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  1. Solutions Unit Honors Chemistry

  2. Solution • Definition: a homogeneous mixture of 2 or more substances in a single physical state • Parts: solute and solvent (usually water) • Types: • Physical states: solid (alloys), liquid, gas • Miscible vs. Immiscible • Miscible: Liquids that dissolve freely in one another in any proportion • Immiscible :Liquid solutes and solvents that are not soluble • Saturated, Unsaturated and Supersaturated • Dilute vs. Concentrated • Electrolyte vs. Nonelectrolyte

  3. Solubility Curves • Saturated – soln containing the max amt of solute • Unsaturated – soln containing less solute than a sat soln under the existing conditions • Supersaturated – contains more dissolved solute than a saturated solution under the same conditions

  4. supersaturated solution(stirred) Supersaturated Solution of Sodium Thiosulfate

  5. Solubility (physical change) • Definition: mass of solute needed to make a saturated solution at a given temperature • solution equilibrium in a closed system • dissolution ↔ crystallization • Unit = g solute/100 g H2O

  6. Solubility of solids in liquids • For most solids, increasing temperature, increases solubility. • In general, “like dissolves like”. Depends on • Type of bonding • Polarity of molecule • Intermolecular forces between solute and solvent

  7. At 20oC, a saturated solution contains how many grams of NaNO3 in 100 g of water? Saturated sol’n 90 g What is the solubility at 70oC? Supersaturated solution 135 g/100 g water What kind of solution is formed when 90 g NaNO3 is dissolved in 100 g water at 30oC? Unsaturated solution unsaturated What kind of solution is formed when 120 g NaNO3 is dissolved in 100 g water at 40oC? supersaturated

  8. Solubility of Gases • Gases are less soluble at high temperatures than at low temperatures • Increasing temperature, decreases solubility. • Increasing pressure, increases solubility.

  9. Increasing pressure, increases solubility. • The quantity of gas that dissolves in a certain volume of liquid is directly proportional to the pressure of the gas (above the solution).

  10. Effervescence – rapid escape of gas dissolved in liquid

  11. Factors Affecting Solubility • Increase surface area of solute (crushing) • Stir/shake • Increase temperature

  12. Dissolution Process electrolyte nonelectrolyte • Ionic Compounds NaCl(s) Na+1(aq) + Cl-1(aq) • For dissolution to occur, must overcome solute attractions and solvent attractions. • Dissociation Reaction: the separation of IONS when an ionic compound dissolves (ions already present) • Try calcium chloride Dissolving NaCl in water hexahydrated for Na+1; most cations have 4-9 H2O molecules 6 is most common Solvation: process of solvent molecules surrounding solute Hydration: solvation with water

  13. Dissolution Process • Molecular Compounds • Nonpolar molecular solids do not dissolve in polar solvents • naphthalene • Polar molecule • C12H22O11(s)  C12H22O11(aq) • Molecular solvation • Nonelectrolyte • Polar molecule • HCl(g)  H+1(aq) + Cl-1(aq) or • HCl(g) + H2O  H3O+1(aq) + Cl-1(aq) • Ionization: ions formed from solute molecules by action of solvent (no ions initially present) • Nonelectrolyte (HCl)  electrolyte (ions)

  14. Electrolyte vs. Nonelectrolyte

  15. Concentration • Percent concentration by mass (mass % • (solute/solution) x 100% = % Concentration • Molarity (M) • Moles of solute/Liters of solution = mol/L • Molality (m) • Moles of solute/mass of solvent = mol/kg • ppm and ppb • Used for very dilute solutions • Dilution – a process in which more solvent is added to a solution • How is this solution different? • Volume, color, molarity • How is it the same? • Same mass of solute, same moles of solute • In Dilution ONLY – M1V1 = M2V2

  16. Net Ionic Equations • Net ionic equations are equations that show only the soluble, strong electrolytes reacting (these are represented as ions) and omit the spectator ions, which go through the reaction unchanged. • Substances that are aqueous break down into ions. • Substances that pure solids, liquids, or gases do not break down in solution. • Hint: Remember to check solubility rules to determine if a precipitate forms.

