Unit #3

Question 1

true or false: weak interparticle forces are distrupted in the phase changes s -> l -> g, but the molecule remains intact

Answer 1

t, no intramolecular bonds are broken

Question 2

what are the physical properties of gases (5) and are they dependent on their chemical composition

Answer 2

1) highly compressible (boyle’s law)
2) thermally expandable (charles’ law)
3) low viscosity
4) low density
5) infinitely miscible (pattern to behaviour (regardless of composition)

• the physical properties of gases are largely independent of their chemical compositon

Question 3

what is the equation for the ideal gas law

Answer 3

PV=nRT

Question 4

what is the relationship between different types of pressure

Answer 4

1 atm = 760 torr = 760 mm Hg = 101.325 kPa

Question 5

what is standard temperature and pressure

Answer 5

0˚C (273.15 K) and 1 atm

Question 6

what are the applications of the ideal gas equation (3)

Answer 6

1) molar mass determination (since moles and mass are related by n=m/M, the idea gas equation can be rearragned to get PV = (m/M)RT or M=mRT/PV)
2) gas density (because d=m/V you can rearrange to get d=MP/RT)
3) gas stoichiometry

Question 7

how do you do combustion analysis

Answer 7

• start by using ideal gas law but changing n to m/M and isolate M then solve
• find the number of moles for C and H (ex. C = 3.613g CO2 x 1 mol / 44.01 g CO2 x 1 mol / 1 mol CO2 = 0.08209 mol C)
• determine if the sample contains oxygen by transforming the moles of C and H into grams and seeing is the sum of each mass equals the same amount of sample present
• hydrocarbon is CxHy so divide both amounts by the smallest amount of moles then find the ratio
• solve for molar mass and change proportions accordingly

Question 8

true or false: gases are not infinitely miscible, therefore, they cannot mix freely in any proportions

Answer 8

false, they are infinitely miscible, meaning that they can mix freely in any proportions, ex. air

Question 9

what are partial pressures

Answer 9

the pressure that a gas would exert if it were alone in the container
Ptot = (ntRT)/V (nt = n1+n2+n3+…)
P1 = (n1RT)/V
P2 = (n2RT)/V (and so on)
Ptot= P1+P2+P3+…

Question 10

what is dalton’s law of partial pressures

Answer 10

in a mixture of non-reacting gases, the total pressure is the sum of the partial pressures of individual gases.
- another way of expressing this law is in terms of the mole fraction of a component i which is: xi = ni/ntot
- from this, we conclude: Pi=xi*Ptot or xi = Pi/Ptot

Question 11

what is the kinetic molecular theory (5 pts.)

Answer 11

to explain the behaviour of gases, we develop a model of idea gas:
1) particles are in random motion (collision of particles w/ walls and other molecules cause pressure)
2) negligible particle volume (volume of individual particles can be neglected)
3) particles collide with each other, and container walls (collisions are elastics (X energy loss))
4) particles move independently and experience no interparticle forces
5) constant total energy (but energy is transferred in collisions, w/ KE conserved)

Average kinetic energy of the collection of gas particles is directly proportional to the temperature of the gas depends only on its temperature

Question 12

how can we calculate the pressure exerted by the gas on the wall

Answer 12

by summing the impacts made by many particles
P=mu1^2/V (u1 = velocity of particle 1)
or for a larger number of N:
P=1/3(NMu^2)(w/ _ on top)/V\ (u^2 w/ _ on top means mean-squared speed = 1/N(u1^2+u2^2+…)

Question 13

what are distributions of particle speeds

Answer 13

a way to depict a sample of gas that has a variety of speeds due to collisions.
goes off of molecules over speed
- note: urms>uavg>um

Question 14

which has the faster urms (avg speed):
1) H2 (at 0˚C) or O2 (at 0˚C)
2) O2 (at 273 K) or O2 (at 1000K)

