Chapter 12: Intermolecular forces: Liquids, Solids, & Phase Changes
Chapter 12: Intermolecular forces: Liquids, Solids, & Phase Changes CH 121
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This 11 page Class Notes was uploaded by Amelia Notetaker on Tuesday December 1, 2015. The Class Notes belongs to CH 121 at University of Alabama - Huntsville taught by Pamela D Twigg (P) in Fall 2015. Since its upload, it has received 37 views. For similar materials see GENERAL CHEMISTRY I - 90514 - CH 121 - 02 in Chemistry at University of Alabama - Huntsville.
Reviews for Chapter 12: Intermolecular forces: Liquids, Solids, & Phase Changes
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Date Created: 12/01/15
Chapter 12 Lecture Notes Things to remember o Remember that properties of gases can be related by PV=nRT (ideal gas law) o Remember that molecules of gases are far apart, while molecules in liquid and solid are closely packed o Remember that molecules can have polar bonds, due to differences in electronegativity of bonded atoms o Intramolecular bonds are strong Covalent: 100-400 kJ/mol Ionic: 700 -1100 kJ/mol o Intramolecular forces: forces holding atoms together to form molecules (covalent bonds, ionic bonds) o Intermolecular forces: electrostatic interactions between molecules that are weaker than forces between oppositely charged ions Intermolecular forces o Directly related to melting point, boiling point, and energy needed for phase changes (convert solid to liquid or liquid to vapor) o Important in determining solubility of gases, liquids, and solids in various solvents o Crucial in determining structures of biologically important molecules Ion-ion forces, for comparison of magnitude o Na+ - Cl- in salt The strongest forces Lead to solids with high melting temperatures NaCl = 800 degrees C Types of intermolecular forces o Ion-dipole o Dipole-dipole o Dipole-induced dipole o Induced dipole-induced dipole Attraction between ions and permanent dipoles o Water is highly polar and can interact with positive or negative ions to give hydrated ions in water o Example: many metal ions are hydrated. This is the reason metal salts dissolve in water o Strength of attraction evaluated by Coulomb's law (just like in ion-ion attractions) E = Q1Q2/D^2 o Therefore, 3 factors are important Distance between ion and dipole (size) Charge on the ion Magnitude of the dipole o Energy associated with hydration of ions is known as solvation energy (Esolv) or enthapy of hydration (delta H) Cannot be measured directly, but estimated Hydration energy o Explain the difference in enthalpies of hydration for Mg+2 (-1922 kJ/mol), Na+ (-405 kJ/mol), and Cs+ (-262 kJ/mol). Why is the enthalpy of hydration of Mg+2 so much more exothermic than the other 2? Determine the relative sizes of the ions. Determine the magnitude of the charges on the ions Sizes: Mg+2 = .66, Na+ = .97, and Cs+ = 1.65 Mg+2<Na+<Cs+ for size, and consequently distance between ion and dipole center (d) Delta H increased with decreasing d Delta H also increases with increasing charge Therefore delta H (Mg+2)>> delta H (Na+)> delta H (Cs+) Dipole-dipole forces o Such forces bind molecules having permanent dipoles to one another Dipole-dipole interactions influence evaporation of a liquid and condensation of a gas Delta Hvap is positive, therefore endothermic The greater the interaction, the greater the energy required for vaporization o Influence of dipole-dipole forces is seen in the boiling points of simple molecules o Attraction between dipoles depends on dipole moment (strength) and molecular size (distance) E= Q1Q2/d^2 Intermolecular forces influence solubility o "like dissolves like" o Polar molecules more likely to dissolve in a polar solvent o Nonpolar molecules more likely to dissolve in a nonpolar solvent Forces involving induced dipoles o If "like dissolves like", then how can nonpolar molecules such as O2 and I2 dissolve in water? o The water dipole induces a dipole in the O2 electron cloud The dipole of water induces a dipole in O2 by distorting the O2 electron cloud Dipole-induced dipole o Solubility increases with mass of the gas because larger molecules easier to polarize o Consider I2 dissolving in alcohol CH3CH2OH The alcohol temporarily creates or induces a dipole in I2 o Formation of a dipole in 2 nonpolar molecules Induced dipole-induced dipole also known as "london dispersion forces" 2 nonpolar atoms or molecules --> momentary attractions and repulsions between nuclei and electrons in neighboring molecules lead to induced dipoles --> correlation of the electron motions between the 2 atoms or molecules (which are now dipolar) leads to a lower energy and stabilizes the system Dispersion forces operate between all molecules o The magnitude of the induced dipole depends on the tendency to be distorted High molecular weight --> larger induced dipoles Greater induced dipoles --> increased interactions (higher BP) Things to remember about the strength of intermolecular forces o E = Q1Q2/d^2 o Strength of ion-dipole bond increases with Increasing charge on ions Decreasing distance between centers (size of