Lecture 8 - Chemistry - Liquids and Intermolecular Forces PDF

Lecture #8, p. 1 Lecture 8: Announcements Today: Brown Ch. 11 Liquids and Intermolecular Forces 11.1 Molecular Comparison of Gases, Liquids, and Solids 11.2 Intermolecular Forces...

Lecture #8, p. 1 Lecture 8: Announcements Today: Brown Ch. 11 Liquids and Intermolecular Forces 11.1 Molecular Comparison of Gases, Liquids, and Solids 11.2 Intermolecular Forces 11.3 Select Properties of Liquids 11.4 Phase Changes 11.5 Vapor Pressure 11.6 Phase Diagrams Problem Set 7: Due before Exercise #8 tomorrow; upload on Moodle link Problem Set 8: Posted on Moodle; due before Exercise #9 next week Study Center: Wednesdays 18:00–20:00 in ETA F 5 Office Hours: NO OFFICE HOUR THIS WEEK Chemistry  Lecture #8, p. 2 Lecture 9 Next Week: Brown Ch. 13 Properties of Solutions 13.1 The Solution Process 13.2 Saturated Solutions and Solubility 13.3 Factors Affecting Solubility 13.4 Expressing Solution Concentration 13.5 Colligative Properties Chemistry  Lecture #8, p. 3 Review In Lecture 7, we discussed chemical bonding and molecular shape Basics of chemical bonding: ionic, covalent, metallic Lewis symbols, octet rule Ionic bonding, covalent bonding Multiple bonds, bond polarity, electronegativity, dipole moments Drawing Lewis structures, alternative Lewis Structures Oxidation number vs formal charge vs partial charge Resonance structures, exceptions to octet rule Bond enthalpies, strengths and lengths to covalent bonds Molecular shapes for ABn VSEPR model, electron domains Chemistry  Lecture #8, p. 4 Today: Phases Lecture 2: gases Recall three states of matter: solid, liquid, vapor Example: H2O Physical properties of phases? Why do substances change phase? Chemistry  Lecture #8, p. 5 Explain Boiling Process? Chemistry  Lecture #8, p. 6 in iii itis is_is Intermolecular Forces Last time: intramolecular bonding Interactions between molecules Also, important ! Why? Intermolecular forces affect many properties we care about Examples? Phase changes Lubrication Viscosity Surface tension Wettability Ex: Gore-tex fabric source: wikipedia Engineers need understanding to control / exploit these effects! Chemistry Iii  Lecture #8, p. 7 Let'sstartwithsomebasicproperties General Properties of Phases Property Gas Liquid Solid Fills container? Yes No No Takes shape of container? Yes Yes No Flows? Yes Yes No Molecules are close? No Yes Yes Molecules are ordered? No No Yes Liquidsan Compressible? Yes No No !! versus !"#$ ? T.EE ii How can we explain all of these differences? Chemistry  Lecture #8, p. 8 !! versus !"#$ ? !! Average kinetic energy of molecules Remember from Lecture #2 Proportional to $% with !$ > 0 !"#$ Potential energy due to intermolecular interactions !!"# is the intermolecular interaction energy If molecules attract each other, !!"# < 0 !!"# set by molecules and their separation distance ' Competition between !! and !"#$ As temperature decreases, !! decreases relative to !"#$ Rationalizes gas to liquid to solid with decreasing " However, we know now that entropy is also involved Chemistry BecauseVrms with Rgasconstant iiiit ia  Lecture #8, p. 9 General Properties of Phases Property Gas Liquid Solid Fills container? Yes No No Takes shape of container? Yes Yes No Flows? Yes Yes No Molecules are close? No Yes Yes Molecules are ordered? No No Yes Compressible? Yes No No !! versus !"#$ ? !! ≫ !"#$ !! ≈ !"#$ !! ≪ !"