Answer 14

1) H2 as lighter gases have greater average speeds than heavier gases
2) at 1000K as gases have greater average speeds at higher temperatures

the distribution is wider at higher temps

Question 15

true or false: higher T = greater KE = greater motion

Answer 15

true, temperature measures the degree of random motion in a gas

Question 16

what does a temperature mean for one particle

Answer 16

temperature is a property that only applies on macroscopic level
- it is NOT a property of fundamental particles
- if you have 1 particle it means you have one speed therefore temperature is not defined using the relation above

Question 17

what does 0 temperature mean

Answer 17

0 temperature means 0 motion, all molecular motion ceases at absolute 0

Question 18

what is the difference between real gases and ideal gases

Answer 18

in reality gases never actually act like ideal gases.
ideal gases only exist in low P and high temps

Question 19

what occurs in real gases that violate the assumptions made in kinetic molecular theory

Answer 19

1) particles in a real gas experience weak interparticle attractions (@ moderately high P, interparticle attractions are important)
2) particles in a real gas occupy a finite volume (@ very high P, particle volumes dominates)

Question 20

when you consider the difference between ideal and real gases, how do you alter IGL to measure properly

Answer 20

you use van der Waals equation
(P+(an^2)/V^2)(V-nb)=nRT
- P+(an^2)/V^2 - adjusting measured pressure up to account for attractive forces, larger a ⇌ stronger attractive forces
- V-nb - adjusting volume down to account for volume of gas particle, larger b ⇌ larger particle volumer (size)
*depends on conditions if IGL is used or van der Walls eq’n

Question 21

true or false: during phase changes, the intramolecular forces are intact, but intermolecular forces are disrupted or formed

Answer 21

true, the chemical behaviour of the 3 states is identical because all of them consist of the same molecule held together by the same bonding forces. the physical behaviour is different because the strength of IMFs differs from state to state

Question 22

what happens to KE when temperature increases and decreases

Answer 22

• as temperature increases, KE avg increases, thus particles move faster and overcome attractive forces more easily
• as temperature decreases, KE avg decreases, thus particles move more slowly and attractions can pull them together more easily (usually succumbing to attractive forces)

Question 23

what are the energy transitions when molecules require energy

Answer 23

melting, vaporization and sublimation, these all have positive enthalpy (endothermic) as they are absorbing heat

Question 24

what are the energy transitions when molecules release energy

Answer 24

freezing, condensation, and deposition, these all have negative enthalpy (exothermic) as they are releasing heat

Question 25

why is ∆H˚fus &laquo_space;∆H˚vap, in general, for all substances

Answer 25

it takes less energy to reduce the IMF enough for the molecules to move out of their fixed positions (melt a solid) than to separate them entirely (vaporize a liquid)

Question 26

what state is a phase equilibrium

Answer 26

a dynamic state (rate of vaporization = rate of condensation @ constant pressure)

Question 27

what effect does temperature have on vapour pressure

Answer 27

a major one, because it changes the fraction of molecules moving fast enough to escape the liquid (and thus the fraction of molecules moving slowly enough to be “recaptured”)
- at the higher temperature, more molecules have enough energy to leave the surface. thus, in general, the higher the temperature is, the higher the vapour pressure

Question 28

what is the effect of intermolecular forces on vapour pressure

Answer 28

the vapour pressure depends only on the type of liquid (not the amount) and the temperature. substances w/ weak interparticle forces in the liquid have high vapour pressures.
- at a given temperature, all substances have the same average Ek. thus, molecules w/ weaker intermolecular forces are held less tightly at the surface of the liquid and therefore will vapourize more easily

Question 29

which has the higher vapour pressure, water or diethyl ether (CH3CH2)2O

Answer 29

the intermolecular forces in diethyl ether are weaker compared to water (which contains strong H-bonds) thus it will have the higher vapour pressure

Question 30

what is aboiling point

Answer 30

temperature at which the vapour pressure (inside bubbles) in the liquid equals the external pressure, which is usually the atmospheric pressure.
vapour pressure = external pressure
- substances w/ high vapour pressures have low boiling points

Question 31

what is a volatile substance

Answer 31

substances that have low boiling points and relatively weak intermolecular forces

