Intermolecular forces Liquids and Solids Intermolecular forces

Chapter objectives Physical properties of liquids: Vapor pressure and boiling Intermolecular forces

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

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

Forces of Attraction within a Liquid Cohesive Forces = forces that try to hold the liquid molecules to each other Molecule  Molecule surface tension Adhesive Forces = forces that bind a substance to a surface Molecule  Surface capillary action meniscus

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 water) can float on water

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)

Evaporation and Vapor Vapor: gaseous form from liquid Evaporation : molecules of a liquid breaking free from the surface: Liquid  Gas also known as vaporization Physical change Vapor: gaseous form from liquid

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

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

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

Evaporation  Condensation Water is just added to the flask and it is capped, all the water molecules are in the liquid. Eventually, Rateevap = Ratecondsn The air in the flask is now saturated with water vapor. Shortly, the water starts to evaporate. Speed of evaporation >> Speed of condensation (Rateevap >> Ratecondsn)

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 don’t change the conditions Vapor pressure: the partial pressure exerted by the vapor of the liquid. Depending on the temperature and strength of intermolecular attractions

Vapor Pressure increases as temperature increases ethanol ether normal boiling point water

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 100°C 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 84°C

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

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

Heat of Vaporization Definition: the amount of heat needed to vaporize one mole of a liquid. DHvap + Liquid  Gas it requires 40.7 kJ of heat to vaporize one mole of water at 100°C _____thermic since Condensation (Gas  Liquid + DHvap ) is the opposite process to evaporation, ____thermic. DHcond = -DHvap DHvap is a Conversion Factor (unit: kJ/mol).

g H2O  mol H2O  heat 5.56 mole H2O 261 kJ 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 Given: 100. g water Find: heat CF: 40.7 kJ = 1 mol; 18.02 g = 1 mol g H2O  mol H2O  heat 5.56 mole H2O 261 kJ

Temperature and Melting For solid, temperature increases until it reaches the melting point. Ice melts at 0°C. 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

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

Heat of Fusion Definition: heat needed to melt one mole of a solid DHfus since freezing (crystallization) is the opposite process to melting, the same amount of energy transferred is the same, but in the opposite direction DHcrystal = -Dhfusion Heat of fusion (kJ/mol) can be used as conversion factor to calculate heat in melting/freezing problems

Heats of Fusion of Several Substances Liquid Chemical Formula Melting Point, °C DHfusion, (kJ/mol) water H2O 0.00 6.02 isopropyl alcohol C3H7OH -89.5 5.37 acetone C3H6O -94.8 5.69 diethyl ether C4H10O -116.3 7.27

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.

Heating Curve: phase changes during heating solid ice at 1 atm s+l s g l l+g

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

Why are molecules attracted to each other?  Attractive forces between opposite electric charges Ionic: + ion to – ion: Na+ Cl– Molecular: (+) end of polar molecule to (-) end of polar molecule. HOd–-Hd+ Od–-Hd+ Ionic or Molecular: larger charge  stronger attraction: Mg2+ Cl– > Na+ Cl– Fd–-Hd+  Cld–-Hd+ > Cld–-Hd+  Cld–-Hd+ How about nonpolar molecules?

Intermolecular forces in Pure liquids Dispersion force (aka London force) Dipole-dipole force Hydrogen bonding

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 + - + - + -

Dispersion Forces: Instantaneous Dipoles Nonpolar Somewhat polar Polar

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. -127°C) vs. CBr4 (b.p. 190°C) Larger molecules. Example: CH4 (b.p. -161°C) vs. C2H6 (b.p. -89°C) Why? more electrons + electron farther from the nuclei  the larger the dipole that can be induced

Practice: Which has higher boiling point? Hint: Compare molecular polarity (use VSEPR theory) the number of electrons in each molecule/atom CH4 vs. CCl4 F2 vs. Br2 Liquid argon vs. liquid xenon

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.

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

Attractive Forces Dispersion Forces – all molecules _ + Dipole-to-Dipole Forces – polar molecules + - + - + - + -

Why iodine chloride (ICl, molar mass = 162, b. p 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)

Hydrogen Bonding  Hydrogen Bond 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

Intermolecular H-Bonding

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–-Hd+  Yd–-

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

Hydrogen Bond in DNA double helix

Types of Intermolecular Forces Type of Force Relative Strength Present in Example DispersionForce weak, but increases with molar mass all atoms and molecules H2 Dipole – Dipole Force moderate only polar molecules HCl Hydrogen Bond strong molecules having H bonded to F, O or N HF

Why dimethyl ether (CH3OCH3, b. p Why dimethyl ether (CH3OCH3, b.p. 249 K) has much lower boiling point than ethanol (CH3CH2OH, b.p. 351 K)? Dispersion force: CH3OCH3 (20 e-) vs. CH3CH2OH (20 e-) Dipole moments: CH3OCH3 (1.30 D) < CH3CH2OH (1.63 D) Hydrogen bonding: CH3OCH3 (absent) vs. CH3CH2OH (present) Hint: compare the number of electrons in each molecule/atom molecular polarity (use VSEPR theory) Hydrogen bonding?

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 H2O Nonpolar molecules dissolve in nonpolar solvents ligroin (hexane), toluene, kerosene, CCl4 if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition

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.

Types of Crystalline Solids

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 DHfusion

Ionic Crystalline Solids Ionic solids: composite units are formula units. NaCl Na+  Cl– Na+  Cl– 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 DHfusion than molecular solids because ionic bonds are stronger than intermolecular forces

Atomic Crystalline Solids Atomic solids: composite units are individual atoms Xe  Xe  Xe  Xe Held together by either covalent bonds, dispersion forces or metallic bonds melting points and DHfusion vary depending on the attractive forces between the atoms

Types of Atomic Solids

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 DHfusion  High hardness

Types of Atomic Solids Nonbonding Nonbonding Atomic Solid: held together by dispersion forces Xe  Xe  Xe  Xe Dispersion forces are relatively weak,  very low melting points and DHfusion

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

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 DHfusion of metals vary as well

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

Practice: Which of the following pairs of substances does the first one have higher boiling point? water vs. hydrogen sulfide sulfur dioxide vs. carbon dioxide liquid oxygen vs. liquid nitrogen CH3CH2OH vs. CH3OCH3 CH3CH2OH vs. CH3CH2SH hydrogen chloride vs. hydrogen fluoride methane vs. carbon tetrafluoride

Practice: Which of the following pairs of substances does the first one have higher boiling point? water vs. hydrogen sulfide sulfur dioxide vs. carbon dioxide liquid oxygen vs. liquid nitrogen CH3CH2OH vs. CH3OCH3 CH3CH2OH vs. CH3CH2SH hydrogen chloride vs. hydrogen fluoride methane vs. carbon tetrafluoride H-bond Dipole-dipole Dispersion H-bond H-bond H-bond Dispersion

Practice: Which of the following pairs of substances is the first one more volatile? water vs. hydrogen sulfide sulfur dioxide vs. carbon dioxide liquid oxygen vs. liquid nitrogen CH3CH2OH vs. CH3OCH3 CH3CH2OH vs. CH3CH2SH hydrogen chloride vs. hydrogen fluoride methane vs. carbon tetrafluoride

Practice: Which is Which? Rock candy (crystalline sugar) Gold nugget Copper(I) sulfide solid Diamond Atomic solid Molecular solid Ionic solid Covalent solid