Liquids and Solids Intermolecular forces Chapter objectives Physical properties of liquids: Vapor pressure and boiling Intermolecular forces 2 Physical Property: Interactions Between Molecules Many of the phenomena we observe are related to interactions between molecules that do not involve a chemical reaction your taste and smell organs work because
molecules interact with the receptor molecule sites in your tongue and nose 3 Why is Sugar a Solid But Water is a Liquid? The state a material exists in depends on the attraction between molecules and their ability to overcome the attraction The attractive forces between Ions or Molecules Their structure the attractions are electrostatic depend on shape, polarity, etc. The ability of the molecules to overcome the
attraction Kinetic energy they possess 4 Forces of Attraction within a Liquid Cohesive Forces = forces that try to hold the liquid molecules to each other MoleculeMolecule surface tension Adhesive Forces = forces that bind a substance to a surface MoleculeSurface capillary action meniscus
5 Surface Tension Surface tension: the tendency of liquids to minimize their surface. Cause: intermolecular force liquids to have a surface that resists penetration Paper clip (denser than wate r) can float on water 6 Viscosity some liquids flow more easily than
others: Soda more fluidy than Syrup Viscosity : the resistance of a liquid to flow. Syrup is more viscous than Soda Attractive forces between the molecules (intermolecular forces) 7 Evaporation and Vapor Evaporation : molecules of a liquid breaking free from the surface: Liquid Gas also known as vaporization Physical change
Vapor: gaseous form from liquid 8 Evaporation: Liquid Molecule escape into Gas Evaporation happens at the surface molecules on the Surface experience a smaller net attractive force than molecules in the Interior NOT all the surface molecules escape at once: only the ones with sufficient kinetic energy (fast enough) to overcome the attractions will escape 9 Condensation: Gas Liquid
Condensation : the vapor molecules (gas state) may bump into and stick to the surface of the container or get recaptured by the liquid. Physical change : Gas Liquid Condensation occurs for molecules with less kinetic energy and/or collides to surface 10 Evaporation vs. Condensation: Dynamic Equilibrium Evaporation and Condensation are opposite processes In a sealed container, eventually, the rate of
evaporation and condensation in the container will be the same Rate evaporation = Ratecondensation Dynamic equilibrium : opposite processes that occur at the same rate in the same system The amount of vapor vs. liquid appears constant 11 Evaporation Condensation Water is just added to the flask and it is capped, all the water
molecules are in the liquid. Shortly, the water starts to evaporate. Speed of evaporation >> Speed of condensation (Rateevap >> Ratecondsn) Eventually, Rateevap = Ratecondsn The air in the flask is now saturated with water vapor. 12
Vapor Pressure Pvap once equilibrium is reached, then the amount of vapor (mole of vapor, nvap) in the container will remain the same as long as you dont change the conditions Vapor pressure: the partial pressure exerted by the vapor of the liquid. Ideal Gas Law : Pvap nvap R T V Depending on the temperature and strength of intermolecular attractions
13 Vapor Pressure increases as temperature increases ether ethanol normal boiling point water 14 Boiling and Boiling Point (b.p.) Boiling: vapor pressure of the liquid is the same as the atmospheric pressure. Pvap = Pair
Rapid evaporation Boiling point: the temperature for boiling process normal boiling point: temperature when Pair = 1 atm b.p. of water is 100C b.p. depends on Pair the temperature of boiling water on the top of a mountain will be cooler than boiling water at sea level On top of Mount Whitney, b.p. of water is about 84C 15 Vapor pressure at given temperature vs. Normal Boiling point At the same temperature, different liquids have different vapor pressure (volatility) Volatile: Liquids having high vapor pressure
Liquids having higher vapor pressure will have lower normal boiling points 16 Energy flow: Evaporation vs. Condensation Evaporation: Liquid _______ heat from its surroundings to evaporate The surroundings cool off Endothermic: heat flows into a system from the surroundings as alcohol evaporates off your skin, it cools your skin Condensation: Gas _______ heat to its surroundings to reduce its temperature The surroundings warms up
Exothermic: heat flows out of a system into the surroundings 17 Heat of Vaporization Definition: the amount of heat needed to vaporize one mole of a liquid. Hvap Liquid Gas it requires 40.7 kJ of heat to vaporize one mole of water at 100C _____thermic since Condensation (Gas Liquid + Hvap ) is the opposite process to evaporation, ____thermic. Hcond = -Hvap
