Prof Gilbert LECTURE 27 CHEM 1211 11 18 10 Last time Chapter 10 Intermolecular Forces Learning goals connecting macroscopic properties to microscopic structure colligative properties of solutions interpreting phase diagrams Ion ion interactions Determining lattice energy of NaCl from the reaction between Na metal and Cl2 gas Broken down into steps The observed value 411 kJ the sum of the enthalpy changes of the steps U 786 kJ NOTE The enthalpy change in Step 4 is called the first electron affinity EA It is the energy released when one mole of gas phase atoms of an element acquires a mole of electrons This time Chapter 10 continued Why salt dissolves in water or How does the salt water system overcome lattice energy Answer strong ion dipole interactions They are not as strong as ion ion interactions ionic bonds but one ion may interact with as many as 6 molecules Plus there is a big gain in entropy Dipole Dipole interactions and H bonds The partial charges on polar molecules create these interactions If N O or F is bonded to H we have the extreme case of dipole dipole interactions among molecules called hydrogen bonds the dashed lines in the structures below H bonds are only about 1 10 the strength of real covalent bonds but hugely important in defining the shapes of large molecules such as proteins and DNA Among nonpolar molecules their can be temporary dipoles or induced dipoles One example is O2 dissolved in water Water distorts the electron cloud around the O atoms in O O Sometimes random motion of the valence electrons does the same thing The result is a kind of interaction called a London force or dispersion force The bigger the size of the cloud the stronger the interaction Boiling points and Vapor Pressures of Liquids The weaker the intermolecular interactions between molecules in the liquid state the easier it is for them to vaporize reflected in high vapor pressures and low boiling points Vapor pressure increases with increasing temperature When vapor pressure reaches 1 00 atm the substance boils Phase diagrams are graphical presentations such as these of CO2 and H2O below of the dependence of the stabilities of the physical states of a substance on temperature and pressure The lines represent combinations of pressure and temperature at which physical states exist in equilibrium with each other All three exist at the triple point The critical point is located at the end of the liquid gas curve Note the differences between the diagram of CO2 and a portion of the phase diagram of water on the left H2O CO2 Supercritical fluids exist at temperatures and pressures above the critical point Unusual properties they have the viscosities and effusivities of gases but they have the solvent properties of liquids Solubility of Gases Henry s law lets us calculate the solubility of a gas Cgas in a liquid where kH is the Henry s law constant for the gas solvent system A few examples at 20 C Note solubility is proportional to partial pressure What does this mean for people living in say Denver where atmospheric pressure is only 83 of that at sea level Vapor pressure of a solution Consider these two beakers in an inverted plastic tub one w distilled water the other w seawater Explanation adding solute to a solvent reduces the vapor pressure of the solvent Raoult s law of the vapor pressure of a solution Psolution Xsolvent Psolvent Inquiry What is the vapor pressure of automobile antifreeze a 50 50 by volume mixture of ethylene glycol d 1 114 g mL M 62 07 g mol in water at 25 C a temperature at which the vapor pressure of pure water is 23 8 torr Colligative properties of Solutions Properties that depend only on the total concentration of solutes not on their chemical properties Thus a 1 M NaCl solution has the same colligative properties as 1 M KCl or 1 M NaNO3 Osmosis the process by which a pure solvent flows through a porous semipermeable membrane separating it from a solution based on the solvent The membrane can be a cell wall such as that surrounding red blood cells suspended in A normal saline B a highly saline solution such as seawater C deionized water To stop the flow requires an opposing back pressure the magnitude of which is the osmotic pressure of the solution R where M is the molar concentration of all particles in solution T is absolute temperature and R is a constant One common value of R 0 0821 L atm mol K yields osmotic pressures in units of atmospheres If more than the required pressure is applied the osmosis process can be forced to run in reverse in a process used to purify water known as reverse osmosis NOTE The pressures involved can be very high Consider the osmotic pressure produced by seawater at 25 C in which the total concentration of dissolved particles is 1 14 M R mol L atm 298 K 28 0 atm L mol K Boiling point elevation and freezing point depression Solutions than pure molecular boil at higher temperatures solvent A view of why Solutions also freeze at lower temperatures Two handy equations relate boiling point elevation and freezing point depression to the concentration of all solutes in solution expressed in molality m number of moles of solute per kg of solvent Tb Kb m Tf Kf m where the K values are boiling point elevation and freezing point depression constants for the solvent and m is molality NOTE Molality is similar to molarity but note the differences kg of solvent vs L of solution Why is molality used Because pressure and temperature changes can affect the volume a solution occupies but the moles of solute and mass of the solvent are not affected and that is the ratio that defines the boiling and freezing points van t Hoff factor i is the number of moles of particles produced by each mole of solute Sometimes it is a little less than you might think Why are the observed values sometimes less than the theoretical values Because sometimes oppositely charged ions aggregate forming ion pairs which reduce the overall number of independent particles Thus the observed elevation in boiling point of a solution is given by the equation Tb i Kb m