Intermolecular forces Liquids and Solids

Chapter objectives Understand the three intermolecular forces in pure liquid in relation to molecular structure/polarity Understand the physical properties of liquids that is relevant to intermolecular force: vapor pressure and boiling

How Gecko can climb upside down on the ceiling? Gecko are capable of climbing vertical or even upsidedown, utilizing micrometer size hairs to adhere to surfaces: Intermolecular force. New research showed gecko can turn on and off the “stickiness”.

Intermolecular forces affect 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 likewise with melting point

Surface Tension Surface tension: Tendency of liquid to minimize total surface. Cause: Molecules on the surface attract to each other. Paperclip (made of iron, density >> water) should sink in water.

Viscosity Viscosity: Internal friction within the liquid. Stronger intermolecular forces increases viscosity Viscous liquid takes longer to flow. Honey has higher viscosity than water Used motor oil has less viscosity

Intermolecular forces affect physical properties of solid and liquid Stronger intermolecular force in the liquid prevent liquid molecules from escaping into gas state  Molecules need more heat energy to escape from the liquid (higher temperature)  Leading to ________ (higher, lower) boiling point.

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

Permanent Dipoles Chapter 4: 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 Similar to attraction between two separated bar magnets

Polar Molecules Can be attracted by Charge Demo: https://www.youtube.com/watch?v=jkYz1WlpRSQ Note: Only works for polar liquid. Nonpolar liquid, such as gasoline, doesn’t work. When two polar molecules are coming close, there are both attraction and repulsion between them. Since the Coulombic force depends on the distance between the charge, the attraction between two polar molecules is STRONGER than repulsion, leading to intermolecular attractive force.

Dispersion Forces Nonpolar molecules also attract each other: London Forces or Induced Dipoles Cause: 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

Dispersion Force: Strength Strength of the dispersion force gets Larger with larger molecules Electron mobility: how easily the electrons can move within a molecule, or be polarized. =O < =S -F < -Cl < -I more electrons + electron farther from the nuclei  the larger the dipole that can be induced

Dispersion Force: Nonpolar molecules F2, Cl2, Br2, I2 Larger molar mass, stronger intermolecular forces leads to higher boiling point, higher melting point.

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

Beyond Dipole-Dipole and London Dispersion forces? London force: high molar mass, stronger force, higher boiling point? Anomaly for water and HF

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  Hydrogen Bond

Intermolecular H-Bonding

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 Bonding 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

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 Gasoline, Paint thinner, toluene, kerosene, CCl4 if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition

Solubility between two liquids: Miscible vs. Immiscible Water and alcohol can mix at any ratio Immiscible: Pentane (C5H12) (C-H and C-C bond, nonpolar) is mixed with water (O-H bond, polar) the two liquids separate

Interaction Between Molecules Affects Physical Property 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 The volatility of liquids depends on the IMF.

Structure Determines Properties: Solids, Liquid and Gases

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

Escaping from the Surface Evaporation : molecules of a liquid breaking free from the surface: Liquid  Gas also known as vaporization Physical change a substance is converted from its liquid form to its gaseous form the gaseous form is called a vapor

Evaporation: Liquid  Gas Molecules of the liquid mix with and dissolve in the air happens at the surface molecules on the Surface experience a smaller net attractive force than molecules in the Interior but all the surface molecules do not escape at once, only the ones with sufficient kinetic energy to overcome the attractions will escape

Condensation: Gas  Liquid in a closed container, after a liquid evaporates, the vapor molecules are trapped and may eventually turn into liquid Condensation : the vapor molecules may eventually bump into and stick to the surface of the container or get recaptured by the liquid. Physical change : Gas  Liquid

Dynamic Equilibrium Evaporation and Condensation are opposite processes eventually, the rate of evaporation and condensation in the container will be the same Dynamic equilibrium : opposite processes that occur at the same rate in the same system

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 nvap) in the container will remain the same Vapor pressure: Pressure exerted by the vapor of the liquid when equilibrium is reached between liquid and gas states. Depending on the temperature and strength of intermolecular forces (IMF): Stronger IMF, more ____________ for liquid to become vapor, ___________ (higher, lower) Pvap.

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. Liquid  Gas. Pvap = Pair 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) Liquids having higher vapor pressure are normally called “more volatile” Liquids having higher vapor pressure will have lower normal boiling points

Energy flow: Evaporation vs. Condensation Evaporation: Liquid absorbs 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 causes your skin to cool Condensation: Gas releases heat to its surroundings to reduce its temperature  The surroundings warms up Exothermic: heat flows out of a system into the surroundings

Temperature and Melting For solid, temperature increases until it reaches the melting point. Ice melts at 0°C. During melting: the temperature remains the same until it all turns to a liquid. solid  liquid all the Energy from the heat source is for overcoming the attractive forces in the solid, not increase the temperature

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

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.

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– 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. Xenon solid (at low temperature) 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”

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 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 expands as it freezes most substances contract as they freeze causes ice to be less dense than liquid water