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Chapter 8 — Solutions and Their Behavior. Some Terminology. Solution A solution is a homogeneous mixture (not possible to see boundaries between components) that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent

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Chapter 8 — Solutions and Their Behavior

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Chapter 8 — Solutions and Their Behavior


Some Terminology

  • Solution

    • A solution is a homogeneous mixture (not possible to see boundaries between components) that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent

  • Solvent is present in the larger amount

  • Solute is present in smaller amount than the solvent

  • Examples:

    Liquids: NaCL (sodium chloride + water)

    0.89% 99.11%

    (solute) (solvent)

    Gases: Air is a solution of N2, O2, few other minor gases


Units of Concentration 1

  • Molality (m) aka “molal concentration”

    Solvent= does not include solute

  • Molarity (M) aka “molar concentration”

    Most common concentration unit in Chemistry

Solution=solute+solvent


Example

  • What is the molarity of a solution prepared by dissolving 45.0 grams of NaCl in enough water to give a total volume of 489 mL?

    Na = 23 a.m.u.

    Cl = +35 a.m.u.

    58 a.m.u. 1mol NaCl = 58g

    45g X 1 mol = 0.7759 mol of NaCL

    58g

    0.7759 mol of NaCl = 1.6 M

    0.489 L


Molality vs. Molarity

  • Molality is never equal to molarity

    • But the difference becomes smaller as solutions become moredilute(denominators are very similar)

  • Molarity is more useful when dealing with solution stoichiometry

  • Molality is more appropriate for dealing with physical chemistry

  • Question: which is temperature-dependent?

    • Molarity depends on temperature. Molarity will decrease as temperature increases since the amount of the solution will decrease (from evaporation). Temp  M

    • Molality does not depend on temperature since mass (kg) does not change with temperature


Units of Concentration 2

  • Mole Fraction ( )

  • Percent by Volume (% w/v) AND Percent by Weight (%w/w)


Example

  • Calculate the percent by weight NaCl in a solution comprised of 45.0 g NaCl and 457 g of water.

    % (w/w) = grams of solute x 100% = 45g x 100% = 8.96%

    grams of solution 502g

    Solute = 45g of NaCl

    Solution = 457g of NaCl (solute) + H2O (solvent) = 502g


Units of Concentration 3

  • Parts per million (ppm): Extremely dilute solutions. Compares amount of solute to a million parts of solution (rather than 100 parts).

  • Parts per billion (ppb) Even more extremely dilute solutions. Compares amount of solute to a billion parts of solution (rather than 1million parts).


Converting Between Units

  • Every concentration unit is a ratio of two quantities

  • Pick a sample size

    This fixes one of the two quantities

  • Use the factor-label method (dimensional analysis a.k.a. “conversion factor”) to systematically convert the given quantities into the desired quantities


Some More Terminology

  • Solubility

    • Solubility is the amount of solute that will dissolve in a given amount of solvent at a given temperature

  • Saturated

    • A saturated solution contains the maximum amount of a solute, as defined by its solubility (no more solute will dissolve in a solution that is already saturated with that solute)

  • Supersaturated

    • A solution contains more solute than allowed by the solubility. Not a stable system since there is more solute dissolved than the solvent can “accommodate” and the excess solute will come out of solution and will crystallize as a solid (e.g., calcium oxalate or CaPO4 will make kidney stone) or it will separate as a liquid or it will bubble out as a gas (e.g., soda)


Some More Terminology

  • Miscible

    • Two liquids are miscible if they are soluble in each other in all proportions

  • Question: Can a solution be both saturated and dilute?

    Yes, just remove some of the solvent.

    If you want to see this for yourself, mix a little salt and water together. Then leave it stand so most of the water evaporates. You have saturated a dilute solution.


Solubility Guidelines

  • Like dissolves like

    • Polar solutes are more soluble in polar solvents

    • Nonpolar solutes are more soluble in nonpolar solvents


Heat of Solution (Energy change or Enthalpy of Solution)

ΔH°solution is defined as the enthalpy (energy) change that accompanies dissolving exactly 1 mole of solute in a given solvent

Enthalpy = Heat (as long as pressure remains constant)

  • Some substances have positive (endothermic) ΔH°solution and some have negative (exothermic) ΔH°solution

  • Endothermic: energy flows INTO the system (gain of energy, temperature of the system decreases)

  • Exothermic: energy flows OUT of the system (loss of energy, temperature of system increases)

    ΔH°solution is the sum of two terms:

    Lattice Energy and Solvation Energy


LE and Solvation

  • Lattice energy is the energy released when molecules or ions settle into a crystalline lattice. It is inherently exothermic (opposite charges coming towards each other)

  • The opposite: Pulling all the ions away from each other in an ionic substance in the solid state requires overcoming the lattice energy (endothermic)

  • Solvation (hydration) energy is the energy released when an ion (or molecules) settles into a sphere of solvent molecules

    • Solvation is inherently exothermic (opposite charges coming towards each other also)

  • Both LE and SE result from IM forces


  • Sodium Chloride: Exothermic DH°solution

    The magnitude of the lattice energy is lower than the magnitude of the solvation energy, in going from solid sodium chloride to a solution of sodium chloride so the energy of this system has decreased (it has undergone an exothermic change)


