chapter 14 liquid and solids

chapter 14liquid and solids

what is the deal with liquids and solids? why do they do the things they do? • unlike gases, why do they cling? and why are they so dense? and why do solids keep their shape? • and, it takes a little E to melt things (sl) but a lot of E to boil (lg); why???

liquids and solids are not too much different from each other • when a liquid becomes a solid it becomes denser (except water!)

14.1 intermolecular forces • but why is water a liquid? • intermolecular forces! forces between - not within - molecules • there are three types…

polar molecules get together b/c of mutual attraction • the ∂+ of one is attracted to the ∂- of the other

they’ll find the best orientation, too, for maximum “stickiness” • called dipole-dipole attraction • only about 1% as strong as covalent

hydrogen bonding • the HeMan of intermolecular forces, a special breed of dipole-dipole, is hydrogen bonding • guess which element plays a big role here? • hydrogen! but only when hooked up to certain other atoms…

when H is hooked up to N, O, or F poor little H is so small that it nearly loses its electron • this makes it an extra strong ∂+ • this makes an really strong dipole-dipole…

then the water molecules, in this case, have a strong attraction for each other • their attraction is called an H-bond • represented by------ lines • [not true bonds, just really strong attractions]

happens also with molecules which have N-H and F-H bonds • above: water, ethanol, ammonia are all H-bonding • it’s the reason they all are liquids when they aren’t suppose to be (i.e. their friends w/o H-bonding are all gases)

it has a great affect on the properties of those H-bonded things by making them cling to each other stronger • so, e.g. boiling points shoot through the roof b/c it is so hard to separate them

londondispersionforces • what about nonpolar substances? how can they be attracted to each other and form a liquid? • they are attracted by london dispersion forces • some atoms, although nonpolar overall can experience temporary, instantaneous dipoles…

this is unbelievably short-lived and extremely weak but it adds up (like Velcro!)

can happen to nonpolar molecules, too • they have to be slooow and close for any significant effect, but… • large critters (e.g. I2) can do it more easily b/c of all the electrons involved

14.2 water and its phase changes • dang if water ain’t… like… everywhere! • oceans, land, atmosphere, life… • it’s colorless, odorless, tasteless • on Earth it is a solid, liquid, & gas

one can heat water and its T rises (duh.) • but when its T reaches 100˚C its T rise stops! • when 1 atm of air is pressing on the water @ 100˚C, it begins to boil • = normal boiling point • what’s this got to do with i/m forces?patience!

this heating thing can be shown on a heating/cooling curve

notice the water follows a left to right path when it is heated • when it is cooled, it goes right to left

notice when cooling that as the water reaches 0˚C its T levels out again (called a plateau) • this is normal freezing point • ice and liquid can co-exist at this T • but take the Tbelow O˚C,it freezes • above, it melts

freezing water takes on unique shape; further apart than liquid •  density of ice is less than liq water • explains potholes, weathering of rocks, water pipes, antifreeze, ice on lakes, etc., etc.

14.3 energy requirements for the changes of state • remember phase changes are physical, not chemical, changes

heating curves can help us keep an inventory on all the energy used to change phases…

let’s pretend! • a piece of ice at-10˚C can be warmed! (A) • notice as it reaches 0˚C, the line goes flat even though energy is still being pumped in… (B) • hmmmm… where is the E going?

water T stays cnst at O°C b/c E is being used to overcome H bonding! • = phase change! • the E required to change 1 mol ice to 1 mol liquid = molar heat of fusion • when all water is melted T begins to rises again... (C)

@ 100 ° C, boiling occurs, T remains cnst (D) • molar heat of vaporization = E to turn 1 mol water liquid to 1 mol vapor • = phase change! • much greater value b/c takes a lot of E to move waters from each other b/c of H-bonding • when all liquid gone, water vapor gets heated…(E)

heating curve = all physical changes! • cooling curve is exact reverse • sweating for cooling & harmful steam burns explained by all these “heats of…”

water absorbs a lot of heat for a given raise in temp - a lot! • heat capacity = thermal energy required to raise temp of one kg of a substance one kelvin • water’s h/c is so great b/c it takes a lot of E to overcome H-bonding •  water takes a longtime to give off heatwhen cooling, a longtime to absorb heatwhen warming • ( nicer weathernear ocean)

14.4 evaporation and vapor pressure • open a container of liquid and the liquid can evaporate • it overcomes the i/m forces and escapes to the atm • this process is called vaporization (aka evaporation)

evaporation can’t “just happen,” molecules need a minimal amt of E • see how there is a whole range of KE above

in a cooler liquid sample a few have enough E to leave • at higher T more have the E to leave

if those critters with enough E happen to be at the surface they go to vapor! • as the higher E molecules leave, the avg E of the leftovers is less, so it is cooler • meaning, evaporation causes cooling (e.g. sweating, fingernail polish on skin, etc.)

vapor pressure • OK, now if we close that sample (aka closed system) it still evaporates, but! • as time goes on some just eventually return - condensation

finally there will come a time when# leaving =# coming back • to our eyes it’s over (macroscopic) • really it’s very dynamic (microscopic) • but all these new guys cause a P

this pressure is called vapor pressure • the greater the i/m force the less likely it is to be a vapor and the smaller the vp

a barometer can be used to measure it

high vp’s are called volatile and have wimpy im forces • lower vp’s mean strong attraction, strong im forces • as to im forces:H-bond > dipole-dipole > london • so as to vp:london > dipole-dipole > H-bond

imf is diff here • water has hiimf • ethanol lower • diethyl ether lowest

example • which has largest VP? H2O or CH3OH • H2O has two OH bonds, methanol just one • so H2O sticks better and loses! • CH3OH has the higher VP

example • which has largest VP? CH3OH or CH3CH2CH2CH2OH • both have OH, but! the big guy has all that london force going for it with all those extra atoms • so CH3CH2CH2CH2OH sticks better and loses! • CH3OH has the higher VP

14.5 boiling point & vapor pressure • what is going on when water boils? • why do the bubbles form? where are they coming from? • remember the atmosphere is pushing down hard on the water…

water has some gas dissolved in it • imagine water molecules “evaporating” into the tiny air bubble which is acting as a “birth place” for the bubble • if it is hot enough (i.e. enough KE) the water molecules can keep pressing out into a bigger bubble

so boiling pointdepends on Tand P • BP is T at which vp = external P on liq • increased P will force them to stay together longer •  harder to boil •  BP goes up (and vise versa) • (like a radiator or pressure cooker!)

so how does elevation affect BP? • the higher you go the less P • the less P, the less E water needs to make a bubble •  water boils at lower T at high altitudes…

for freezing point P has little effect here, mostly a T thing • (reverse of freezing = melting; fp is same as mp) • wussy imf don’t hold the particles together well ( lo MP) • stronger imf hold the particles together longer ( hi MP) • (reverse of freezing =fusion)

14.6 the solid state:types of solids • solids are everywhere; some are cmpds, some elements • some mixtures, some pure • when a solid has a regular, patterned arrangement = crystalline solid • what makes up the crystals and how they are held together determines properties

put together with ions and held by ionic bonds = ? • ionic solid! • when dissolved in water, they often break into ions

molecular solids are molecules held together by i/m forces • when broken up they split into their molecules (not atoms!)

chapter 14 liquid and solids