ion) Increasing magnitude of the dipole o Solubility of non-polar gases in dipolar solvents increases with Increasing size of gas molecule (larger molecules easier to polarize) o Strength of induced-dipole/induced-dipole interactions increases with Increasing size of molecule o When we walk about strength of intermolecular bonds, we think in terms of "heat vaporization" or detla Hvap This can be visualized as the amount of energy it takes to free a molecule from the surface of a liquid into the gas phase If the intermolecular bonds are strong, then more energy is required (higher delta Hvap) to accomplish this In general, delta Hvap is positive, so the process is endothermic Which molecule would you expect to have the highest boiling point? o Br2, N2, or I2? First identify the type of molecule - all are nonpolar Next, identify possible bond types - induced dipole - induced dipole Finally look at size - I2 > Br2 > N2 Larger size = larger heat or vaporization (higher boiling point) Hydrogen bonding o A special form of dipole-dipole attraction, which enhances dipole-dipole attractions o Hydrogen bonding is the strongest when X and Y are N, O, or F They have some of the largest electro negativities of any element Have unpaired electrons Creates a very strong dipole Hydrogen bonding in H2O o H bonding is especially strong in water because o There are 2 lone pairs on the O atom (each can bond to a H on another molecule) o The O-H bond is very polar o This accounts for many of the water's unique properties o In liquid state, water form irregular, continuous networks of hydrogen bonds o On the other hand… In the solid state, water (ice) forms a regular, open lattice structure The density of ice is lower than that of liquid water Ice has open lattice-like structure Ice density is < liquid and so solid floats on water This is the reason lakes freeze over from the top down instead of form the bottom up The tetragonal bonding arrangement in ice leads to the hexagonal symmetry seen in snowflakes Strong hydrogen bonds in water lead to an unusually high heat capacity This is why water is used to put out fires, why lakes/ocean control climate, and why thunderstorms release large amounts of energy Hydrogen bonding in biology o Hydrogen bonding in protein folding Stabilizes secondary structure Alpha-helix Beta-sheet Beta-turns o H-bonding is also important in stabilizing the double helix structure of DNA o DNA - helical chains of phosphate groups and sugar molecules. Chains are helical because P, C, and O have tetrahedral geometry o 2 chains held together by specific hydrogen bonding between pairs of bases Adenine with thymine Guanine with cytosine Intermolecular forces summary o Dipole-dipole: dipole moment (depends on atom electro-negativities and molecular structure) o Hydrogen bonding: very polar X-H bond (where X = F,N,O) and atom Y with lone pair electrons. An extreme form of dipole-dipole interaction o Dipole/induced dipole: dipole moment of polar molecule and polarizability of nonpolar molecule o Induced dipole/induced dipole (london dispersion forces): polarizability Properties of liquids o In a liquid Molecules are in constant motion There are appreciable intermolecular forces Molecules are close together Liquids are almost incompressible Liquids do not fill the container o Evaporation: process by which a liquid become a gas (also known as vaporization) and its opposite --> condensation o Evaporation --> add energy break IM bonds o Condensation --> make IM bonds remove energy Liquid - evaporation o To evaporate molecules must have sufficient energy to break IM forces o Breaking IM forces requires energy o The process of evaporation is endothermic o At higher temperature, a much larger number of molecules have high energy to break IM forces and move from liquid to vapor state Liquids - distribution of energies o Molecules is a liquid have a distribution of kinetic energies o The average kinetic energy is proportional to temperature Liquids - equilibrium vapor pressure o When molecules of liquid are in the vapor states, they exert a vapor pressure o Equilibrium vapor pressure is the pressure exerted by the vapor over a liquid in a closed container (at a given temperature) when the rate of evaporation = the rate of condensation o Vapor as a function of temperature The curves show all conditions of P and T where liquid and vapor are in equilibrium The vapor pressure rises with temperature When vapor pressure = external pressure, the liquid boils, This means the boiling points of liquids change with altitude Liquid boils when its vapor pressure equals outside pressure o If external pressure = 760 mm Hg (standard atm pressure), temperature of boiling is called the normal boiling point o Vapor pressure of a given molecule at a given temperature depends on IM forces Boiling Point at lower pressure o When pressure is lowered (as at higher elevations, or under a vacuum), the vapor pressure will equal the external pressure at a lower temperature, so boiling point is lowered Liquids o Heat of vaporization is the heat required (at constant pressure) to vaporize the liquid o Molecules at surface