Intermolecular forces Liquids and Solids

Transcription

1 Intermolecular forces Liquids and Solids

2 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 2

3 Intermolecular Attractive Forces

4 Intermolecular forces affect physical properties of solid and liquid Stronger intermolecular force in the liquid prevent liquid molecules from escaping into gas state, causing lower vapor pressure Recall: substance with lower vapor pressure will have (higher, lower) boiling point. 4

5 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 5

6 Intermolecular forces in Pure liquids Dispersion force (aka London force) Dipole-dipole force Hydrogen bonding 6

7 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

8 Dispersion Forces: Instantaneous Dipoles Somewhat polar Nonpolar Polar 8

9 Dispersion Force: Strength 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 strength of the dispersion force gets Larger with larger molecules 9

10 Permanent Dipoles Chapter 4: Electronegativity difference & Molecular Geometry some molecules have a Permanent Dipole: (+) (-) all polar molecules have a permanent dipole. H 2 O, NH 3, HCl, etc. 10

11 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 11

12 Attractive Forces Dispersion Forces all molecules Dipole-to-Dipole Forces polar molecules

13 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 13

14 Intermolecular H-Bonding 14

15 Hydrogen Bonding A very electronegative atom X (X = F, O, N) is bonded to hydrogen, the bonding electrons is pulled toward X. X δ -H δ+ Since hydrogen has no other electrons, the nucleus becomes deshielded ( stripped ): -H δ+ exposing the proton The exposed proton H δ+ (center of positive charge) attracting all the electron clouds from neighboring molecules X δ -H δ+ Y δ - 15

16 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 16

17 Hydrogen Bond in DNA double helix 17

18 Types of Intermolecular Forces Type of Force Relative Strength Present in Example Dispersion Force weak, but increases with molar mass all atoms and molecules H 2 Dipole Dipole Force moderate only polar molecules HCl Hydrogen Bond strong molecules having H bonded to F, O or N HF 18

19 Attractive Forces and Solubility Like dissolves Like miscible = liquids that do not separate Polar molecules dissolve in Polar solvents water, alcohol, isopropanol, CH 2 Cl 2 H-bond: molecules with O or N higher solubility in H 2 O Nonpolar molecules dissolve in nonpolar solvents ligroin (hexane), toluene, kerosene, CCl 4 if molecule has both polar & nonpolar parts, then hydrophilic - hydrophobic competition 19

20 Solubility between two liquids: Immiscible Liquids Pentane (C 5 H 12 ) (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. 20

21 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 21

22 Structure Determines Properties: Solids, Liquid and Gases 22

23 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 23

24 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 24

25 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 25

26 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 26

27 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 27

28 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 (Rate evap >> Rate condsn ) Eventually, Rate evap = Rate condsn The air in the flask is now saturated with water vapor. 28

29 Vapor Pressure P vap once equilibrium is reached, then the amount of vapor (mole n vap ) in the container will remain the same as long as you don t change the conditions Ideal Gas Law : Vapor pressure: the partial pressure exerted by the vapor of the liquid. Depending on the temperature and strength of intermolecular attractions P vap = n vap R V 29 T

30 Vapor Pressure increases as temperature increases ether ethanol normal boiling point water 30

31 Boiling and Boiling Point (b.p.) Boiling: vapor pressure of the liquid is the same as the atmospheric pressure. Liquid Gas. P vap = P air Boiling point: the temperature for boiling process normal boiling point: temperature when P air = 1 atm b.p. of water is 100 C b.p. depends on P air 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 31

32 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 32

33 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 33

34 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 34

35 Sublimation vs. Deposition Sublimation: the Solid form changes directly to the Gaseous form. Solid Gas without going through the liquid form Dry ice (solid CO 2 ) gas CO 2 like melting, sublimation is endothermic Deposition is the reverse of Sublimation, exothermic. 35

36 Heating Curve: phase changes during heating solid ice at 1 atm Temperature of water at Constant Heating Temperature ( o C) s s+l g l l+g Time 36

37 Types of Crystalline Solids 37

38 Molecular Crystalline Solids Molecular solid: composite units are molecules. CO 2 CO 2 H 2 O H 2 O H 2 O Held together by intermolecular attractive forces dispersion, dipole-dipole, or H-bonding generally low melting points and ΔH fusion 38

39 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 ΔH fusion than molecular solids because ionic bonds are stronger than intermolecular forces 39

40 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 ΔH fusion vary depending on the attractive forces between the atoms 40

41 Types of Atomic Solids 41

42 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 ΔH fusion High hardness 42

43 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 ΔH fusion 43

44 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 44

45 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 45