#$ All of these properties can be explained by intermolecular interactions Chemistry  Lecture #8, p. 10 SowhatistheoriginofEint Intermolecular Interactions? If molecules are neutral ⇒ three types of interactions A. Dispersion (London dispersion force) nucleus e− Electron e− induces electron moves randomly e− e− on second molecule to right electron electron to move away cloud cloud molecule #1 molecule #2 Induced-dipole– Attractive at "% "& "% "& induced-dipole short distances! attraction Chemistry 1 I affities Iiiii.EEiiiit  Lecture #8, p. 11 In dispersioninteractions an instantaneous chargefluctuation ononeatominducesattractivedipoledipoleinteractions Butinpolarmolecules adipoleisalreadypresent Intermolecular Interactions? Endathole B. Dipole–Dipole !! H F !" #⃗ = %& Recall polar bonds ⇒ dipole moment, #⃗ #⃗ But for #⃗$ ! ! f Molecules can also have an overall dipole moment H !" O #⃗%&% ⇒ vector sum of all bond dipole moments H ! #⃗# ! Ex Cay Polar molecules have #⃗%&% ≠ 0 Molecular dipoles interact ⇒ attractive *'(% ! " ! ! on different molecules attract over short distances: Chemistry s nt ar morepolar he.fi fti efr  Lecture #8, p. 12 JohannesvanderWaals 18371923 Intermolecular Interactions? Dispersion and Dipole–Dipole Together known as van der Waals forces Ex: ⇒ Always attractive ⇒ Always present ⇒ Increase with molecular size ⇒ Affected by molecular shape Linear versus branched shape pentane dimethylpropane ⇒ More van der Waals interactions possible C5H12 C5H12 between linear pentane molecules Tbp = 309 K Tbp = 282 K i t Chemistry sameMw cangetme Stiersnext Thus Eintinliquidishigher us figitaining mterature  Lecture #8, p. 13 Onetypeofdipoledipoleinteraction is so important it has its ownname Intermolecular Interactions? C. Hydrogen Bonding In molecules with N H, O H, or F H bonds !" !! !" !! !" !! N, O, and F are very electronegative ⇒ Very polar bonds with H ! H! Attracted to unbonded electrons on another nearby N, O, or F Small so can get close ⇒ Large "!"# O H H H Very important in biology Ex: H2O O H H Explains unusual properties of H2O O H Chemistry H2Ohasexceptionallyhigh boiling temperature for t.sientii i iIti'it  Lecture #8, p. 14 Above was forneutralmolecules Wecan alsoconsiderions Ion–Dipole Interactions µ Very important for solutions (Lecture 3) Ions solvated by polar liquid (e.g. NaCl in H2O) Iiii Polar molecules can solvate both positive and negative ions by orientation is Polar molecules are good solvents for ions This is true not just for H2O Chemistry More in Lecture 9  Lecture #8, p. 15 IE IE effire Strength of Intermolecular Interactions Weaker than real bonds tailitant able8.3 Recall covalent C H bond ⇒ 413 kJ/mol romBrown Ion–Dipole H-Bonding Dipole–Dipole ≈ Dispersion ~50 kJ/mol > ~25 kJ/mol > ~10 kJ/mol I fi ies These effects can also be additive ! Note: If 2 molecules are close enough ⇒ Repulsive interaction molecule #1 molecule #2 Chemistry Ifntignfingtoaf.is iectIci dtie leitriitsorp each other  Lecture #8, p. 16 Sowhydowecare about intermolecularinteractions Properties of Liquids Explained by intermolecular interactions Units Viscosity ⇒ Resistance to flow m ⇒ Interactions slow flow Surface tension ⇒ Energy per unit area of liquid surface 1 ⇒ If "#$% strong, liquid “wants” to “ball up” Molecules minimize surface area where intermolecular interactions are missing ! Chemistry Surface tension is connected to somanycool phenomena Wetting capillary rise soapbubbles waterwalking insects  Lecture #8, p. 17 Water-Walking Insects physicsworld.com Chemistry  Lecture #8, p. 18 Trampolining Droplets T. Schutzius, S. Jung, T, Maitra, G. Graeber, M. Köhme & D. Poulikakos. Spontaneous droplet trampolining on rigid superhydrophobic surfaces. Nature 527, 82 (2015). Chemistry  Lecture #8, p. 19 Properties of Liquids Vapor pressure Above any liquid, some molecules enter gas phase ⇒ Molecules with "& > "#$% If interactions are strong ⇒ Only tiny fraction in gas phase Picture for intuition? ⇒ Velocity distribution iiiiii.in Depends on temperature !!"# Molecules Fraction escaping increases with temperature that can escape liquid Kinetic Energy Chemistry i ofmolecules of molecules  Lecture #8, p. 20 So what is definition of vapor pressure then Properties of Liquids Vapor pressure Consider closed container with liquid at fixed temperature !!"