0. Matter can exist in different states A. The state depends on the (and ). Of interest to living things.

Transcription

1 Kinetic Molecular Theory, Weak Interactions, States of Matter (Why do liquids & solids exist?) ead chapters 5.6 (Intermolecular forces), 7.1 (first four paragraphs), 7.2 thru 7.5. Lab Manual Appendix: I.1, I.2, and 1.3 Themes for the chapter: 1. What is the Kinetic Molecular Theory? 2. ow do #1 & weak interactions help us with (s), (l), & (g)? emember that KMT being a theory means it is considered to be at the highest level of development of an idea. 1. A theory is supported by many, many experiments (facts?) 2. A theory provides an intellectual framework that often leads us to new intellectual territory. 3. A theory may change through time. (Little in science is completely static, very much like life itself.) owever, it is not common for theories to be modified substantially. 0. Matter can exist in different states A. The state depends on the (and ). Melt ing vaporization Bose-Einstein ondensate Solid Liquid Gas Plasma Freezing ondensation What direction are we moving along diagram, if heat is being added? (l or r) What is sublimation? B. The state (of matter) also depends on the interactions that hold particles together. For molecules, the interactions that hold them together are called weak interactions or First we will introduce the Kinetic Molecular Theory (KMT), and then weak interactions. Then we will see how KMT & weak interactions relate to states of matter (s, l, & g), boiling point and melting point. I. The Kinetic Molecular Theory A. This theory applies very directly to gases f interest to living things. Gas: lots of, lots of space between molecules/atoms gas phase 1

2 B. Tenets (or assertions?) of the Kinetic Molecular Theory: 1. In the gas phase molecules move randomly in straight lines, at a. 2. Average energy of the molecules is related to the of the gas. 3. The molecules exchange energy when they hit each other, but the total energy is conserved. (Elastic collisions, momentum transfer.) 4. The of a gas molecule is negligible, and the space between the molecules is very large. 5. The gas molecules don t stick to each other. (There are no attractive forces.) ne characteristic of a gas is that it exerts on the walls of the container. With the kmt, we can understand pressure at the molecular level as being a result of gas molecules colliding with the container walls. So amount of pressure is related to: the number of collisions per unit time, and how fast the molecules are moving. Take a quick look at Boltzmann. First note: 1. movement of atoms 2. collisions 3. speed of atoms Variables to examine: 1. how many atoms? 2. temperature of system? 3. volume of system?. Temperature, T, is a measure of the. 1. Temperature scales: Fahrenheit, o F elsius, o Kelvin, K 2(s) melts 32 F K 2(l) boils 212 F K i) At 0 K, the motion of the molecules. This is called absolute zero ii) Must use Kelvin scale for all equations/formulas that have their origins in Kinetic Molecular Theory. There are many! 2

3 2. emember the temperature conversion formulas: = 5 9 ( F 32) K = Problem to do on your own: onvert 98.6 F into and K. D. Pressure, P 1. P is a function of the collisions between the gas and the container walls. a) ow many collisions per unit time? More collisions = (more or less?) P b) ow fast is the molecule moving? igher speed = (more or less?) P 2. Units of pressure a) atmospheres (atm) harleston, S is 1atm. b) (in medicine & science) 1 atm = mm g, No need to memorize! 1 mm g = 1 torr E. The Ideal Gas Law 1. The Ideal Gas Law comes directly from kinetic molecular theory. The Ideal Gas Law is: PV = nt, where P = pressure V= volume n= # of moles (1 mole = x molecules [atoms?]) = the Universal Gas onstant = (L atm) (mol K) This helps with units!!! T= temperature For P units must be For V units must be For n units must be For T units must be 2. un through logic associated with PV=nT Take a closer look at Boltzmann. a) Why does P increase as n increases? b) Why does P increase as T increases? c) Why does P decrease as V increases? 3