Question 32

what is the solid-liquid equilibria

Answer 32

the particles in a crystal are continually vibrating about their fixed positions. as the temperature rises, the particles vibrate more rapidly until some have enough Ek to break free of their position -> melting begins
- as more molecules enter the liquid phase, some collide with the solid and become fixed in position again. because the phase remains in contact, a dynamic equilibrium is established when the melting rate equals the freezing rate. the temperature at which this occurs is called the melting point
- because liquids and solids are nearly incompressible, pressure has little effect on the rates of melting and freezing. a plot of pressure vs. temperature for a solid-liquid phase change is nearly a vertical line

Question 33

what is the solid-gas equilibria

Answer 33

a substance sublimes rather than melts because the intermolecular attraction are not great enough to keep the molecules near each other when they leave the solid state
- some solids do have high enough vapour pressures to sublime at ordinary conditions, including dry ice, iodine, etc… all of which consists of nonpolar molecules/substances w/ weak IMFs

Question 34

what are phase diagrams

Answer 34

depict which phase is most stable at different conditons of temperature and pressure. (solid top left, liquid top middle, gas bottom right and left bottom and top right corner)
- solids - stable at low temperatures and high pressure
- gas - stable at high temperatures and low pressure

Question 35

what do the lines indicate on phase diagrams and what do the slopes represent

Answer 35

they indicate which two phases coexist in equilibrium.
- positive slopes indicate that the solid is more dense than the liquid, an increase in pressure converts the liquid to the solid
- negative slopes indicate that the solid is less dense than the liquid, an increase in pressure converts the solid to the liquid

Question 36

what is the critical point on phase diagrams

Answer 36

the point at which the l-g line terminates. above the P and T of the critical point, the vapour cannot be liquified, this state is called a supercritical fluid.
- when a liquid is heated in a closed container, it’s density decreases. at the time, more of the liquid vapourize, so the density of the vapour increases. as the critical point, the 2 densities become equal and the phase boundary disappears. the average KE is so high at this point that the vapour cannot be condensed anymore

Question 37

what is the triple point on phase diagrams

Answer 37

the point at which 3 phases can coexist in equilibrium

Question 38

true or false: phase diagrams only indicates the thermodynamically stable phase, it is not possible for other phases to exist

Answer 38

false: phase diagrams only indicates the most thermodynamically stable phase. it is possible for other phases to exist (under metastable conditions)

Question 39

what is an example of supercritical substances and processes

Answer 39

decaffeinating coffee using supercritcal CO2, supercritical fluid extraction using CO2 is now being widely used as a more effective and environmentally friendly decaffeination method.
- it penetrates deep into the coffee beans, allowing it to effectively dissolve certain substances.
- because CO2 is a gas under STP, it’s removal from the extracted coffee beans is easily accomplished along w/ the recovery of the caffeine.
also freeze-dried foods

Question 40

what properties does intermolecular forces influence

Answer 40

they influence only the physical (NOT CHEMICAL) properties of a substances. their strengths are reflected in:
- mp, bp
- heats of fusion, vaporization
- surface tension, viscosity -> liquids
- solubility -> solids

Question 41

what are substances that are not molecular (do not contain molecules)

Answer 41

• simple ionic solids
• extended network solids
• the entities in a substance can thus be atoms, ions, or molecules

Question 42

rank the type of bonding chemical forces in order of strength

Answer 42

ionic (cation-anion attraction, s: 400-4000) > covalent (nuclei-shared e- pair, s: 150-1100) > metallic (cations - delocalized electrons, s: 75-1000)r

Question 43

rank the type of non-bonding chemical forces in order of strength (van der Waals)

Answer 43

ion-dipole (ion charge - dipole charge, s: 40-600) > dipole-dipole (dipole charges, s: 5-25) > H-bond (polar bond to H - dipole charge, s:10-40) > ion-induced dipole (ion charge - polarizable e- cloud, s: 3-15) > dipole -induced dipole, s: 2-10) > london dispersion/induced dipole-induced dipole (polarizable e- clouds, s: 0.05-40)

Question 44

what is ionic bonding and how do you determine strength

Answer 44

ion is attracted by other ions of opposite charge, strength is determined by F=-kq1q2/r^2
- stonger the force, the more energy needed to overcome forces therefore higher melting points in ionic solids

