Liquids & Aqueous solutions

Ideal Gas Equation R = 0.08206 atm L/mol K Ideal gas equation predicts gas behavior

A Model for Gas Behavior Ideal gas law describes what gases do, but not why. Kinetic Molecular Theory of Gases (KMT): model that explains gas behavior. developed in mid-1800s based on concept of an ideal or perfect gas

Ideal gas Tiny particles in constant, random, straight-line motion Molecules collide w/ each other & w/ walls of container Gas molecules are points; gas volume is empty space between molecules Molecules independent of each other (no attractive or repulsive forces between them).

Molecular motion & temperature Moving molecule has kinetic energy KE depends on mass (m) and speed (u) Temperature (in K) proportional to average molecular KE At higher T, average speed higher At lower T, average speed lower At T = 0, speed = 0 (molecules stop moving)

Different gases at same temperature All have same average KE (same temperature) Heavier gases are slower; lighter gases are faster

Molecular motion and pressure Molecules colliding with container → gas pressure What if there are more molecules? More collisions → higher pressure

Molecular motion and pressure Molecules colliding with container → gas pressure What if the container is smaller? More collisions → higher pressure

Molecular motion and pressure Molecules colliding with container → gas pressure What if the molecules are moving faster? Harder, more frequent collisions → higher pressure

Molecular motion and volume Moving molecules fill the container Light molecules escape faster, heavy molecules more slowly Large spaces between molecules allow gas to be compressed

KMT & liquids Ideal gas remains a gas when cooled, even to 0 K Real gases condense to liquid state when cooled How do we explain condensation? Pressure (atm) Temperature (K) Ideal gas pressure decreases steadily & becomes zero at absolute zero Real gas pressure decreases abruptly to zero when gas condenses to liquid

Condensation KMT ignores attractions between gas molecules Gas molecules are too far apart & too fast for attractions to act BUT . . . attractive forces do exist between all molecules! At low enough T, attractions overcome kinetic energy & molecules stick together to form a liquid

Vaporization At liquid surface, faster molecules have enough kinetic energy to escape (evaporate) As higher-energy molecules leave liquid, average kinetic energy of liquid decreases The temperature of liquid decreases (evaporative cooling)

What if the vapor can’t escape? (a) Molecules evaporate (b) Some vapor molecules return to liquid (c) Rates of evaporation & condensation equal When rate of vaporization = rate of condensation, system has reached dynamic equilibrium vapor pressure at equilibrium is constant vp is balance of KE (temperature) & attractions vp affected only by temperature liquids with high vp at room temp are volatile

Vapor pressure always increases as temperature increases Vapor pressure curves Vapor pressure always increases as temperature increases

Evaporation & boiling

Evaporation & boiling Evaporation occurs only at surface As temperature increases, evaporation increases At some point, evaporation begins to occur throughout the liquid instead of just at the surface: boiling! Vapor bubbles form throughout the liquid Bubbles rise to surface, burst, and release vapor

Boiling point Boiling begins when liquid’s vapor pressure matches atmospheric pressure Temperature at which this occurs is boiling point b.p. at standard pressure is normal boiling point When atmospheric pressure is lower, bp is lower When atmospheric pressure is higher, bp is higher

Intermolecular attractions Attractive forces exist between all atoms/molecules Strength of attractions indicated by boiling point When comparing two substances, Low b.p. ⇒ weaker intermolecular attractions High b.p. ⇒ stronger intermolecular attractions

Trends in boiling points

Intermolecular attractions For molecules of similar structure, boiling point increases as molar mass increases intermolecular attractive forces increase as molar mass increases

Then there’s hydrogen . . .

The O–H bond in water is very polar, and the atoms are very small The dipoles are close together, so their attraction is very strong An H atom is covalently bonded (red-white) to its own O and weakly bonded (dotted line) to the neighboring O This weak bond to a neighboring O is called a hydrogen bond

Hydrogen bonding Hydrogen bonding occurs only between molecules containing N–H, O–H, and F–H bonds Hydrogen bonding is much stronger than ordinary intermolecular attractions ⇒ very high boiling points for their mass Hydrogen bonds are not as strong as covalent bonds (15-40 kJ/mol, vs >150 kJ/mol)

Boiling point & attractions CH3-CH2-CH3 CH3-O-CH3 CH3-CH2-NH2 CH3-CH2-OH molar mass 44 g/mol 45 g/mol 46 g/mol normal bp 231 K 248 K 290 K 352 K Similar masses⇒ predict similar boiling points 1st molecule is nonpolar 2nd molecule is polar, a little stickier: higher bp 3rd & 4th molecules have hydrogen bonding between molecules, much stickier: much higher bp

Heating curve Add energy Temperature

Heating curve Temperature (g) boiling condensing (l) melting freezing Add energy Temperature (s) melting freezing boiling condensing (l) (g) point melting/freezing point

Heating curve Melting & boiling are ENDOTHERMIC Freezing & condensing are EXOTHERMIC

Changing temperature changes KE (#1, 3, 5) Changing state changes potential energy (#2, 4) energy to melt or freeze = heat of fusion (∆Hfusion) energy to vaporize or condense = heat of vaporization (∆Hvaporization)

Energy in phase changes