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Understanding Solutions: Homogeneous Mixtures and Solvation

Learn about solutions, the formation of homogeneous mixtures consisting of a solvent and solutes, and the process of solvation. Explore factors affecting dissolution and the solubility of gases and liquids, as well as the concepts of molality and mole fraction.

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Understanding Solutions: Homogeneous Mixtures and Solvation

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  1. Solutions • Solution - A homogeneous mixture of substances that consists of a solvent and one or more solutes. • Solvent - Generally the most abundant substance in a solution. • Solute - The impurity in a solution and less abundant than the solvent. • Dissolution • Dissolution - The formation of a solution by the addition of a solute to a solvent. This may occur with or without reaction. • Solvation - Dissolution without the occurrence of a reaction. • 2Na(s) + 2H2O  2NaOH(aq) + H2(g) • H2O • NaCl(s)  NaCl(aq)

  2. Ease of Dissolution • Two factors that affect the ease of dissolution: • 1) The change in energy (enthalpy) • 2) The change in order (entropy) • Dissolution is favored by exothermic processes that increase disorder • Interactions that favor dissolution • 1) weak solute-solute intermolecular interactions • 2) weak solvent-solvent intermolecular interactions • 3) strong solvent-solute intermolecular interactions

  3. (page 544) • Enthalpy of solution - Hsol - The heat involved in the solvation of a solute. • Generally dissolution accompanies an increase in disorder so, endothermic dissolutions can occur spontaneous depending on the magnitude of Hsol.

  4. Solvation of solids • Crystal lattice energy (solute-solute) - The energy change involved in the formation of one mole of formula units in a crystalline state from it’s particles in the gaseous state. (always exothermic) • Na+(g) + Cl-(g) NaCl(s) + heat • Hsol = (heat of solvation) - (crystal lattice energy) • Hydration - Solvation where the solvent is water. • The larger the ion the more waters hydrate it (4-9 waters, 6 average)

  5. Dissolution of liquids • miscibility - the ability of one liquid to dissolve in another. • Polar liquids tend to dissolve in polar liquids • Non-polar liquids tend to dissolve in non-polar liquids • The dissolution of liquids is generally exothermic

  6. Dissolution of gases in liquids • Gases that are capable of hydrogen bonding or ionize are generally soluble in water (HF, HCl, HBr ...) • Some non-polar gases are slightly soluble in water • O2 due to dispersion forces between O2 and H2O • CO2 due to reaction CO2(g) + H2O  H2CO3 • H2CO3  HCO3- + H+ • HCO3- CO32- + H+ • An increase in pressure increases • the solubility of a gas

  7. Effect of temp. on solubility • Exothermic: reactant  product + heat • Endothermic: reactant + heat  product • negative Hsol (exothermic) solubility decreases with increasing temp. • positive Hsol (endothermic) solubility increases with increasing temp. • gas dissolutions in liquid are exothermic • soda goes flat at room temperature faster than in the refrigerator • (Don’t put your fish tank near a window... )

  8. Molality • molality - moles of solute per 1 kg of solvent. • m = mol of solute/kg of solvent (mol/kg) • Why molality instead of molarity? • Molality is used when there is a change in temperature involved because it is temperature invarient. • Molarity (moles/L) is dependent on the volume which changes with a change in temperature. • What is the molality of a solution prepared from 50 g of sucrose (C12H22O11) and 117g of water? • m = (50g sucrose * 1mol/342g sucrose)/0.117 kg H2O = 1.25 m

  9. Mole Fraction • Mole Fraction of A - XA = moles of A/total moles • dimensionless quantity • What is the mole fraction of CH3OH and H2O in a solution composed of 128 g CH3OH and 108 g H2O? • mol CH3OH = 128gCH3OH * (1 mol/32.0gCH3OH) = 4.00 mol CH3OH • mol H2O = 108 g H2O * (1 mol/18.0 g H2O) = 6.00 mol H2O • XCH3OH= 4.00 mol/(4.00 mol + 6.00 mol) = 0.400 • XH2O = 6.00 mol/(4.00 mol + 6.00 mol) = 0.600 • The mole fraction of all species in a solution add up to 1

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