Question 45

what are ion-dipole attractions

Answer 45

attraction between the charge of the ion and an opposite dipole (ex. NaCl in H2O)
- ions become separated because attractions between the ions and the oppositely charged poles of the water molecule are stronger (or equal to) than the attractions between the ions themselves

Question 46

what are dipole-dipole attractions

Answer 46

molecules w/ permanent dipole moments experience dipole-dipole interaction (relatively weak forces)
- molecules will orient to maximize positive and negative attractions or minimize like-charge repulsions these interactions are not important in gases due to molecules being far apart
these forces depend on the magnitude of the dipole moment

Question 47

what is hydrogen bonding

Answer 47

a special type of dipole-dipole force between molecules that have an H atom bonded to a small, highly electronegative atom
- the H-N, H-O, and H-F bonds are very polar. when the partially positive H of one molecule is attracted to the partially negative lone pair on the N, O, or F of another molecules -> a hydrogen bond forms

Question 48

what happens when water solidifies in terms of H-bonding

Answer 48

H-bonding between the molecules line up in a way that creates empty space between the molecules, increasing volume of the solid. this is why the ice is less dense than liquid water

Question 49

what is polarizability

Answer 49

the ease with which the electron cloud can be distorted to create a charge separation. polarizability increases with greater number of electrons and larger molecular size. in some instances, large molecules with high molar mass may experience quite substantial dispersion forces

Question 50

what are the periodic trends of polarizability

Answer 50

• down a group it increases because atomic size increases, and larger electron clouds are easier to distort
• across a period it decreases because increasing Zeff makes the atoms smaller and holds the electrons more tightly
• cations are less polarizable than their parent atoms because they are smaller; anions are more polarizable because they are larger

Question 51

what are london dispersion forces

Answer 51

this interaction is caused by the instantaneous position of an electron in a molecule, which temporarily makes that point of the molecule negatively charges and the rest of the molecule positively charged
- all molecules have it, more important in molecules w/o permanent dipoles
- disturbance of the electron cloud, caused by applying and electric/magnetic field or by asymmetric movement of molecules.
- strength depends on polarizability of particles

Question 52

what is surface tension

Answer 52

refers to the resistance of a liquid to increase its surface area. it is the energy required to increase the surface area by a given amount. this property is dependent on the intermolecular forces

Question 53

what are the forces acting on a particle at the surface vs. in the interior of a liquid and how does this result in surface tension

Answer 53

• an interior molecule is attracted by others on all sides
• a surface molecule is only attracted by others below and to the sides so it experience a net attractive force downward (tend to want to move into molecule)
• the stronger the intermolecular forces between particles, the more energy it takes to increase the surface area, so the greater the surface tension

Question 54

true or false: surface tension decreases with increasing temperatures

Answer 54

true: at higher temperatures, the liquid molecules have increase kinetic energy w/ which to break attraction in the interior.

Question 55

what is capillarity and wetting and general trends of them

Answer 55

capillarity refers to the rising of a liquid in a tube while wetting is the spreading of a liquid across a surface. here, we consider the competition between cohesive forces (w/i) and adhesive forces (b/w)
- when cohesive > adhesive, liquid retains shape, greater capillarity
- when cohesive < adhesive, liquid spreads out, greater wetting

Question 56

what is viscosity and trends w/ temperature and shape

Answer 56

refers to the resistance of a liquid to flow (high viscosity implies strong IM forces)
- viscosity decreases with increasing temperature, faster moving molecules overcome intermolecular forces more easily, so the resistance to flow decreases with increasing temperature
- small, spherical molecules make little contact and pour easily whilst long molecules make more contact and become entangled and pour slowly (think spaghetti and meatballs)

Question 57

what is a solution and solubility

Answer 57

a solution is a homogeneous mixture in which the solute (minor) is dissolved in the solvent (major). solubility is generally “like dissolves like”
- attractive forces between the same molecules are replaced by similar attractive forces w/ solvent molecules