  17. Energy Changes • Heat of solution = Hsoln • Endothermic • Solute particles separating in solid • Solvent particles moving apart to allow solute to enter liquid • Energy absorbed • Exothermic • Solute particles separating in solid • Solvent particles attracted to solvating solute particles • Energy released

  18. Colligative Properties • Definition: physical properties of solutions that differ from properties of its solvent. • Property depends upon the number of solute particles in solution. • Types: • Vapor Pressure • Boiling Point ELEVATION • Freezing Point DEPRESSION

  19. Vapor PressureA measure of the tendency of molecules to escape from a liquid • For nonvolatile liquids or solid solutes • A nonvolatile solute will typically increase the boiling point and decrease the freezing point. • Adding a nonvolatile solute lowers the concentration of water molecules at the surface of the liquid. • This lowers the tendency of the water molecules to leave the solution and enter the gas phase. • Therefore the vapor pressure of the solution is LESS than pure water.

  20. H2O solution 100 80 H2O H2O H2O Sugar H2O Vapor Pressure (kPa) 60 40 20 100 Temperature (ºC) Same Temperature

  21. Boiling Point Elevation • tb = boiling point elevation • tb = iKbm • i = molality conversion factor; for ionic compounds adjust for # of ions actually present in solution (dissociation process) • Kb = molalbp elevation constant • Kb = 0.512°C·kg H2O moles of solute (ions or molecules) • m = molality = moles solute kg of solvent • bp of solution = bp of solvent + tb

  22. Freezing Point Depressionwhen a solution freezes, the solvent solidifies as a pure substance; deviates for more concentrated solutions • tf = freezing point depression • tf = iKfm • i = molality conversion factor; for ionic compounds adjust for # of ions actually present in solution (dissociation process) • Kf = molal freezing point depression constant • Kf = 1.858°C·kg H2O moles of solute (ions or molecules) • m = molality = moles solute kg of solvent • fp of solution = fp of solvent - tf

  23. Boiling Point Elevation and Freezing Point Problems • At what temperature will a solution begin to boil if it is composed of 1.50 g potassium nitrate in 35.0 g of water? • Solute: • At what temperature will a solution begin to freeze when 18.0 g ammonium phosphate is dissolved in 200.0 g water? • Solute:

  24. Naming Acids Review: A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid Ex) HCl hydrochloric acid Ex) H3Phydrophosphoric acid B. Tertiary –H + polyatomic anion no Prefix “hydro” (oxo) end “ate” = “ic” acid end “ite” = “ous” acid Ex) H2SO4sulfuric acid Ex) H2SO3 sulfurous acid

  25. Properties of Acids and Bases:

  26. Indicators: Turn 1 color in an acid and another color in a base. • Litmus Paper: Blue and Red An aciD turns blue litmus paper reD A Base turns red litmus paper Blue. • Phenolphthalein: colorless in an acid and pink in a base • pH paper: range of colors from acidic to basic • pH meter: measures the concentration of H+ in solution

  27. Reactions • Neutralization: A reaction between an acid and base. When an acid and base neutralize, water and a salt (ionic solid) form. Acid + Base → Salt + Water Ex) HCl + NaOH → NaCl + HOH

  28. Arrhenius Definition (1884): A. An acid dissociates in water to produce more hydrogen ions, H+. HCl  H+1 + Cl-1 B. A base dissociates in water to produce more hydroxide ions, OH-. NaOH  Na+1 + OH-1 C. Problems with Definition: • Restricts acids and bases to watersolutions. • Oversimplifies what happens when acids dissolve in water. • Does not include certain compounds that have characteristic properties of acids & bases. Ex) NH3 (ammonia) doesn’t fit

  29. Bronsted-Lowry Definition (1923): A. An acid is a substance that can donate hydrogen ions. Ex) HCl → H+ + Cl- • Hydrogen ion is the equivalent of a proton. • Acids are often called proton donors. • Monoprotic (HCl), diprotic (H2SO4) , triprotic (H3PO4) B. A base is a substance that can accept hydrogen ions. Ex) NH3 + H+→ NH4+ • Bases are often called proton acceptors. C. Advantages of Bronsted-Lowry Definition •Acids and bases are defined independently of how they behave in water. •Focuses solely on hydrogen ions.

