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Mixtures: Solutions and Colloids

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Mixtures: Solutions and Colloids Mixtures: Solutions and Colloids Intermolecular Forces in Solution – Solutions – homogeneous mixtures (a single phase) • All types of IFs in pure substances also occur in – Colloids – heterogeneous mixtures (two or more phases) solutions (in addition, ion-dipole forces are very common) 13.1 Types of Solutions and Solubility • Ion-dipole forces – present in solutions of ionic – Solvent – the substance that dissolves; usually compounds in polar solvents such as H2O the most abundant component of the mixture; has – Hydration (solvation) the same physical state as the solution • The water dipoles pull the ions – Solute – the substance that is dissolved away from the ionic crystal and surround them (typical coord. – Solubility (S) – the maximum amount of solute numbers 4 or 6) that can be dissolved in a given amount of • There is a short range order solvent or solution around the ions consisting of H- – Concentration of solute – various units are used bonded shells of water molecules • Dipole-dipole forces – present in solutions of •The “like dissolves like rule” – substances with similar types of IFs dissolve in each other well polar molecules in polar solvents such as H2O • H-bonding forces – present in solutions of O- – Strong solute-solvent IFs lead to better solubility and N-containing molecules (sugars, alcohols, (lower the total energy of the system) amino acids, …) in protic solvents such as H O – The solute-solvent IFs created during dissolution 2 must have comparable strength to the solute-solute • Ion-induced dipole forces – solutions of ionic and solvent-solvent IFs destroyed in this process compounds in less polar or non-polar solvents • Dipole-induced dipole forces – present in solutions of polar molecules in non-polar solvents or non-polar molecules in polar solvents • Dispersion forces – present in all solutions (most important for solutions of non-polar molecules in non-polar solvents) Example: Example: ¾ Water dissolves well alcohols (ROH) with short What is a better solvent for diethyl ether hydrocarbon chains (R) → methanol, ethanol, … (CH CH -O-CH CH ), water or propanol ¾ Strong H-bonding IFs in the solute and the solvent are 3 2 2 3 replaced with strong solute-solvent H-bonding IFs (CH3CH2CH2OH)? ¾ Water does not dissolve well alcohols (ROH) with ¾ Both water and propanol interact with ether long hydrocarbon chains (R) → hexanol, … through H-bonding ¾ Strong H-bonding in water is replaced with weaker IFs ¾ Propanol and ether interact well through dispersion between the water dipole and the large non-polar forces (similar non-polar hydrocarbon chains) hydrocarbon chains (R) → (H-bonding with the OH part ¾ Water can’t interact well through dispersion forces of the alcohol is a smaller fraction of the total IFs) with the hydrocarbon portion of ether ⇒ Propanol is a better solvent for ether 1 • Soaps –Na+ salts of long chain carboxylic acids • Liquid solutions – the solvent is a liquid – Long, non-polar hydrocarbon “tail” → hydrophobic – Solid-liquid and liquid-liquid solutions – Small, polar-ionic carboxyl “head” → hydrophilic • Some salts and polar molecular compounds dissolve – The polar head dissolves in water; the non-polar tail well in water and short chain alcohols dissolves in grease → washing action • Less polar solids dissolve well in less polar solvents like acetone, chloroform, ether, … • Detergents – contain surfactants (surface-active • Non-polar substances dissolve best in non-polar compounds) → similar to soaps solvents like hexane, benzene, … Tail Head – Gas-liquid solutions • Non-polar gases have poor solubility in water • Some gases dissolve in water through chemical reactions (HCl, CO2, SO2, …) • Gaseous solutions – the solvent is a gas – Gas-gas solutions – gases mix in all proportions • Solid solutions – the solvent is a solid Example: – Gas-solid solutions – gas molecules penetrate the Quick cleaning of laboratory glassware: crystal lattices of some metals (H2/Pd, O2/Cu, …) ¾ Cleaning a sample tube with a salt residue – Solid-solid solutions – homogeneous alloys, Wash with water → dissolves salt (ion-dipole forces) waxes, … Wash with ethanol → dissolves water (H-bonding) ¾Substitutional alloys – Wash with acetone → dissolves ethanol (dipole- the atoms of one element dipole, dispersion forces) take some of the positions Dry → acetone evaporates easily (low T ) in the lattice of another b element ¾ Cleaning a sample tube with an oily residue ¾Interstitial alloys – the Wash with hexane → dissolves oil (dispersion forces) atoms of one element fit Wash with acetone → dissolves hexane (dispersion into the gaps of the lattice and dipole-induced dipole forces) of another element Dry → acetone evaporates easily (low Tb) 13.2 