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From Flashcardlet > AP Chem Ch 10 > Flashcards

AP Chem Ch 10 Flashcards

0
Q

Exceptions to octet rule

A

B, Be, and Al (as a non-metal) are stable with 6 or 4 electrons
N is often stable with 7 electrons
P and larger atoms are often stable with 10 or 12 electrons

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1
Q

Method for Lewis Dot Structure

A
  1. Add up valence electrons
  2. Re count valence electrons
  3. Arrange atoms. Least electronegative atom in the center
  4. Use pairs of electrons to draw in single bonds
  5. Put the rest of the electrons on the non central atoms trying to form octets
  6. Put extra electrons on central atom
  7. Check if central atom has octet
  8. Form double/triple bonds to give central atom an octet
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2
Q

Formal charge

A

valence e- - #bonded electrons/2 - # of electrons in a lone pair

Ex. Oxygen in a double bond:
6-4/2-4

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3
Q

London dispersion forces

A

Induced/temporary dipole
Increases with size–larger the molecule, bigger the LDF
Type of IMF for non-polar molecules
As you go down a group, freezing and boiling points increase because bigger LDF. Stronger IMF allows for this increase

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4
Q

Polarizability

A

The freezing point rises going down the group. Principle reason for this is because as the atomic number goes up, number of electrons GO up, so there is an increased chance of the occurrence of momentary dipole interactions. We describe this phenomenon using the term polarizability, which indicates the ease with which the electron cloud of an atom can be distorted to give a dipolar charge distribution. Thus we say that large atoms with many electrons exhibit a higher polarizability than small atoms. So, LDF increases as size increases

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5
Q

Dipole-Dipole

A

Molecules with a permanent dipole (polar species) desire to align to reduce repulsions. More favorable–less energy required–when a polar molecule interacts with another polar molecule.

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6
Q

H-bonding

A

Extreme case of dipole-dipole where Hydrogen is attracted to a lone pair in an N, O, or F atom.
Very strong IMF because of the big electronegativity difference–strong polar bond between H-N, H-O, H-F
Also occurs because the dipoles are able to closely approach one another because hydrogen is so small.

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7
Q

Effects of hydrogen bonding

A

Since the IMF is so strong, higher boiling and melting points, less vapor pressure than dipole-dipole or LDF species

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8
Q

Why do polar molecules mix with polar molecules and not with non-polar molecules?

A

Polar– dipole produced with a partial positive and a partial negative side. Attracted to other polar molecules where partial positive is attracted to partial negative because this is a favorable reaction, can just bond, doesn’t take too much energy
However, this partially charged molecule won’t interact with a non-charged non-polar molecule unless a lot of energy is induced, so this is not favorable

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9
Q

Surface tension

A

The resistance of a liquid to an increase in its surface area. Liquids with large IMFS have high surface tension.

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10
Q

Capillary action

A

Spontaneous rising of a liquid in a narrow tube.
Caused by 2 different types of forces:
1. Cohesive forces– IMF among the molecules of the liquid
2. Adhesive forces– forces between liquid and container
Adhesive forces occur when the container is made of a substance with polar bonds so the water can attract to it.
The ability of water to wet the glass makes it creep up the walls of the tube where the water surface touches the glass. Increases the surface area of the water, which is opposed by the cohesive forces trying to minimize the surface. So, with the strong adhesive and cohesive forces, water pulls itself up a glass capillary tube to a height where the weight of the column of water just balances the water’s tendency to be attracted to the glass surface.

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11
Q

Cohesive vs adhesive forces

A

Cohesive among molecules of a liquid

Adhesive with liquid and the container

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12
Q

Concave vs convex meniscus and reason why

A

Polar has a concave meniscus because of strong adhesive forces.
Attracted to the glass
Nonpolar is convex because of strong cohesive forces. Won’t interact with the polar glass tube.

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13
Q

Viscosity

A

Measure of the resistance to flow.
Larger IMFS have larger viscosities
Also more sticky with higher viscoscities

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14
Q

Solid vs gas vs liquid

A

Solid is rigid structures with essentially no molecular motions
Gas has lots of motion. No IMFs
Liquid the hardest to analyze because of both motion and IMFS

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15
Q

X - ray diffraction

A

Diffraction occurs when beams of light are scattered from a regular array of points in which the spacing between the components are comparable with the wavelength of the light.
Allows for structure of crystalline solid to be determined

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16
Q

Constructive vs destructive interference

A

Constructive when the waves of parallel beams are in phase. Produces a spot on the detector
Destructive when the waves are out of phase. No image produced.