  30. Hydronium Ion: Hydronium Ion – H3O+ This is a complex ion that forms in water. H+1 + H2O  H3O+1 To more accurately portray the Bronsted-Lowry, the hydronium ion is used instead of the hydrogen ion.

  31. STRONG Acid/Base versus WEAK Acid/Base Strength refers to the % of molecules that form IONS. A strong acid or base will completely ionize (>95% as ions). This is represented by a single ()arrow. HNO3 + H2O  H3O+ + NO3- A weak acid or base will partially ionize (<5% as ions). This is represented by a double (↔) arrow. HOCl + H2O ↔ H3O+ + ClO-

  32. HF < HCl < HBr < HIincreasing strength 7 Strong Acids HNO3 H2SO4 HClO3 HClO4 HCl HBr HI 8 Strong Bases LiOH NaOH KOH RbOH CsOH Ca(OH)2 Sr(OH)2 Ba(OH)2

  33. Strength vs. Concentration • Strength refers to the percent of molecules that form ions • Concentration refers to the amount of solute dissolved in a solvent. Usually expressed in molarity.

  34. Ionization of Acids & Bases • H2SO4 2 H+ + SO4-2 • Sulfuric acid • H3PO3  • Phosphorous acid • Ca(OH)2 • Calcium hydroxide 3 H+ + PO3-3 Ca+2 + 2 OH-1

  35. Conjugate Acid-Base Pairs: A pair of compounds that differ by only one hydrogen ion • Acid donates a proton to become a conjugate base. • Base accepts proton to become a conjugate acid. • A strong acid will have a weak conjugate base. • A strong base will have a weakconjugate acid.

  36. Acid (A), Base (B), Conjugate Acid (CA), Conjugate Base (CB) NH3 + H2O ↔ NH4+ + OH- HCl + H2O ↔ Cl- + H3O+ • Base and Conjugate Acid are a Conjugate Pair. • Acid and Conjugate Base are a Conjugate Pair. B A CA CB A B CB CA

  37. AciDonates & Bases accept • H2O + H2O ↔ H3O+ + OH− B A CA CB • H2SO4 + OH− ↔ HSO4− + H2O A B CB CA • HSO4− + H2O ↔ SO4−2 + H3O+ A B CB CA • OH− + H3O+ ↔ H2O + H2O B A CA CB

  38. The Self-ionization of Water & pH 1. Water is amphoteric, it acts as both an acid and a base in the same reaction. Ex) H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq) Keq = equilibrium constant = [H3O+] [OH-] Because reactants and products are at equilibrium, liquid water is not included in the equilibrium expression @ 25C, [H3O+] = 1 x 10-7M and [OH-] = 1 x 10-7M Kw = ion product constant or equilibrium constant for water Kw = [H3O+] [OH-] = 1 x 10-14M2 1.0 x 10-14 M2= [1.0 x 10-7 M] [1.0x10-7 M] 1.0 x 10-14 = [H3O+] [OH-]

  39. Acids: [H3O+] > 1 x 10-7 M Bases: [OH-] > 1 x 10-7 M Using Kw in calculations: If the concentration of H3O+ in the blood is 4.0 x 10-8M, what is the concentration of OH­ ions in the blood? Is blood acidic, basic or neutral? Kw = [H3O+] [OH-] 1.0 x 10-14M2 = [4.0 x 10-8M] [OH-] 2.5 x 10-7 M = [OH-] slightly basic

  40. The pH scale (1909): the power of Hydrogen • Measure of H3O+ in solution. • pH = -log[H3O+] • Range of pH: 0-14 pH < 7: acid pH > 7: base pH = 7: neutral • pOH = -log[OH-] • pH + pOH = 14

  41. 14 H+ OH- 1

  42. Significant Digits Rule • The number of digits AFTER THE DECIMAL POINT in your answer should be equal to the number of significant digits in your original number • Ex -log[8.7x10-4M] • Calc Answer = 3.0604807474 • Sig Fig pH = 3.06