ch11 imfs

Question 1

characteristics of gas

Answer 1

assumes volume and shape of container, is compressible, flows readily, diffusion within a gas occurs rapidly

Question 2

characteristics of liquid

Answer 2

assumes shape of the portion of container it occupies, does not expand to fill container, incompressible, flows readily, diffusion within occurs slowly

Question 3

characteristics of solid

Answer 3

retains own shape and volume, incompressible, does not flow, diffusion within occurs extremely slowly

Question 4

condensed phases

Answer 4

Solids and liquids are condensed phases because particles are close together. Solids that possess highly ordered structures are crystalline.

Question 5

how to change state

Answer 5

The state of a substance depends largely on the balance between the kinetic energies of the particles and the interparticle energies of attraction. KE keeps particles apart, interparticle attractions draw particles together.Gases have weaker interparticle attractions than liquids
We can change a substance from one state to another by heating or cooling which changes the average kinetic energy of the particles. As temp decreases, KE decreases.
Increasing pressure of a gas forces molecules closer together, which strengthens intermolecular forces of attraction

Question 6

types of van der waals and 1 other4

Answer 6

dipole-dipole, london dispersion, and hydrogen bonding. The other important force is ion-dipole force. All of these forces are electrostatic, so involve positive-negative attractions

Question 7

ion dipole force

Answer 7

between ion and partial charge on the end of a polar molecule (dipole). Positive ions attracted to negative end of dipole, and vice versa. Magnitude of attraction increases as charge of the ion or magnitude of dipole moment increases

Question 8

dipole dipole

Answer 8

positive end of one molecule is near negative end of another. Found in polar molecules. Weaker than ion-dipole. For molecules of approximately equal mass and size, the strengths of intermolecular attraction increase with increasing polarity=increased boiling point.

Question 9

ldf

Answer 9

motion of electrons in molecule can create an instantaneous dipole moment. In nonpolar molecules. (if you could freeze moment, at that instant, the electrons could be on one side). Temporary dipoles on one atom can induce dipoles on adjacent atoms, causing LDF. Strength of LDF depends on ease with which charge distribution can be distorted=polarizability. Higher molecular weight=More polarizable=stronger LDF. Longer/cylinder molecules=higher boiling point because more contact is possible than sphere shaped

Question 10

comparing strengths of forces in molecules3

Answer 10

• ldf in all molecules
• When molecules have sameish weight and shape, LDF is equal. Attractions are due to dipole-dipole interactions
• When molecules differ in weight, more massive molecule has strongest attraction due to LDF

Question 11

hbonding

Answer 11

unique dipole-dipole attraction between an H atom in a polar bond ( HF, HO, HN) and an unshared electron pair on a nearby small electronegative ion/atom (F, O, or N atom in another molecule). Happens because F O N are so electronegative, and H only has 1 electron. Stronger than dipole-dipole bonds

Question 12

h bonding in water

Answer 12

ice assumes structured shape with lots of open space, so liquid water is more dense than ice. (ice floats)

Question 13

viscosity

Answer 13

resistance of a liquid to flow. Greater viscosity=flows slower. Measured by time it takes to travel through thin tube under gravitational force. Increases with stronger molecular forces, so Viscosity increases as molecular weight increases and decreases with increasing temperature. SI units for viscosity are kg/m-s.

Question 14

surface tension

Answer 14

energy required to increase the surface area of a liquid by a unit amount. Water has a high surface tension because it has strong H bonds.
molecules at the surface are attracted only by molecules below them while interior molecules are attracted equally in all directions, so there is a net downward attraction into the interior of the liquid. Because spheres have smallest surface area for their volume, water droplets form sphere shape.

Question 15

cohesive force

Answer 15

intermolecular forces that bind similar molecules to one another, such as H bonding in water.

Question 16

adhesive force

Answer 16

intermolecular forces that bind a substance to a surface. Water in a glass tube sticks to the glass on the sides because of adhesive forces, which is why water has a meniscus (curved upper surface) that is U shaped.

Question 17

capillary action

Answer 17

rise of liquids up very narrow tubes (capillaries). Adhesive forces increase the surface area of the liquid, but surface tension reduces the area, pulling the liquid up the tube, until the cohesive and adhesive forces are balanced by the force of gravity on the liquid.

Question 18

sublimation

Answer 18

solid→ vapor state

molecules of a solid transformed directly into gaseous state. Hsub=Hfus+Hvap

Question 19

fusion

Answer 19

melting process where the units that made up a solid freed to move with respect to one another. Heat/enthalpy of fusion: Hfus, adding E to system to melt a solid

Question 20

vapor pressure3

Answer 20

• pressure of gas-phase molecules over a liquid that increases with increasing temp. VP increases with increasing T until it equals external pressure, the point where the liquid boils. Energy for this is heat/enthalpy of vaporization Hvap
• constant value of pressure attained when a liquid is evaporated in a closed container and the pressure of the area above the liquid increases.
• The weaker IMF and higher temp=larger number of molecules able to escape and higher vapor pressur

Question 21

hvap

Answer 21

heat of vaporization (boiling point

greater than Hfus vs transition from liquid to vapor, molecules must sever all intermolecular attractions

Question 22

deposition

Answer 22

vapor–> solid
Heat of deposition is exothermic to the same degree as heat of sublimation, heat of freezing and condensation equal and opposite to heat of fusion and vaporization.

