Solutions and their properties

1. Solutions and their properties Part I: What’s a solution: The process of dissolving Factors affecting solubility Ways of expressing concentration Part II: Colligative properties

2. Learning objectives • Describe a solution • Describe factors influencing solubility • Predict solubility based on molecular properties • Describe effect of temperature on solubility • Predict solubility of gases using Henry’s law • Perform concentration calculations using molarity, molality, mole fraction and percent by mass and volume

3. Mixtures are a fact of life • We learn about matter by studying pure substances • In nature, most things are mixtures – nothing is “pure”: • Air • Water • Soil • Living systems

4. Solution or colloid? • Both contain a solid dispersed in a liquid • In a solution, the particle size is on the molecular scale (< 1 nm) • In a colloid, the particles are larger (2 – 500 nm) • Colloids scatter light – milk, fog

5. Why do solutions form? • Things tend to get mixed up (entropy) • Without any interactions between molecules, gases always mix • Intermolecular forces complicate matters • May improve mixing if forces in mixture are stronger • May oppose mixing if forces in pure substances are stroner • Three intermolecular forces: • Solvent – solvent • Solute – solute • Solvent – solute • Like dissolves like: • Polar solvents dissolve polar solutes • Nonpolar solvents dissolve nonpolar solutes

6. Review of intermolecular forces

7. Ion - dipole • Characteristic of interactions in solutions of ionic compounds in polar solvents • Negative ion with the positive dipole end • Positive ion with the negative dipole end

8. Dipole - dipole • Important attractive force in polar substances • Strength of the order of 3 – 4 kJ/mol (compared with 200 – 400 kJ/mol for covalent bonds)

9. Dipole force manifested in boiling points: • Nonpolar substances have much lower boiling points • Acetone (polar) 56ºC butane (nonpolar) -0.5ºC • Boiling point increases with dipole strength

10. London calling • Even molecules with no net dipole moment attract each other. • Electrons are not static but mobile: • Fluctuation creates dipole in one molecule which induces dipole in another molecule • Effect increases with atomic number – as atom becomes more polarizable • Boiling increases with molar mass • For small molecules, dispersion forces are weaker than other inter-molecular forces. For large molecules this is not true. Large molecules are solids because of dispersion forces

11. Solute – solvent interactions • The stronger the interactions between solute and solvent, the greater the solubility • Dispersion forces, dipolar interactions and liquids • Ion-dipole interactions and ionic compounds in water

12. Dissolving an ionic compound • Attractive forces between solute and solvent compensate attractive forces between solute – solute and solvent - solvent

13. Solution on the atomic scale • The ions held tightly in the crystal lattice become solvated by water molecules • Ion-ion interactions are replaced by ion-dipole interactions in the solution

14. Profits and losses • Energy costs • Solvent – solvent interactions • Solute – solute interactions (lattice energy) • High lattice energy → low solubility • Energy gain • Solvent – solute interactions • Small, more highly charged ions have stronger interactions

15. Solutions, disorder and entropy • Substances dissolve with either gain or loss of enthalpy • Dissolving usually results in an increase in entropy • Exothermic processes tend to occur spontaneously • Processes that result in increased entropy tend to occur spontaneously

16. Strong interactions complicate predictions about solubility • Weak interactions: • Gases mix completely over all compositions • Moderate interactions: • Liquids exhibit limits on range of miscibility • Strong interactions: • Solubility of ionic compounds varies • Some ionic compounds are completely insoluble

17. Super-saturate me • Saturated solution is in equilibrium with undissolved solute and cannot contain any more solute Supersaturation: a state of the solution where it contains more solute than allowed by saturation. Non-equilibrium condition • Crystallization and rainfall depend on supersaturation

18. Temperature and solubility • Solubility is amount of solute present in solution at saturation point • Solubility varies widely from compound to compound • Solubility varies with T • Most substances increase with T (KNO3) although the dependence can vary • Some are unaffected by T (NaCl) • A few decrease with T (Na2SO4) • Results can be explained by sign of ΔH solution and Le Chatelier’s principle (later)

19. Gases always decrease solubility with temperature • Solubility of gases always decreases as temperature increases

20. Solubility and pressure • For solutions of solids in liquids and liquids in liquids, pressure has almost no effect. • For gases in liquids: • Henry’s Law operates Solubility = k P • Solubility of gas depends on partial pressure above solution • k is function of gas for given solvent

21. Henry’s law and equilibrium • Amount of gas in solution is in equilibrium with gas above solution • When P is increased, density of molecules above solution increases • More molecules enter solution to restore equilibrium

22. Concentration: ways of expressing it • Molarity • Mole fraction • Mass percent/volume percent • Molality

23. Molarity (M) Molarity (M) = • Stoichiometry calculations are easy • Amounts of solution required are volumetric • Concentration varies with T • Amount of solvent requires knowledge of density

24. Mole Fraction (X) Mole fraction (X) = • Dimensionless • Independent of temperature • Used for gas mixtures and solid solutions

25. Mass percent (mass %) Mass percent = • For very dilute solutions: Parts per million (ppm) = • Independent of temperature • Need to work with masses rather than volumes

26. Molality (m) Molality (m) = • Used in calculations of colligative properties • Independent of temperature • Need to know mass of solution • Need to know density to convert to molarity