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Solutions

Solutions. Solution = a homogeneous mixture of 2 or more substances that does not scatter light Solvent = substance with same physical state as solution (or substance in greatest quantity) Solute = substance dissolved in solvent Solvents and solutes can be solid, liquid or gas

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Solutions

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  1. Solutions • Solution = a homogeneous mixture of 2 or more substances that does not scatter light • Solvent = substance with same physical state as solution (or substance in greatest quantity) • Solute = substance dissolved in solvent • Solvents and solutes can be solid, liquid or gas • In lab, solvents are most commonly liquids • Water is the main solvent in living systems • “Like dissolves like”: polar substances dissolve in polar solvents (dipole-dipole) and nonpolar substances dissolve in nonpolar solvents (dispersion)

  2. Types of solvents • Gases: gases can be mixed readily since the particles are too far apart to interact • Solids: solid solvents are usually melted so that mixing can take place in liquid phase • Liquids: solute must be “solvated” by solvent • When an ionic solid dissolves in water, the ions separate and are stabilized by surrounding water molecules • Ionic solids often have water molecules associated with them (there is water in the air), these are called hydrates (when the water is removed they’re called anhydrates)

  3. Electrolytes • Electrolyte = a compound that forms ions when dissolved so that the solution conducts electricity • Electricity is a flow of electrons • Strong electrolyte = dissolves and completely ionizes in water, strong conductor, usually an ionic compound NaCl  Na+ + Cl- • Weak electrolyte = dissolves but only partially ionizes in water, weak conductor, usually a polar covalent compound (a weak acid or weak base) CH3CO2H + H2O  CH3CO2- + H3O+ • Nonelectrolyte = dissolves but does not ionize in water, not a conductor, usually a polar covalent compound (like glucose, C6H12O6)

  4. Solubility • Solubility = amount of solute that can dissolve in a given amount of solvent at a given temperature • Usually use units of grams solute/100 g solvent • When maximum amount of solute is dissolved, solution is “saturated” (less than max is dissolved it’s “unsaturated”) • In a saturated solution, there is an equilibrium between dissolving and precipitating at the surface of any un-dissolved solid • Example: if the solubility of KCl is 34.0 g/100.0 g H2O at 20ºC, how many grams of KCl are needed to make a saturated KCl solution with 250 mL of water? (density of H2O = 1 g/mL) 250 mL H2O x (1 g/1 mL) = 250 g H2O 250 g H2O x 34.0 g KCl/100.0 g H2O = 85 g KCl

  5. Factors that Affect Solubility • Adding heat: - Increases solubility of most solids in liquids - Decreases solubility of gases in liquids (higher KE so more gas molecules escape) • Adding pressure: - Has little effect on solubility of solids in liquids - Increases solubility of gases in liquids • Henry’s Law: solubility of a gas in a liquid is dependent on the vapor pressure of the gas above the liquid

  6. Solubility of Ionic Compounds in Water • Some ionic compounds (or salts) are insoluble in water • When these form in an aqueous reaction there is a precipitate • Can predict solubility of salts based on rules (see table 9.7, p. 295): - Salts containing Li+, Na+, K+, NH4+, NO3- or CH3CO2- are always soluble - Salts containing Cl-, Br- or I- are soluble unless combined with Ag+, Pb2+ or Hg22+ - Salts containing SO42- are soluble unless combined with Ba2+, Pb2+, Ca2+ or Sr2+ -Salts containing CO32-, S2-, PO43- or OH- are insoluble unless combined with Li+, Na+, K+, NH4+, NO3- or CH3CO2- • Examples: state whether each of the following compounds are soluble in water: a) AgCl b) NaBr c) CaOH d) PbCl2 e) PbNO3 a) no b) yes c) no d) no e) yes

  7. Concentrations of Solutions • Concentration = amount of solute dissolved in a solution • Common ways to write concentration: Mass % (m/m) = (grams solute/grams solution) x 100% Volume % (v/v) = (mL solute/mL solution) x 100% Mass/volume % (m/v) = (g solute/mL solution) x 100% Molarity (M) = mol solute/L solution *Note: solution = solute + solvent • Examples: 40% (m/m) = 40 g solute in 100 g solution 25% (v/v) = 25 mL solute in 100 mL solution 10% (m/v) = 10 g solute in 100 mL solution 2.0 M = 2.0 mol solute in 1.0 L solution

  8. To make a 50.0 g solution of 16% (m/m) KCl in water: 50.0 g x (16 g KCl/100 g solution) = 8.0 g KCl

  9. To make 250.0 mL of a 2.0% (m/v) solution of KI in H2O: 250.0 mL x (2.0 g KI/100 mL solution) = 5.0 g KI

  10. Calculations using Concentration • Concentrations can be used as conversion factors • Example 1: how many moles of HCl are in 0.500 L of a 2.0 M HCl solution? Plan: L HCl  mol HCl 0.500 L x (2.0 mol HCl/1 L) = 1.0 mol HCl • Example 2: Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) How many L of a 6.0 M HCl solution must be reacted to produce 10.0 L of H2 gas at STP? L H2(g)  mol H2(g)  mol HCl(aq)  L HCl(aq) 10.0 L H2 x (1 mol H2/22.4 L H2) = 0.4464 mol H2 0.4464 mol H2 x (2 mol HCl/1 mol H2) = 0.8929 mol HCl 0.8929 mol HCl x (1 L HCl/6.0 mol HCl) = 0.15 L HCl

  11. Colloids and Suspensions • Recall: Solutions are clear and homogeneous, they have small particles dissolved that don’t scatter light • Colloids have larger particles (like proteins) that can scatter light, but they are still homogeneous (solutes don’t settle over time) • Four types of colloids: Aerosol = liquid or solid dispersed in gas (fog) Foam = gas dispersed in liquid or solid (styrofoam) Emulsion = liquid dispersed in liquid or solid (milk) Sol = solid dispersed in liquid or solid (paint) • Suspensions are heterogeneous mixtures containing large particles, often particles are large enough to see by eye - look cloudy and particles settle out over time - examples: salad dressing, medicines (like antacids)

  12. Osmosis • Osmosis = movement of water across a semi-permeable membrane from region of low solute concentration to region of high solute concentration (or high water concentration to lower water concentration) • Semi-permeable membrane allows some particles to pass through, but not others • As the volume of the side with more solute goes up, it exerts “osmotic pressure” that eventually prevents further volume change due to osmosis • Isotonic = same osmotic pressure (same # of particles) • Hypotonic = less osmotic pressure (less particles) - cell placed in hypotonic solution swells, can burst • Hypertonic = more osmotic pressure (more particles) - cell placed in hypertonic solution shrivels

  13. Dialysis • Dialysis is similar to osmosis, can be used to separate solutions from colloids • A dialysis membrane (like cellophane) allows water and small solutes (salts, glucose) through, but not larger particles (proteins, DNA) • Example: if you fill a dialysis bag with a solution containing DNA, starch, NaCl and sucrose, and place it in a beaker of pure water, what happens? NaCl and sucrose will flow out of the bag, and water will flow in, until the concentrations are the same inside and outside the bag. The DNA and starch will remain inside the bag.

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