behave differently than those in the interior Water molecules on the surface are not completely surrounded by other water molecules Water molecules under the surface and completely surrounded by other water molecules o Molecules a surface experience net inward force of attraction This leads to surface tension - the energy required to break the surface o Intermolecular forces also lead to capillary action and to the existence of a concave meniscus for a water column Adhesive forces between water and glass: concave meniscus Cohesive forces between water molecules: H2O in glass tube Surface tension o Prevents a paper clip from sinking when placed in water o Also leads to spherical water drops Capillary action o Movement of water up a piece of paper depends on H-bonds between H2O and the OH groups of the cellulose in the paper Viscosity - the resistance of liquids to flow o Strongly dependent on strength of intermolecular forces Properties of solids o Molecules or atoms cannot change their relative positions o Solid state characterized by long-range order - regular, repeating pattern Cubic unit cell o Unit cell is the simplest repeating unit in a crystal lattice o Cubic unit is easy to visualize and fairly common o 3 types Primitive (simple cubic- sc) 1 atom Each corner atom is 1/8 inside the unit cell Body centered cubic (bcc) 2 atoms Each corner atom is 1/8 inside unit cell and 1 atom is inside the cell Face centered cubic (fcc) 4 atoms Each corner atom is 1/8 inside the unit cell, each face atom is 1/2 inside the cells (6 faces and 1/2 per face = 3 net atoms) Finding the lattice type o Al has density of 2.7 g/cm^3 and radius = 143 pm. Verify that Al is fcc Calculate density for a fcc cell with Al atoms and compare with the reported density Density = mass/volume Need to know volume of cell and number of atoms per cell to calculate density For fcc, the number of atoms per cell is 4 To calculate volume you need the length of cube side Use pythagorean theorem: a^2 + b^2 = c^2 (Al diameter =2)(radius = 2) (143 pm) Therefore, (c=2)(2(143 pm)) = 572pm To calculate cube volume, the side of a must be known 2a^2=c^2 Sqrt2(a) = c A=(572 pm)/Sqrt2 = 404 pm V=a^3, convert length to cm A= 404 pm = 4.04 x 10^-8 cm V= (4.04 x 10^-8 cm)^3 V= 6.62 x 10^-23 cm^3 Mass 1 Al atom = 1 mol 26.98 g 6.022 x 10^23 atoms 1 atom = 4.48- x 10^-23 g 4 atoms = 17.92 x 10^-23 g Density = mass/volume (17.92 x 10^-23 g)/6.62 x 10^-23 cm^3 = 2.707 g/cm^3 Structure and formulas of ionic solids o Many ionic compounds are formed by taking a simple cubic cell of one ion and placing ions of opposite charge in the lattice holes o The chemical formula for Cesium chloride can be determined by counting the number of each ion in the unit cells: 1 Cl- and 8(1/8 Cs+) = 1 Cs + Chemical formula is CsCl Phase diagrams o Substances can exist as solid, liquid, or gas, depending on temperature and pressure conditions o Under some conditions, 2 or all 3 phases can exist in equilibrium o Phase diagrams are graphs showing the relationships between phases of matter and T and P o Transitions between phases Lines connect all conditions of T and P where equilibrium exists between the phases on either side of the line At equilibrium, atoms or molecules move from liquid to gas as fast as they move from gas to liquid for example o Triple point - water P = .6 mmHg T = .01 degrees C At the triple point all 3 phases are in equilibrium o Important points for water Normal boiling point : 100 degrees C to 760 mmHg Normal freezing point: 0 degrees C to 760 mmHg Triple point: .0098 degrees C to 4.58 mmHg o Critical T and P As P and T increase, you finally reach the critical T and P Note that line goes straight up Above critical T no liquid exists no matter how high the pressure Solid-liquid equilibria o In any system, if you increase P the density will go up o Therefore as P goes up equilibrium favors phase with the greater density (or smaller volume/gram) o At solid/liquid equilibrium, raising P squeezing the solid o It responds by going to phase with greater density the liquid phase o Raising the pressure at constant T causes water to melt (think ice skating) o The negative slope of the S/L line is unique to H2O Almost everything else has positive slope Solid-vapor equilibria o At P < 4.58 mmHg and T < .0098 degrees C solid H20 can go directly to vapor Sublimation Clausius-calpeyron equation o clapeyron showed that for a pure liquid, there is a linear relationship between ln(P) (Vapor pressure) and 1/T (in kelvin) Y=mx+b Ln P = -(delta H degrees vapor/RT)+C o Where delta H vap is enthalpy of vaporization, R is ideal gas constant, and C is a constant characteristic of the compound Slope of the line = -delta H vap/R o By measuring vapor pressure of a liquid at 2 temperatures, you can calculate the enthalpy of vaporization Ln P = -(delta H degrees vap/RT) + C Ln (P1) = -(delta H degree vap/R) [1/T1] +C Ln (P2) = -(delta H degree vap/R) [1/T2] +C Ln (p2)-ln(P1) = [-(delta H degree vap/R)[1/T2]+ C] - [-(delta H degrees vap/R)[1/T1] + C] Or ln (P2/P1) = (delta H degrees vap/R)[1/T1-1/T2]
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