    Ammonium Chloride: Endothermic DH°solution

    The magnitude of the lattice energy is greater than the magnitude of the solvation energy, in going from solid ammonium chloride to a solution of ammonium chloride so the energy of this system has increased (it has undergone an endothermic change)


    Effect of Pressure on Solubility

    • Gaseous solute

      • As P increases, solubility increases

      • Henry’s Law:

        S = kHPgas

        S=solubility, k=Henry’s constant, P=partial pressure of the gas

    • Liquid and Solid Solutes

      • P has negligible effect (liquids and solids are not very compressible)


    Example

    • The Henry’s law constant for oxygen in water is 0.042g/L/atm at 25°C. What is the solubility (in grams per mL) of O2 in pure water when the atmospheric pressure is 740 torr. In air, the mole fraction of oxygen is 0.21.

      Ptotal = 740 torr x 1 atm = 0.974 atm

      760 torr

      PO2 = 0.21 x 0.974 atm = 0.204 atm

      S = 0.042g/L/atm x 0.204 atm = 0.0086 g/L = 8.6 mg/L


    Colligative Properties of Solutions

    • A colligative property depends only on the number of solute particles, not the identity of the solute particles

    • Colligative properties include:

      • Vapor pressure of a solution decreases with increasing [solute]

      • Boiling point of a solution increases with increasing [solute]

      • Freezing point of a solution decreased with increasing [solute]

      • Osmotic pressure of a solution increases with increasing [solute]


    Vapor Pressure Decreases

    • Introduce a solute (solute will “plug escape sites” of molecules that are at the surface of the solution “ready to jump out” and become gas)


    Raoult’s Law (vapor pressure of a volatile component of a solution (P) is equal to the vapor pressure of the pure substance Po) times the mole fraction (X) of that substance)

    • The vapor pressure of a solution is given by Raoult’s Law


    Example (Textbook page 206)

    • What is the vapor pressure of a 5% (by mass) sugar solution at 45°C? (P°water = 71.9 torr)


    Example (Textbook, page 206)

    • What is the total vapor pressure of a solution comprised of 0.75 moles of benzene and 0.45 moles hexane at 60°C? (P°benzene= 390 torr; P°hexane= 580 torr)


    Boiling Point Elevation(BP: Temperature at which the vapor pressure of the material is equal to the ambient pressure)

    • As vapor pressure goes down, boiling point goes up

      ΔTbp = Kbpmsolute

      ΔTbp is the boiling point elevation

      Kbp is the boiling point (ebullioscopic) constant

      msolute is the molality of all solute particles


    Freezing Point Depression(Freezing point: temperature at which the liquid phase of the material is in equilibrium with the solid phase (aka melting point)

    ΔTfp = Kfpmsolute

    is ΔTfp the freezing point depression

    Kfp is the freezing point (cryoscopic) constant

    msolute is the molality of all solute particles


    Ionic Solutes

    • When ionic solutes dissolve, they dissociate into solvated ions

    • Each ion counts as a particle for colligative properties


    Osmosis

    • Osmosis is diffusion of water through a semipermeable membrane

    • Solute particles are too big (or too polar) to make it across the membrane

    • This is how water gets moved around cells


    Tonicity

    • Isotonic solutions have equal concentrations of solute particles

    • A hypertonic solution has a greater concentration of solute

    • A hypotonic solution has lower concentration of solute


    Example

    Hypertonic solution

    Less water

    Hypotonic solution

    More water


    Osmotic Pressure (Increases with increasing solute concentration)

    • Osmotic pressure (P) results from the potential drive for the concentration of water to equalize

      P V = nRT

      Or

      P = MRT

      A 1.0 M solution of glucose exerts an osmotic pressure of

      22.4 atm at 25°C


    Question

    • What will happen to a red blood cell when it is placed into pure water? Cells are isotonic with normal saline (0.89% NaCl)


    Question

    • What will happen to a red blood cell when it is placed into 10% aqueous sodium chloride? Cells are isotonic with normal saline

      (0.89% NaCl)


    Colloids

    • Colloids are not true solutions

    • Particle size is on the order of 10 to 200 nm

      • Might be super-sized molecules (e.g., proteins) or aggregates of ions

    • Colloidal particles cannot be filtered and do not settle out of solution

    • Colloids exhibit the Tyndall effect (whereas solutions don’t)

      • The particles in a colloid are large enough to scatter light passing through)

    • Examples:

      • blood, milk, jelly, propofol


    Surfactants

    • A surface active agent breaks the surface tension of water

    • Surfactants improve a solvent’s ability to be a solvent

    • Soaps and detergents are common surfactants


    polar head

    greasy tail

    Soaps and Detergents

    • A soap is prepared by hydrolyzing fat with alkali

    • A detergent is a synthetic chemical with a structure similar to soap

    • Both have a polar (hydrophilic) head and a non-polar (hydrophobic, greasy) tail


    Monolayers

    • The greasy tails stick out of the surface of water

    • This breaks down the surface tension of water


    Bilayers

    • The tails can dissolve in each other in a bilayer

    • This structure is used in cell membranes


    Micelles

    • The tails can dissolve in each other forming a sphere

    • This creates a non-polar microenvironment in the water


    Thank you!


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