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Understanding Intermolecular Forces: Polarization and Hydrogen Bonding in Liquids, Lecture notes of Chemistry

University of Detroit Mercy Chemistry

The concept of intermolecular forces and their impact on the properties of liquids. Topics include polar and non-polar intermolecular forces, polarizability, dispersion forces, dipole-dipole interactions, hydrogen bonding, and their effects on surface tension, viscosity, and phase changes. Learn about the differences between polar and non-polar molecules and how their intermolecular forces contribute to various liquid properties.

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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

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Download Understanding Intermolecular Forces: Polarization and Hydrogen Bonding in Liquids and more Lecture notes Chemistry in PDF only on Docsity!

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 L^2 atm/mol^2 ) vs O 2 (a = 1.36 L^2 atm/mol^2 ) ‐‐ water is polar (draw diagram) and O 2 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: NH 3 , H 2 O, SO 2 , SF 4 , XeOF (^4) ‐‐‐ molecules that do not possess a permanent dipole: CBr 4 , BF 3 , BeCl 2 , PCl 5 , I 3 ‐, SF 6 , XeF 4

 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 L^2 atm/mol^2 for O (^2) ‐‐ 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 L^2 atm/mol^2 ) versus Ar (a = 3.59 L^2 atm/mol^2 ) ‐‐ 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. N 2 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

 11.3 Select Properties of Liquids

‐ surface tension: energy needed to separate the molecules of unit area on the surface of a liquid ‐‐ the reason a cold needle floats on the surface of water is because it is not dense enough to break the hydrogen bonds btwn the individual water molecules ‐‐ a hot needle will sink because the added temperature is enough to break the H‐bonds ‐ meniscus: curved surface of a liquid as a result of cohesive (H‐bonds ‐ e.g. water with water) or adhesive (dipole‐dipole ‐ e.g. water with glass) forces btwn solvent and container molecules ‐ capillary action: rise of a liquid up a narrow tube ‐‐ result of cohesive and adhesive forces ‐‐‐ cohesive forces occur btwn liquid molecules ‐‐‐ adhesive forces occur btwn liquid and solid molecules ‐‐ the way in which water flows upwards into trees and plants from the soil ‐ viscosity: resistance of a fluid to flow ‐‐ molasses is very viscous ‐‐ water is not ‐‐ heating a fluid causes the viscosity to lower

 11.4 Phase Changes

‐ Energy Changes Accompanying Phase Changes

‐‐ All phase changes require energy input (endothermic) or release (exothermic)

‐‐ heat of vaporization,  Hvap is the energy

required to vaporize 1 mole of liquid at a pressure of 1 atm ‐‐ Other changes of state ‐‐‐ aside from going back and forth from liquid to gas we also have solid state transitions ‐‐‐ sublimation is another endothermic process in which solid goes to gas ‐‐‐ solidification is the opposite effect in which liquid/gas is solidified ‐‐‐ a solid may melt to form a liquid

‐‐‐ for each one of these processes we have accompanying enthalpies e.g.  Hsub

‐‐‐ one final note: when a change of state is performed the intermolecular forces which led to the initial state must be overcome (e.g. to boil something the intermolecular forces in the

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 = mCT, 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=mCT 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.

 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