Hvap is a Conversion Factor (unit: kJ/mol). 18 Example: Calculate the amount of heat is required to vaporize 100. g water at its boiling point. Molar heat of vaporization = 47.01 kJ/mol g H2O Given: 100. g water Find: heat CF: 40.7 kJ = 1 mol; 18.02 g = 1 mol
mol H2O 5.56 mole H2O heat 261 kJ 19 Temperature and Melting For solid, temperature increases until it reaches the melting point. Ice melts at 0C. During melting: the temperature remains _________ until it all turns
to a liquid. solid liquid Why temperature remains constant? All the added heat is for overcoming the attractive forces in the solid, not increase the temperature 20 Energy of Melting and Freezing Melting: Solid absorbs heat from its surroundings: _____thermic as Heat flows out of the surroundings the surroundings cool off as ice in your drink melts, it cause the liquid to cool Freezing: Liquid releases heat into its surroundings: _____thermic
as heat flows into the surroundings the surroundings warm up 21 Heat of Fusion Definition: heat needed to melt one mole of a solid Hfus since freezing (crystallization) is the opposite process to melting, the same amount of energy transferred is the same, but in the opposite direction Hcrystal = -fusion Heat of fusion (kJ/mol) can be used as conversion factor to calculate heat in melting/freezing problems 22
Heats of Fusion of Several Substances Melting Point, C 0.00 Hfusion, (kJ/mol) 6.02 Liquid water Chemical Formula H2O
diethyl ether C4H10O -116.3 7.27 23 Sublimation vs. Deposition Sublimation: the Solid form changes directly to the Gaseous form. Solid Gas without going through the liquid form Dry ice (solid CO2) gas CO2
like melting, sublimation is endothermic Deposition is the reverse of Sublimation, exothermic. 24 Heating Curve: phase changes during heating solid ice at 1 atm Temperature (oC) Temperature of water at Constant Heating 160 140
120 100 80 60 40 20 0 -20 s s+l g l 0
10 20 30 40 50 l+g 60 70
80 90 100 Time 25 Intermolecular Forces (IMF) affects physical properties of solid & liquid stronger intermolecular forces increases surface tension and viscosity
stronger intermolecular forces reduces vapor pressure (retaining molecules in liquid state), thus increases boiling point. Stronger IMF also increases melting point 26 Why are molecules attracted to each other? Attractive forces between opposite electric charges Ionic: + ion to ion: NaCl Molecular: (+) end of polar molecule to (-) end of polar
molecule. HOd-HdOd-Hd Ionic or Molecular: larger charge stronger attraction: MgCl > NaCl Fd-HdCld-Hd > Cld-HdCld-Hd How about nonpolar molecules? 27 Intermolecular forces in Pure liquids Dispersion force (aka London force) Dipole-dipole force Hydrogen bonding 28
Dispersion Forces also: London Forces or Induced Dipoles Electrons on one molecule distorting the electron cloud on another ALL molecules have Dispersion Forces Dispersion force is especially important among nonpolar molecules - - + - - - -
Somewhat polar Nonpolar Polar 30 Strength of Dispersion Force More electrons (more molar mass): Electrons can move more easily within a molecule: =O < =S -F < -Cl < -I Example: CF4 (b.p. -127C) vs. CBr4 (b.p. 190C) Larger molecules.
Example: CH4 (b.p. -161C) vs. C2H6 (b.p. -89C) Why? more electrons + electron farther from the nuclei the larger the dipole that can be induced 31 Practice: Which has higher boiling point? 1. CH4 vs. CCl4 2. F2 vs. Br2 3. Liquid argon vs. liquid xenon Hint: Compare molecular polarity (use VSEPR theory)
the number of electrons in each molecule/atom 32 Permanent Dipoles Recall molecular polarity: Electronegativity difference & Molecular Geometry some molecules have a Permanent Dipole: (+) (-) all polar molecules have a permanent dipole. H2O, NH3, HCl, etc. 33
Dipole-to-Dipole Attraction Polar molecules have a permanent dipole a + end and a end the + end of one molecule will be attracted to the end of another 34 Attractive Forces Dispersion Forces all molecules + +
+ - + - 35 Why iodine chloride (ICl, molar mass = 162, b.p. 97 C) has higher boiling point than bromine (Br2, molar mass = 160, b.p. 59 C)? Dispersion force: ICl (70 e-) vs. Br2 (70 e-) Dipole moments: ICl (polar molec.) vs. Br2
(nonpolar molec.) Hint: compare the number of electrons in each molecule/atom molecular polarity (use VSEPR theory) 36 Hydrogen Bonding Molecules that have HF, -OH or NH groups have particularly strong intermolecular attractions unusually high melting and boiling points unusually high solubility in
water Not for all molecules with hydrogen atom Hydrogen Bond 37 Intermolecular H-Bonding 38 Cause of Hydrogen Bonding A very electronegative atom X (X = F, O, N) is bonded to hydrogen, the bonding electrons is pulled toward X.