Question 58

what are the two types of structures that solids can be

Answer 58

• crystalline: has well-defined shape due to their particles (ions, atoms, molecules) occurring in an orderly arrangement
• amorphous: have poorly defined shapes because their particles lack orderly arrangement

*for our purposes crystalline solids are focused on only

Question 59

what are different types of crystalline solids (2 categories: 5 structures total) and what are their properties

Answer 59

discrete:
- non-polar (LDFS, properties: soft, low mp, poor conductors)
- polar (dipole-dipole, H-bonding, and LDFs, properties: fairly soft, low to moderate mp, poor conductors)
extended:
- ionic (ion-ion, properties: hard and brittle, high mp, good conductivity (in molten state)
- covalent network (covalent bonding, properties: very hard, very high mp, poor conductivity)
- metallic (metallic bonding, properties: soft to hard (depends on the metal), usually high mp, lustrous, ductile, malleable, excellent conductors)

Question 60

what is the structure of crystals

Answer 60

the highly symmetrical external form of crystals suggests a regular ordered arrangement of atoms. describe contents by unit cells (small repeating units - there are a very large number of atoms, so we need a compact way to describe their arrangement)

Question 61

what are lattices

Answer 61

an ordered array of points which describes the arrangement of particles that form a crystal. it consists of all points w/ identical surroundings
a lattice is imaginary, the crystal structure is what is real

Question 62

what are the 3 main types of cubic lattices

Answer 62

note: all planes are equidistant and mutually perpendicular
1) simple cubic: 8 identical particles define the corners, coordination #: 6
2) body centred cubic: 8 identical particles at the corners and 1 at the center, coordination #: 8
3) face centred cubic: 8 identical particles at the corners and 6 at the faces, coordination #: 12

Question 63

how are atoms that lie on a boundary of the unit cell classified

Answer 63

they are shared, so we have to count a “fraction of a sphere” as their contribution to the contents of the unit cell
- a corner is shared by 8 cells ∴ 1/8 contribution
- a edge is shared by 4 cells ∴ 1/4 contribution
- a face is shared by 2 cells ∴ 1/2 contribution
- a particle w/i a cell are not shared ∴ 1/1 contribution

Question 64

what are the 2 different ways that molecules can be packed in each layer

Answer 64

1) grid like - further packing ∴ less efficient
2) hexagonal - closer packing ∴ more efficient

Question 65

what are the different types of packing sequences? (4) what are trends in efficiency?

Answer 65

a) simple cubic - stack on top of each other
b) body centred cubic - 2nd layer over diamond spaces in 1rst
c) hexagonal closest packing (ababab…) HCP 6 atoms/cell
d) cubic closest packing (abcabc…) FFC 4 atoms/cell

sc > bcc> hcp = ccp

Question 66

what is the structure of metallic solids

Answer 66

it is nondirectional (sea of electrons), hence the atoms pack tightly to fill space. if atoms are rigid and non-interacting, the most efficient packing of spheres in 2D gives a closest-packing layer. In 3D, these layers can stack in diff. sequences.

Question 67

what are the three common types of structures found for metallic elements

Answer 67

hcp, stacking sequence (ss): abab…, primitive hexagonal
ccp, ss: abcabc…, fcc
bcc, ss: (not closest-packed), bcc
- combinations of two or more metals can give improved physical properties (strength, hardness, resistance to corrosion)

Question 68

what are the strucutres of ionic solids

Answer 68

it is non-directional, but the structures now contain 2 types of spheres:
- larger anions found in close-packing arrangements
- smaller cations are found in holes
structure can be thought of as a fcc array of anions w/ interpenetration fcc cation lattice.

Question 69

what are the different allotropes of carbon in respect to structure, bonding, and properties

Answer 69

diamond: fcc arrangement
- bonding: adopts a fcc lattice, C has a tetrahedral arrangement in this endless array. ELECTRONS ARE LOCALIZED
- properties: transparent, poor conductivity
graphite: hexagonal rings
- bonding: flat sheets of hexagonal carbon rings w/ strong sigma bond framework and DELOCALIZED π BONDS
- properties: soft and slipper, good conductivity