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17
Q

Importance of diffraction

A

Allows for bond lengths to be determined from the Bragg eq

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18
Q

Bragg equation

A 
n (lambda) = 2d * sin (theta)
n is an integer
Lambda is the wavelength
d is the disfrace between points
Theta is the angle of incident light
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19
Q

Example of Bragg equation
If the wavelength of an Al crystal is 1.54 * 10^-10 m and it produces reflection at an angle of 19.3 degrees and we assume n=1, calculate the distance between the planes of atoms producing this reflection

A

Use Bragg equation:
1(1.54*10^-10)=2dsin(19.3)
d = 2.33 * 10^-10 m

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20
Q

Molecular orbital theory for metals

A

Bonding in most metals is both strong and nondirectional because they are durable and have high melting points.
Electron sea model explains how there is a regular array of metal cations in a sea of valence electrons. Mobile electrons can conduct heat and electricity and the metal ions can be easily moved around as the metal is hammered into a sheet or pulled into a wire
A related model with more detailed view of the electron energies and motions is the band model, or molecular orbital model, for metals. In this model, the electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal atoms. Large number of resulting molecular orbitals are closely spaced and finally form a virtual continuum of levels, called bands.

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21
Q

How does temperature affect conductivity of a metal?

A

T up, so more collisions, so less conductivity.

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22
Q

Band theory to explain metals conducting

A

Conduct electricity and heat very efficiently because of the availability of highly mobile electrons. Electrons in partially filled bonds are mobile, so they can flow and create charge. The molecular orbitals occupied by these conducting electrons are called conduction bands. These account for the efficiency of the conduction of heat thru metals. When one end of a metal rod is heated, the mobile electrons can rapidly transmit the thermal energy to the other end.

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23
Q

Which has a bigger gap between conductive and valence band?

A

Insulator bigger gap than semiconductor

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24
Q

Closest packing

A

The structure in crystalline solid materials. Packing uniform hard spheres in a manner that most efficiently uses the available space.

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25
Q

Aba arrangement

A

Hexagonal closest packed (hcp) – hexagonal unit cell. Third layer lies directly over the first because of dimples between first and second, and second and third layers.

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26
Q

ABC arrangement

A

Cubic closest packed (ccp) structure. Face centered cubic unit cell. Fourth layer is like the first. Abcabc
FCC has 4 atoms in 1 unit cell.
And one side you can do by the diagnaol, and then pi that, so d^2 + d^2 = (4r)^2

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27
Q

Simple cubic, body-centered cubic, face centered cubic

A

SC– 1 atom per unit cell. 2r= one side
BCC– 2 atoms in a unit cell.
FCC– 4 atoms in a unit cell. One side = 4r/SQRT2

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28
Q

Ex. Silver crystallizes in a CCP structure. Radius is 144 pm

Find the density of solid silver

A

CCP structure, so 4 atoms in a unit, so (4r)^2 = d^2 + d^2
d = 144SQRT 8 = 407 pm
Volume = side cubed = 6.7410^7 pm^3 * 1/10^30 cm^3 = 6.7410^-23 cm^3
Mass–we know that there are four atoms–> 107.9 g/mol * 1 mol/6.0210^23 afoms * 4 afoms = 7.17*10^-22 g
Mass / volume = 10.6 grams per cubic cm

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29
Q

Substitutional alloy

A

When a metal has equal in size secondary metal.
Host atoms are replaced by other metal atoms of similar size
Brass, sterling silver

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30
Q

Interstitial alloy

A

Some holes in the closest packed metal structure are occupied by small atoms, most of the time, carbon. Changes properties and makes it stronger

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31
Q

Network solids

A

Many atomic solids contain strong directional covalent bonds to form a solid that might best be viewed as a giant molecule.
Brittle, do not efficiently conduct heat or electricity.
Carbon– diamond, graphite
Silicon– silicates, silica
Big networks of Si and O