4 3. The Ideal Gas Law works well for gases not near the condensation point. Given any 3 variables, you can always calculate the other. Problem: What volume will 9.00 mol of e occupy at 20.5 o & a pressure of atm? F. ther Gas Laws that are special cases of Ideal Gas Law (where one or more variables are kept constant.) 1. The following three laws are 2 condition problems. All assume n is constant. Begin at condition1. ow does changing one variable alter another? 2. Boyle s Law: P x V is constant a) Assumes T and n are kept constant. b) A useful way to express this is: c) P and V are inversely proportional d) If V going from state1 to state2, would P or? e) P1V1 = P2V2 can be derived from the Ideal Gas Law P1V1 = n1t1 hange P & V, keep n & T constant: P2V2 = n1t1 Because n1t1 = n1t1 then P1V1 = P2V2 Problem: If an 2 tank used in a health care situation has a volume of 4.50 L and a pressure of 33.5 atm, what is the volume of the same amount of gas if the pressure is 0.98 atm? Assume the temperature is constant. 3. harles Law: V is a constant. T a) This assumes that P and n are constant. b) A useful way to express this law is V 1 = V 2 T 1 T 2 c) V and T are directly proportional. d) As T, would V or? (What must always be done with temperatures in a gas laws problem?) 4

5 4. Gay-Lussac s Law: P T is constant. a) Assumes n and V must be constant! P1 b) As T increases, P increases. (Industrial applications!) 5. The ombined Gas Law (a special case of the ideal Gas Law): P x V T is constant G. Dalton s Law 1. The total pressure is just the sum of the pressures exerted by each of the components in a mixture of gases. 2. Algebraically: PT = P1 + P2 + P This means that the Ideal Gas Law applies individually to each of the components in a mixture of gases. Air PT = Poxygen + Pnitrogen + P What other P? would be important for a patient receiving respiratory gases? II. Weak Interactions (Important!!!) A. Introduction 1. Liquids and solids form because item #5 of the kinetic molecular theory does not completely apply to real (as opposed to Ideal) gases. There are 2. All materials (even e) can exist in (l) [or (s)] state if temperature is low enough. 3. Discuss three weak interactions or inter-molecular forces, IMF a.. London forces (or dispersion forces) b. dipole-dipole interactions c. hydrogen bonding These forces are additive. All molecules have London forces. nly polar molecules have dipole-dipole forces. ydrogen bonding requires particular structures in a molecule. B. Temporary dipole-induced dipole (a.k.a. London forces or dispersion forces) 1. All atoms exhibit London Forces. e, a single atom, can be liquefied. (b.p. ~ 4 K) ow do e atoms stick together? Imagine 2 cold helium atoms that happen to be next to each other: (see next page) 5

6 Even charge 2p + e - 2p + e - Every so often distribution gets uneven. ere, the left atom forms a temporary dipole δ - 2p+ δ + 2p+ e - The left hand dipole causes a dipole to form in the right atom. This is the induced dipole δ - 2p+ δ + δ - 2p+ δ + Electrons are very mobile, so the temporary dipole goes away in a relatively short time. So does the attraction. e - 2p+ 2p+ e - 2. London Forces increase with increasing mass/size. a) molecule mass bp F-F ~ o F l-l ~71-34 o F Br-Br ~ o F. b) Atoms or molecules with more electrons can be more easily distorted to form temporary dipoles. Discussed more below 6

7 . Dipole-dipole interactions 1. emember our previous interest in polarity? Polar? (Y or N) Polar? (Y or N) butane acetone an you locate δ + and δ - on either structure? 2. This is a result of standard electrostatic attraction (Positive charge attracts negative charge.) Draw a picture below showing how 2 acetone molecules would stick together: D. ydrogen Bonds (Important!!!) 1. It is not a standard covalent bond. (Not every bond with is a ydrogen Bond ). 2. To form a ydrogen Bond, you must have: a) a highly e deficient atom (bonded to ) b) a non-bonding e pair on a N,, or F atom. c) appropriate orientation & distance of (a) & (b). ontrast w/ dipole-dipole interactions. ydrogen bond components donor site acceptor site X Y N,, or F N,, or F Notes: i) Use a dashed or dotted, not solid, line to show this bond type Why shouldn t you use a solid line? ii) Bond strength is greater with larger electronegativity differences (Which has higher electronegativity, N or?) iii) Attraction between molecules is stronger if more ydrogen Bonds are involved. Two arm vs. one arm pullup? iv) In a biologically interesting ydrogen Bond, both X and Y have electronegativity values greater than or equal to 3.0. Which elements have electronegativity 3.0? 7