Energy Changes in the Solution Solute (aggregated) + heat → Solute (separated) ∆Hsolute>0 Solvent (aggregated) + heat → Solvent (separated) ∆H >0 Process solvent Solute (separ.) + Solvent (separ.) → Solution + heat ∆Hmix<0 • Solution cycle – the solution process can be – Heat of solution (∆Hsoln) – the total enthalpy divided into three steps: change during the solution cycle (can be either exothermic or endothermic) 1) Separation of solute particles to make room for the solvent – endothermic (energy is needed to overcome the IFs of attraction) 2) Separation of solvent particles to make room for the solute particles – endothermic (energy is needed to overcome the IFs of attraction) 3) Mixing of solvent and solute particles – exothermic (solute-solvent IFs lower the energy) ∆Hsoln = ∆Hsolute + ∆Hsolvent + ∆Hmix 2 Solutions of Ionic Solids • Individual ionic heats of hydration – ∆H for • For dilute solutions of ionic solids, the solute the hydration of 1 mol of separated gaseous separation has ∆H equal to the lattice enthalpy cations (or anions) + - M+(g) → M+(aq) or X-(g) → X-(aq) MX(s) → M (g) + X (g) ∆Hsolute = ∆Hlattice> 0 •The hydration of the separated solute ions is a – ∆Hhydr of ions is always exothermic (∆Hhydr< 0) combination of steps 2 and 3 in the solution cycle because the ion-dipole forces that replace some of → (solvent separation + mixing) the H-bonds in water are stronger – The overall heat of hydration is a sum of the heats – Heat (enthalpy) of hydration (∆H ) – ∆H for hydr for the cations and the anions hydration of separated gaseous ions M+(g) + X-(g) → M+(aq) + X-(aq) ¾Trends in ∆Hhydr of ions (same as for ∆Hlattice) • ∆H is larger (more exothermic) for ions with ∆Hhydr = ∆Hsolvent + ∆Hmix hydr ⇒∆H = ∆H + ∆H greater charges and smaller sizes → ions with soln solute hydr higher charge density ⇒∆Hsoln= ∆Hlattice + ∆Hhydr • The charge factor is more important Some ionic heats of hydration in kJ/mol The Tendency Toward Disorder Li+(-558) Be2+(-2533) F-(-483) – ∆Hsoln< 0 favors the solution process since the Na+(-444) Mg2+(-2003) Al3+(-4704) Cl-(-340) total energy of the system is lowered; but many K+(-361) Ca2+(-1657) Br-(-309) ionic solids have ∆Hsoln> 0 and still dissolve ¾ The sign of ∆Hsoln is hard to predict, since both readily in water ∆Hlattice and ∆Hhydr depend on the ionic charge and ⇒There is a second factor that affects the solution size and tend to cancel each other’s effects process and acts in addition to the enthalpy factor Example: NaCl • Disorder – systems have a natural tendency to ∆H = 787 kJ/mol lattice become more disordered ∆Hhydr = (-444) + (-340) = -784 kJ/mol – Entropy – a measure of the disorder in the system ↑Entropy ↔↑Disorder ∆Hsoln = (787) + (-784) = +4 kJ/mol – Mixing leads to greater disorder and ↑ the entropy • The solution process is governed by the 13.3 Solubility as an Equilibrium Process combination of the enthalpy and entropy factors • Typically solutes have limited solubility in a – Compounds with ∆Hsoln< 0 are typically soluble since both factors are favorable given solvent – The solution process is countered by – Compounds with ∆Hsoln> 0 are soluble only if ∆Hsoln is relatively small so the favorable entropy factor recrystallization of the dissolved solute dominates – Dynamic equilibrium – the rates of dissolution – Compounds with ∆Hsoln>> 0 are typically insoluble and recrystallization become equal since the unfavorable enthalpy factor dominates Solute (undissolved) ↔ Solute (dissolved) Example: Benzene is insoluble in water because – Saturated solution – no more solute can be the solute-solvent IFs are weak so the (-) ∆Hmix dissolved (can be produced by equilibrating the is very small compared to the (+) ∆Hsolut & ∆Hsolv solution with an excess of the solute; the excess ⇒ ∆Hsoln= ∆Hsolute + ∆Hsolvent + ∆Hmix >> 0 solute can be filtered out after saturation) 3 – Molar solubility (S) – the concentration of the Solute + Solvent + heat ↔ Saturated solution saturated solution in mol/L ¾Increasing T provides the heat needed for the – Unsaturated solution – more solute can be forward reaction to proceed (S↑) dissolved – Supersaturated solution – the concentration is •Some solids are more soluble at higher T higher than the solubility, S (unstable, non- despite their exothermic (-) ∆Hsoln (Why?) equilibrium state; disturbances or addition of seed ¾Tabulated ∆Hsoln values refer to the standard state crystal lead to crystallization) (1M solution), but saturated solutions can be Effect of Temperature on Solubility much more concentrated •Most solids are more soluble at higher T due ¾At very high concentrations the hydration process is hindered so ∆H becomes less negative; thus to their endothermic (+) ∆Hsoln hydr ∆Hsoln can change dramatically and even switch ¾(+) ∆Hsoln means that heat is absorbed during dissolution (the heat
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