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32
Q

Ionic solids

A 
Metal and nonmetal
Ionic bonds
Latrice
High boiling and melting points because of high electrostatic attraction
Very soluble in aqueous solution
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33
Q

Molecular solid

A

Water
Covalent bonds but really IMFS holding together the structure, hence the weak structure and lower boiling and melting points.
Molecular structure, not formula units

34
Q

Semiconductor

A

Silicon is a semiconductor–a few electrons can cross the gap between the filled and empty molecular orbitals at 25 Celsius, making silicon a semiconducting element, or a semiconductor. Additionally, at higher temperatures, when more energy is available to excite electrons into the conduction bands, the conductivity of silicon increases. Typical for semiconductor, contrast to that of metals, whose conductivity decreases with increasing temperature.

35
Q

p-type semiconductor

A

Positive. Species introduced with fewer electrons than host to increase conductivity. Creates holes–absensce of an electron. Emory orbital.
Holes are places of net positive charge

36
Q

p-type semiconductor

A

Positive. Species introduced with fewer electrons than host to increase conductivity. Creates holes–absensce of an electron. Emory orbital.
Holes are places of net positive charge

37
Q

P-n junction

A

Where p and n semiconductors meet

38
Q

P-n junction

A

Where p and n semiconductors meet

39
Q

What happens when p and n type semiconductors are next to each other with no battery?

A

Electrons migrate from the n type to the p type by filling in the holes. Leaves the n type with holes and a net positive charge and the p type with extra electrons and net negative charge. Further migration of electrons is prevented–charge buildup results in stop of migration– contact potential

40
Q

What happens when p and n type semiconductors are next to each other with no battery?

A

Electrons migrate from the n type to the p type by filling in the holes. Leaves the n type with holes and a net positive charge and the p type with extra electrons and net negative charge. Further migration of electrons is prevented–charge buildup results in stop of migration– contact potential

41
Q

What happens if we put the negative end of a battery by a p-type semiconductor and the positive end by the n-type

A

Then, electrons in the n-type are drawn toward the positive terminal. Holes in the p-type towards the negative terminal
Electrons don’t flow across the junction
Reverse bias
No current flow

42
Q

What happens if we put the negative end of a battery by a p-type semiconductor and the positive end by the n-type

A

Then, electrons in the n-type are drawn toward the positive terminal. Holes in the p-type towards the negative terminal
Electrons don’t flow across the junction
Reverse bias
No current flow

43
Q

What happens when positive end of battery placed by p-type and negative by n-type?

A

The electrons in the n-type are drawn toward the positive terminal but must flow across the junction filling the holes in the p type.
Holes move towards the negative terminal
Forward bias
Results in current flow
Movement of holes and electrons is favorable.
Important for circuits

44
Q

What happens when positive end of battery placed by p-type and negative by n-type?

A

The electrons in the n-type are drawn toward the positive terminal but must flow across the junction filling the holes in the p type.
Holes move towards the negative terminal
Forward bias
Results in current flow
Movement of holes and electrons is favorable.
Important for circuits

45
Q

Rectifiers

A

Devices that produce a pulsating direct current from an alternating current because current only flows in forward bias.

46
Q

Rectifiers

A

Devices that produce a pulsating direct current from an alternating current because current only flows in forward bias.

47
Q

Dc vs ac

A

D.C. In one direction

AC in both directions

48
Q

Dc vs ac

A

D.C. In one direction

AC in both directions

49
Q

What happens when water is placed in a eurleymeyer flask?

A

Initially, some water evaporates/vaporizes into has but then some gas turns back into liquid–>condensation.
Process continues back and forth until equilibrium reached, when the rate of the forward reaction= rate of the reverse reaction

50
Q

Equilibrium

A

When the rage of vaporization = raye of condensation

Dynamic process but no net change

51
Q

Vapor pressure

A

The pressure of the vapor present at equilibrium
Patm = Pvapor + P Hg column
Pvap= Patm-P Hg

52
Q

Rates of evaporation and consenation in a closed container

A

As time goes on, rate of evaporation remains constant the whole time but rage of condensation increases until the rates are equal
Because it starts off with more liquid to gas, but then gas to liquid rises until equal. No net change

53
Q

Heat of vaporization

A

The energy required to vaporize 1 mol of a liquid at 1 atm

Requires energy, endothermic, because energy is needed to overcome IMFS

54
Q

How does temperature influence vapor pressure?