8 3. Identify the donor (d) and acceptor (a) sites: N 4. Show one of the structures above forming a ydrogen Bond: 5. Which would be stronger, and why: a) N b) N N N N c) E. omparison Non polar atoms/molecules Polar molecules Polar molecules w/ hydrogen bonding capabilities Example Example: 4 Example: 2 Example: 2 IMF London dispersion London dispersion dipole-dipole London dispersion dipole-dipole hydrogen bonding 8

9 F. ow much E to break an individual: Energy (kj mol) Example 1) Ionic bond(s) ,000 Na + l - (in a crystal) (lattice energy) 2) ovalent bonds: or 3) Ionic bonds(aq): N3 (in 2) 4) ydrogen Bond: N2 5) Dipole-dipole: 9 = δ δ + = 6) London Dispersion 0.3 (Small but mounts up with many atoms.) III. States of Matter A. Gas (g) (See Boltzmann) 1. Microscopically: lots of movement, lots of space between molecules/atoms 2. Macroscopically: fills the container, whatever its size (7 molecules) 3. Is density high or low in a gas? (ompressibility!) B. Liquid (l) 1. Microscopically: lots of movement, (Molecules sliding around each other.) very little space between molecules/atoms 2. Macroscopically: adapts to shape of container, and has a fixed density (at a given T) (24 molecules) gas phase liquid phase, imagine little tiny arrows! 3. ompare density with gas phase above.. Solid (s) The one shown is crystalline. 1. Microscopically: only movement in place (oscillate, vibrate), very little space between molecules solid phase, movement in place only! 2. Macroscopically: has fixed shape, fixed density is usually greater than liquid phase & relatively T independent. (24 molecules) 3. Apparent density relative to (g) and (l)? A summary: D. If you start from the gas phase, why do liquids and solids form? 1. Because of Intermolecular attractive forces (a.k.a., weak interactions) 9

10 2. What does it mean when a type of molecule in the gas phase condenses at a relatively high temperature? It has E. Increasing mass/size, increases London Forces and increases boiling point. a) Molecules have more electrons. b) They can be more easily distorted to form temporary dipoles. c) ompare ethane, butane and hexane b.p. ethane o butane o hexane o F. Shape of molecules: Effects on strength of London forces and b.p. 1. Molecules must be close together for strong London forces 2. Long, thin molecules have stronger London forces than shorter, rounder molecules of the same size. 3. Which molecule has the strongest IMF and highest bp? a. b. c. Determining weak interactions and boiling points 1. What IMF does each molecule have? Which molecule has the stronger IMF forces overall and which bp belongs to which molecule (-0.4 o, -42 o )? propane butane 2. What IMF does each molecule have? Which molecule has the stronger IMF forces overall and which has the lower boiling point? butane 2-methylpropane 10

11 When focusing on IMF in addition to London forces, it is best to compare molecules of similar size. 3. What IMF does each molecule have? Which molecule has the stronger IMF forces overall and which has the lower boiling point? ethanol propane 4. What IMF does each molecule have? Which molecule has the stronger IMF forces overall and which has the lower boiling point? propanal 1- propanol 5. What IMF does each molecule have? Which molecule has the stronger IMF forces overall and which has the lower boiling point? acetic acid propanol F. What occurs (at molecular level) when 2 boils? 1. Defining vapor pressure initial Molecules start to equilibrate with (g) phase. evaporation occurs at equilibrium evap = condens a) Start w/ a liquid in a closed container. (left panel, above) 11

12 vapor pressure (mm g) b) The liquid phase molecules are moving, and some are moving fast enough to escape from their neighbors and go into the gas phase (center panel) i) This process is called evaporation ii) Which molecules are most likely to escape liquid (Boltzmann distribution?) c) As molecules build up in the vapor phase, they start colliding with the surface of the liquid i) This process is called condensation ii) When the evaporationrate = condensationrate, we have reached equilibrium. d) Vapor pressure is the pressure exerted by the gas phase due to evaporation. Vapor Pressure of 2 e) This varies with temp. As T, the vapor P mm g 2. ow does vapor pressure vary as a function of T? temperature ( o ) 3. What happens when the vapor pressure = the external atmospheric pressure? 4. So why it takes longer to cook an egg by boiling in Denver, than in Greer, S? In harleston, atm press. = 760 mm g. In Denver, atm press. = 560 mm g n Mt Everest (top), atm press = 200 mm g. Water boils at Water boils at Water boils at 12