Xd-Hd Since hydrogen has no other electrons, the nucleus becomes deshielded (stripped): -Hd exposing the proton The exposed proton Hd (center of positive charge) attracting all the electron clouds from neighboring molecules Xd-HdYd- 39 H-Bonds vs. Chemical Bonds Hydrogen bonds are not chemical bonds Hydrogen bonds are attractive forces between molecules Chemical bonds are attractive forces that
make molecules 40 Hydrogen Bond in DNA double helix 41 Types of Intermolecular Forces Type of Force Relative Strength
Present in Example weak, but all atoms Dispersion increases H2 and Force with molar molecules mass Dipole
Dipole Force Hydrogen Bond moderate only polar HCl molecules strong molecules having H
HF bonded to F, O or N 42 Why dimethyl ether (CH OCH , b.p. 249 K) has 3 3 much lower boiling point than ethanol (CH3CH2OH, b.p. 351 K)? Dispersion force: Hint: compare CH3OCH3 (20 e-) vs. CH3CH2OH
the number of (20 e-) electrons in each molecule/atom Dipole moments: molecular polarity CH3OCH3 (1.30 D) < (use VSEPR CH3CH2OH (1.63 D) theory) Hydrogen Hydrogen bonding: bonding? CH3OCH3 (absent) vs. CH3CH2OH (present) 43
Attractive Forces and Solubility Like dissolves Like miscible = liquids that do not separate Polar molecules dissolve in Polar solvents water, alcohol, isopropanol, CH2Cl2 H-bond: molecules with O or N higher solubility in H 2O Nonpolar molecules dissolve in nonpolar solvents ligroin (hexane), toluene, kerosene, CCl4 if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition 44
Solubility between two liquids: Immiscible Liquids Pentane (C5H12) (C-H and C-C bond, nonpolar substance) is mixed with water (O-H bond, polar) the two liquids separate they are more attracted to their own kind of molecule than to the other. 45 Types of Crystalline Solids
46 Molecular Crystalline Solids Molecular solid: composite units are molecules. CO2 CO2 H2O H2O H2O Held together by intermolecular attractive forces dispersion, dipole-dipole, or H-bonding generally low melting points and Hfusion 47
Ionic Crystalline Solids Ionic solids: composite units are formula units. NaCl NaClNaCl Held together by Electrostatic forces between Cation and Anion cations and anions arranged in a geometric pattern called a crystal lattice to maximize attractions generally higher melting points and Hfusion than molecular solids because ionic bonds are stronger than intermolecular forces
48 Atomic Crystalline Solids Atomic solids: composite units are individual atoms XeXe XeXe Held together by either covalent bonds, dispersion forces or metallic bonds melting points and Hfusion vary depending on the attractive forces between the atoms 49
Types of Atomic Solids 50 Types of Atomic Solids Covalent Covalent Atomic Solids : atoms attached by covalent bonds. Diamond Carbon (tetrahedral, C-C bond). effectively, the entire solid is one, giant molecule Covalent bonds are strong very High melting points and Hfusion High hardness
51 Types of Atomic Solids Nonbonding Nonbonding Atomic Solid: held together by dispersion forces XeXe XeXe Dispersion forces are relatively weak, very low melting points and Hfusion 52 Types of Atomic Solids Metallic
Metallic solids: held together by metallic bonds How: metal atoms release some of their electrons to be shared by all the other atoms in the crystal Metallic bond: the attraction of the metal Cations M+ for the mobile electrons e often described as islands of cations in a sea of electrons 53 Metallic Bonding Model of metallic bonding explain: luster, malleability, ductility, electrical and thermal conductivity the mobility of
the electrons in the solid the strength of the metallic bond Charge and Size of the cations so the melting points and Hfusion of metals vary as well 54 Water: A Unique and Important Substance found in all 3 states on the Earth: Ice, Liquid, Vapor the most common solvent (liquid) found in nature: seawater as largest sample of solution. without water, life as we know it could not exist the search for extraterrestrial life starts with the search for water
relatively high boiling point: mostly as liquid expands as it freezes most substances contract as they freeze causes ice to be less dense than liquid water 55 Practice: Which of the following pairs of substances does the first one have higher boiling point? a) b) c) d) water vs. hydrogen sulfide
sulfur dioxide vs. carbon dioxide liquid oxygen vs. liquid nitrogen CH3CH2OH vs. CH3OCH3 e) CH3CH2OH vs. CH3CH2SH f) hydrogen chloride vs. hydrogen fluoride g) methane vs. carbon tetrafluoride 56 Practice: Which of the following pairs of substances does the first one have higher boiling point? a) b) c) d)
water vs. hydrogen sulfide sulfur dioxide vs. carbon dioxide liquid oxygen vs. liquid nitrogen CH3CH2OH vs. CH3OCH3 e) CH3CH2OH vs. CH3CH2SH H-bond Dipole-dipole Dispersion H-bond H-bond f) hydrogen chloride vs. hydrogen fluoride g) methane vs. carbon tetrafluoride Dispersion
H-bond 57 Practice: Which of the following pairs of substances is the first one more volatile? a) b) c) d) water vs. hydrogen sulfide sulfur dioxide vs. carbon dioxide liquid oxygen vs. liquid nitrogen
CH3CH2OH vs. CH3OCH3 e) CH3CH2OH vs. CH3CH2SH f) hydrogen chloride vs. hydrogen fluoride g) methane vs. carbon tetrafluoride 58 Practice: Which is Which? Rock candy (crystalline sugar) Gold nugget Copper(I) sulfide solid Diamond
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