# ≡ pressure of molecules in gas phase equilibrium Weak interactions ⇒ High vapor pressure ⇒ Volatile liquid Tia Dynamic equilibrium ⇒ Molecules leaving/entering gas phase are balanced Iii If T changes ⇒ Equilibrium must be re-established If T increases ⇒ More molecules leave liquid If T decreases ⇒ Excess gas molecules return to liquid Milow Chemistry explained I i e formliquid Theseeffects bychanges inthedistribution of kineticenergy ofthemoleculeswithchanges in temperature Note uchlowerthanforatypicalliquidatthesametemperature III If t.tnsgursiitainesenflge  Lecture #8, p. 21 Phase Changes Another related property where intermolecular interactions play a big role Intermolecular interactions change with phase Energy absorbed or released ∆" > 0 for solid to liquid to gas boiling ∆" < 0 for gas to liquid to solid ∆"!"# ≡ Heat of fusion ∆"$%& ≡ Heat of vaporization ∆"#"' ≡ Heat of sublimation melting with ∆"#"' = ∆"!"# + ∆"$%& and ∆"$%& > ∆"!"# Makes sense! For liquid to gas, all intermolecular interactions are lost For solid to liquid, intermolecular interactions remain significant Chemistry  Lecture #8, p. 22 Tohelpus understand let'sdoanexampleproblem Heating Curves Ex: How much energy is required to heat 1 mol H2O from −25.0°C to 125°C at 101.3 kPa? Molar heat capacities: Ice = 38.1 J mol K −1 −1 Liquid water = 75.3 J mol−1 K−1 Molar Enthalpies ∆"!"# = 6.01 kJ mol−1 ∆"$%& = 40.6 kJ mol−1 F Steam = 37.5 J mol−1 K−1 Enthalpy Changes: A to B = 0.953 kJ Entire Process B to C = 6.01 kJ ∆" = 56.0 kJ C to D = 7.53 kJ D to E = 40.6 kJ E to F = 0.938 kJ Notice how heat curve dominated by vaporization!! Chemistry  Lecture #8, p. 23 Phase Diagram Phase transition temperatures depend on pressure Thus, we often represent phase changes by plotting the phase diagram ! versus " Notes Lines show phase changes Depend on ! and " TP ≡ triple point At TP, three phases co-exist! C ≡ critical point Beyond C, no distinct liquid phase Orange arrow: liquid to gas without ever boiling! Critical point dryer "!" ≡ typically increases with P Tmp 7 Era "#" ≡ typically increases with P Tip Era Chemistry sensel Makes BoilingoccurswhenPrap P LargeEnt night seebelow t.fi ii Seebelow  Lecture #8, p. 24 Explain Boiling Process? Note: we can boil H2O at room temperature Great demo Chemistry  Lecture #8, p. 25 Chemistry  Lecture #8, p. 26 Thevaporpressureof a liquid at a given temperature can bedeterminedusing Calculating Vapor Pressure ln &!"# Clausius–Clapeyron Equation −∆(!"# Plot ti Iii ln $!"# = +, Slope = !∆#!"# )* $ # = $% + ' (Formula for line) # ≡ Gas constant ! ≡ Absolute temperature 0 0 1! $ ≡ Constant # Also says food will cook more slowly in boiling water at higher elevations (%& = 95 °C at 1500 m (%& = 90 °C at 3000 m Chemistry  Lecture #8, p. 27 Becauseoftheirimportance we should sketchphasediagrams forH2oandCO2 Sketches of Important Phase Diagrams H2O CO2 73atm H2O has negative slope for solid–liquid line CO2 has triple point above 1 atm Tmp decreases with increasing pressure Dry ice sublimes! Ice is less dense than liquid H2O Critical point is at mild conditions I Chemistry metathisen cite.it i is.moatmsueisijiitfa aiest  Lecture #8, p. 28 Aerogel source: wikipedia Chemistry li i iiiiiiiiiiii manatees  Lecture #8, p. 29 What We Learned Intermolecular interactions: due to forces between molecules Van der Waals interactions: dispersion and dipole–dipole Molecular dipoles Hydrogen bonding Ion–dipole interactions, polar solvents Relative strengths of intermolecular interactions Properties of liquids: viscosity, surface tension, vapor pressure Phase changes Heat of fusion, heat of vaporization, heat of sublimation Heating curves Phase diagrams, critical point, triple point Clausius–Clapeyron equation Chemistry