polar bonds, allowing them to dissolve in both water and oil/grease.

Temperature and Solubility

• Effect of Temperature: Changes in temperature can affect solubility, with higher temperatures resulting in more frequent and energetic collisions between solvent and solute particles.

• Standard Conditions: Solubilities are often stated at standard ambient temperature and pressure (SATP), which is 25°C and 100 kPa.

Effect of Temperature on the Solubility of Solids

• General Trend: Solubility of solid solutes in liquid solvents usually increases with temperature. For example, caffeine solubility in water increases from 2.2 g/100 mL at 25°C to 40 g/100 mL at 100°C.

• Solubility Curves: A solubility curve is a graph that shows the relationship between the solubility of a solute and the temperature of the solvent.

Effect of Temperature on the Solubility of Liquids and Gases

• Liquids in Liquids and Gases in Gases: Temperature usually has little effect on the solubility of one liquid in another liquid or one gas in another gas.

• Gases in Liquids: The solubility of a gas in a liquid depends on both temperature and pressure. Increasing the temperature of a solution containing dissolved gas molecules provides energy for the gas molecules to escape, decreasing solubility.

• Carbonated Drinks: The decrease in solubility explains why carbonated drinks go "flat" more quickly at room temperature. Carbon dioxide reacts with water to form carbonic acid: $H<em>2O(l) + CO</em>2(g) \rightleftharpoons H<em>2CO</em>3(aq)$.

Environmental Effects of Increased Temperatures

• Thermal Pollution: Heat (or thermal) pollution and the warming of Earth can have serious consequences due to decreased gas solubility in water.

• Manufacturing Processes: Many manufacturing and electricity-generating plants release waste heat into the environment, often by pumping water from a nearby source through a heat-exchange system and then returning the heated water.

• Ecological Impact: Warmer water decreases oxygen solubility, making aquatic organisms more susceptible to disease and potentially causing death from oxygen shortage.

Pressure and Solubility

• Effect on Liquids and Solids: Pressure has very little effect on the solubility of a liquid or a solid.

• Effect on Gases: The solubility of a gas in a liquid is directly proportional to the pressure of that particular gas above the liquid (Henry's Law). It is not affected by the pressure of other gases.

• Carbonated Drinks: The pressure of carbon dioxide in a sealed soft drink is much higher than atmospheric pressure, which greatly increases its solubility in water.

Pressure, Solubility, and Scuba Diving

• Scuba Diving and Gas Solubility: The solubility of gases is vital to scuba divers who breathe compressed air. Deeper dives mean greater water pressure, and the diver breathes air at a greater pressure.

• Nitrogen Dissolution: More gases, especially nitrogen, dissolve in the diver's blood at higher pressures.

• Decompression Sickness (The Bends): Divers must surface slowly to allow dissolved nitrogen to come out of solution gradually. Surfacing too quickly can cause bubbles to form in the blood, leading to a painful and dangerous condition known as the "bends."

• Flying After Diving: Divers should not fly within 24 hours of completing a dive because air pressure in airplanes is lower than sea-level air pressure, increasing the risk of the bends.

Factors That Affect the Rate of Dissolving

• Definition: The rate of dissolving measures how quickly a solute dissolves in a solvent, distinct from solubility, which measures the amount of solute that dissolves in a given volume of solvent.

• Importance: Rate of dissolving is important in various applications (cooking, labs, engineering).

• Factors Affecting Rate:

  • The rate of dissolving depends on the collisions between solute and solvent particles.

    1. Agitation or Mixing: Increases collisions between solute and solvent particles.

    2. Temperature: Higher temperatures increase the kinetic energy of solvent particles, leading to more frequent collisions.

    3. Surface Area: A greater surface area of the solute increases the amount of solute in direct contact with the solvent, increasing the rate of collisions.