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Chapter 11 – Liquids & Intermolecular Forces 11.1 A Molecular Comparison of Gases, Liquids, and Solids The state of a substance is a balancing act between how fact the molecule is moving (kinetic energy) and interactions between particles (intermolecular forces) ‐ The fundamental difference between states is the strength of the intermolecular force ‐ Stronger forces bring molecules closer together ‐ Solids and liquids are referred to as condensed phases <Recall Polar Covalent Bonds & Dipole Moments> ‐ van der Waals constant for water (a = 5.28 L2atm/mol2) vs O2 (a = 1.36 L2atm/mol2) ‐‐ water is polar (draw diagram) and O2 is non‐polar ‐‐‐ recall the Electronegativity (EN) Trend ‐‐ interaction between water molecules are electrostatic ‐ polar bonds and polar molecules ‐‐ bond dipole ‐‐‐ change in EN between 2 atoms makes the bond connecting them polar ‐‐‐ this phenomenon leads to a bond dipole (arrow head points to the more EN atom) ‐‐ permanent dipole moment (see figure ) ‐‐‐ a molecule has a permanent dipole moment when it possesses an asymmetric orientation of polar bonds ‐‐‐ molecules that possess a permanent dipole: NH3, H2O, SO2, SF4, XeOF4 ‐‐‐ molecules that do not possess a permanent dipole: CBr4, BF3, BeCl2, PCl5, I3‐, SF6, XeF4 11.2 Intermolecular Forces ‐ Intramolecular =inside a single molecule versus Intermolecular = between two or more molecules ‐‐ Intramolecular forces will impact bond energies (polar versus covalent) ‐‐ Intermolecular forces will impact things like melting/freezing and boiling points ‐ Dispersion Forces Why do van der Waals constants have nonzero values for nonpolar species? ‐‐ recall a = 1.36 L2atm/mol2 for O2 ‐‐ polarizability: refers to the distortion of the electron cloud around the atom's nucleus as another atom or molecule approaches ‐‐ this distortion occurs as a result of electron‐electron repulsion btwn the atom and the approaching species ‐‐ the larger a molecule is the less tightly the electrons are held to the nucleus ‐‐‐ this makes it easier to distort the electron cloud ‐‐‐ therefore larger molecules are more polarizable ‐‐ comparison btwn He (a = 0.0341 L2atm/mol2) versus Ar (a = 3.59 L2atm/mol2) ‐‐ London or dispersion forces: interactions btwn induced dipoles ‐‐‐ when an atom is polarized in the presence of another species this induced dipole occurs ‐‐‐ this is the type of interaction which happens btwn two non‐polar species ‐ e.g. N2 molecules ‐‐‐ factors that impact the strength of this force: ‐‐‐‐ molecular weight (aka size) – the larger the more polarizable and therefore the larger the force ‐‐‐‐ molecular shape – when two molecules have the samemolecular formula then the shape that maximizes surface area will have a greater induced dipole ‐‐ Usually this intermolecular force is considered to be the weakest ‐‐ It is also the only force that is present in ALL neutral molecules ‐ Dipole‐Dipole ‐‐ we have attraction btwn the negative (pink) and positive “poles” in blue shown with solid red lines ‐‐ we also have repulsion btwn the “poles” which are charged the sames shown with dashed blue lines ‐‐ these dipoles happen because electron density is pulled from the less electronegative atoms toward the more electronegative ones liquid must be overwhelmed with our heat to the point molecules escape from the liquid to the gas phase) ‐ Heating Curves ‐‐ AB we are heating up to the freezing point of water, Tf (recall Ch 5 q = mCT, q = heat, m = mass, C = specific heat of solid, T = change in temperature) ‐‐ BC represents the heat of fusion, Hfus, which allows a phase change from s to l ‐‐ CD we are heating up the liquid from Tf to Tb (boiling point, use q=mCT where C is for liquid) ‐‐ DE represents the heat of vaporization, Hvap, which allows a phase change from l to g ‐‐ EF we are heating up the gas from Tb to final temperature ‐ Critical Temperature & Pressure ‐‐ All substances have a T & P in which the liquid and gas phases are completely indistinguishable this is called the critical point ‐‐‐ the density is the same for both states ‐‐‐ the liquid phase is less dense due to high T ‐‐‐ the gas phase is more dense due to high P ‐‐ The name we give to this state is supercritical fluid ‐‐ We will talk more about this in 11.6 11.5 Vapor Pressure ‐ vapor pressure is a result of molecules escaping from the liquid phase as gas ‐ vaporization/evaporation is an endothermic process because energy/heat must be added to the system for a molecule to escape the liquid phase ‐ when the rate of the liquid escaping to the gas is equal to the rate of a gas returning to liquid we have an example of equilibrium ‐ Volatility, Vapor Pressure & Temperature ‐‐ A volatile liquid is one that evaporates and does not readily return to liquid ‐‐ hot water will evaporate more quickly than cold because there is more energy present in the form of heat to break the H‐bonds between water molecules ‐ Vapor Pressure & Boiling Point ‐‐ As the temperature is increased so is the vapor pressure ‐‐ when T increases so do the molecular motions and the ability for a molecule to escape from the liquid and go into the gas phase ‐‐ since the pressure of the atmosphere is lower at higher elevations ‐ less temperature is required for water to boil ‐‐ One of the way we can use vapor pressure is to calculate the heat of vaporization by plotting the ln of the vapor pressure versus the inverse of the corresponding T: ‐‐‐ This produces a linear equation in the slope encompasses Hvap 1 ln vap vap b y xm H P C R T ‐‐‐ This is called the Clausius‐Clapeyron Equation and you will use this in the Heat of Vap lab ‐‐‐ We can alsoHvap get by measuring the vapor pressure at two different temperatures using: 1 2 , , 2 1 1 1 ln vap T vap vap T P H P R T T 11.6 Phase Diagrams ‐ the strength of intermolecular forces, temperature and pressure contribute to the phase (g, l, or s) of a molecule ‐ at lower temperature and pressure gases are preferred whereas solids are more likely at higher temperatures and pressures ‐ phase diagrams: graphical representation of the physical states as a function of T&P ‐‐ lines in the diagram: 1) melting point line: the state corresponding to this line is both solid and liquid 2) boiling point line: both gas and liquid present 3) sublimation point line: gas and solid present ‐‐ note for water this line has a negative slope ‐ due to H‐bonds ‐‐ usually this is positively sloped ‐‐ the sublimation from ice to gas is what causes ice cubes to get smaller in the freezer ‐‐ triple point: point at which all three phases are present ‐‐ critical point: point at which liquid and gaseous state is indistinguishable ‐‐‐ the liquid phase is less dense due to high temperature ‐‐‐ the gas phase is more dense due to high pressure ‐‐‐ densities of the two states are the same ‐‐ normal points occur at 1 atm of pressure ‐‐‐ normal bpt for water is 100C ‐‐‐ normal fpt for water is 0C ‐‐ above this critical temperature and pressure we get a supercritical fluid ‐‐‐ has the physical properties of gas ‐‐‐ has the ability to dissolve substances like a liquid ‐‐‐ e.g. supercritical CO2 to create decaffeinated coffee 11.7 Liquid Crystals – Skip it!