Question 23

heating curve

Answer 23

graph of temperature of the system versus the amount of heat added. In a heating curve from a solid to a gas, there are 3 segments of verticalish lines (heating up phase to higher temp), and 2 segments of horizontal lines (phase change, temp remains constant)

Question 24

specific heat in heating curve

Answer 24

The greater specific heat=more heat needed to accomplish a certain temp increase. Since specific heat of water is greater than ice, slope of heating liquid segment is less than slope of heating ice.

Question 25

supercooling

Answer 25

rapidly cooling a substance below its freezing point without forming a solid.

Question 26

critical temp

Answer 26

highest temp at which a distinct liquid phase can form (highest temp at which a liquid can exist). Critical pressure: pressure required to bring about liquefaction at critical temp.
More polar=greater MW=Greater IMF=greater critical temperature of a substance

Question 27

calculating energy from heating curve

Answer 27

AB: heating solid Q=mcT
BC:melting Hfus*mols
CD:heating liquid Q=mcT
DE: evaporation Hvap*moles
EF: heating gas Q=mcT

Question 28

dynamic equilibrium

Answer 28

condition in which two opposing processes are occurring simultaneously at equal rates. A liquid and its vapor are in dynamic equilibrium when evaporation and condensation occur at equal rates. The vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor state are in dynamic equilibrium

Question 29

critical point

Answer 29

beyond critical point liquid and gas phases become indistinguishable from each other.

Question 30

triple point

Answer 30

all three phases in equilibrium

Question 31

melting point h2o

Answer 31

when liquid form is more dense than solid form, melting line will slant to left

Question 32

crystalline solid

Answer 32

atoms/ions/molecules are ordered in well-defined arrangements. Have flats surfaces/faces that make angles with one another and have highly regular shapes (quartz, diamond)

Question 33

amorphous solid

Answer 33

no orderly structure, no faces or defined shapes. Mixtures of molecules that do not stack together well or large, complicated molecules. (rubber, glass)
Amorphous solids do not melt at a specific temperature, because intermolecular forces vary throughout a sample

Question 34

unit cell3

Answer 34

• repeating unit of a solid.
• Unit cells are parallelepipeds (6 side, faces are parallelograms). Unit cells are described by their lengths of edges and angle between edges (cubic unit cell=sides are equal in length and angles are 90º)
• The total cation-to-anion ratio of a unit cell must be the same as that for the entire crystal. (NaCl must have equal # of Na and Cl)

Question 35

crystal lattice

Answer 35

3D array of points to represent a crystalline solid. Each point=lattice point. Form a crystal structure by arranging unit cells repeatedly on crystal lattice.

Question 36

types of cubic unit cells3

Answer 36

• primitive cubic(simple cubic)8 corners (C#=6)
• Body centered cubic: 1center, 8corners (C#=8)
• Face-centered cubic: 8corners, 6faces (C#=12)

Question 37

hexagonal close packing

Answer 37

spheres of third layer placed in line with those of the first layer. Fourth layer repeats the second layer (ABAB)
C#=12

Question 38

cubic close packing

Answer 38

spheres of third layer do not sit above first layer. Fourth layer repeats second layer (ABCA)
c#=12

Question 39

coordination number

Answer 39

umber particles immediately surrounding a particle in the crystal structure. As coord # increases, % of total volume occupied by spheres increases (more tightly packed)

Question 40

molecular solids

Answer 40

atoms/molecules held together by weak intermolecular forces. Soft and low melting points, poor conductivity. (Ar, H2O, CO2)

Question 41

covalent-network solids

Answer 41

atoms held together in large networks/chain by covalent bonds. Harder and higher melting point than molecular solid, because covalent bonds are strong. (Diamond, SiO2) Often poor conductors

Question 42

ionic solids

Answer 42

ions held together by ionic bonds. Higher charges on ions=stronger bond and higher melting point. Hard, brittle, poor conduction

Question 43

structures of ionic solids4

Answer 43

1) NaCl/KCl/CaO,etc (ions have +/-1 charge), coordination #6, face centered about Cl, with Na in open edges
2) CsCl-primitive cubic arrangement, coordination#8. Ions have +/-1 charge but Cs is bigger than Na. body centered unit arrangement
3) ZnS-face centered cubic arrangement (ions have +/-2 charge)
4) CaF2-face centered cubic but there are twice as many Ca than F. more e than 3. 4 molecular units in a unit cell

Question 44

metallic solid

Answer 44

metal atoms, hexagonal lose-packed/cubic/body-centered cubic structures (8 or 12 adjacent atoms). Valence electrons are delocalized throughout the entire solid. Have wide range of physical properties. Strength of bonding increases as number of electrons available (valence electrons) for bonding increases. Good conductors. Also malleable and ductile

Question 45

clausius clapeyron equation

Answer 45

lnP= -Hvap/RT+C
as T increases, VP increases
Hvap= -slope*R