A

Increases significantly with temperature.
To overcome the IMFS in a liquid, a molecule must have sufficient kinetic energy. As the temperature is increases, the fraction of molecules having the minimum energy needed to overcome these forces and escape to the vapor phase increases markedly. Thus the vapor pressure of a liquid increases dramatically with temperature.

55
Q

How do IMFS relate to vapor pressure?

A

Inversely proportional
Liquids with large IMFS have low vapor pressures because the molecules need high energies to escape to the vapor phase.
Strong H-bond in water causes it to have a low vapor pressure.
For non-polar, more massive have less vapor pressures because with the extra electrons, it is more polarizable, and has bigger ldfs

56
Q

Temperature and vapor pressure equation

A

We can make a linear relation plotting ln(Pvap) vs 1/T (1/Kelvin)
ln(Pvap) = -heat of vap/R (1/T) + C

57
Q

Slope of ln(Pvap) vs 1/T graph

A

Slope is - heat of vaporization / R,

So the steeper slope will have the larger heat of vaporization

58
Q

Practice problem:
Vapor pressure of water at 25 degrees Celsius is 23.8 torr, delta H vap = 43.9 kJ/mol
Calculate vapor pressure of water at 65 degrees Celsius

A

At 25 degrees, ln(23.8)=-43.9/8.31410^-3 (1/298) + c
C= 20.89
Now plug into same equation at 65 degrees
ln(Pvap) = -43.9/8.31410^-3 (1/338) + 20.89
Pvap = 194.4 torr

59
Q

Hearing curve for water

A

Plot of temperature vs heat added at a constant rate
First starts out as a solid, temp increases to the melting point, then remains constant as melts, then increases to the boiling point, then remains constant as it vaporizes, then increases as a gas

60
Q

Phase change

A

Alters physical properties, not chemical

Involves IMFS, not bonds

61
Q

Enthalpy of fusion

A

Enthalpy change that happens at the melting point

62
Q

Solid and liquid water vapor pressure graph vs temperature

A

Melting and boiling points for a substance determined by the vapor pressures of the solid and liquid states. Below 0 degrees, solid has lower vapor pressure but above 0 degrees, solid has larger vapor pressure. At 0 degrees they are equal. So, the solid vapor pressure has a larger slope–increases more rapidly for a given rise in temperature

63
Q

Case 1: a temperature at which the vapor pressure of the solid is greater than that of the liquid

A

The solid requires a higher pressure than the liquid does to be in equilibrium with the vapor. So, as the vapor is released from the solid to try to achieve equivilibrium, the liquid will absorb vapor to try to reduce the vapor pressure to its eq value. Net effect is a conversation from solid to liquid thru the vapor phase. No solid can exist under these conditions. Solid decreases as volume of liquid rises until there is only liquid and it will come to eq with the water vapor. No further changes. Must be above the melting point

64
Q

Case 2: when the temperature is such that the vapor pressure of the solid is less than that of the liquid

A

Here, the liquid requires a higher pressure to be in equilibrium with its vapor phase. So, the liquid will gradually disappear as the solid increases. Only solid will remain. Must be below melting point.

65
Q

Case 3: temperature at which the vapor pressure of the solid = vapor pressure of the liquid

A

Here, they have the same vapor pressure, so they can coexist at eq simultaneously with the vapor. This is the freezing point. Both states coexist.

66
Q

Normal melting point

A

The temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure = 1 Atm

67
Q

Normal boiling point

A

Boiling occurs when the vapor pressure of a liquid equals the pressure of its environment. Normal boiling point is the temperature when the vapor pressure of the liquid = 1 atm
Boiling point lowers as altitude increases

68
Q

Supercool

A

Water can be supercooled–it can be cooled below 0 degrees Celsius at 1 atm pressure and remain in liquid state. Supercooling occurs because as it is coooled, the water may not achieve the degree of organization necessary to form ice at 0 degrees, and continues to exist as liquid. At some point, the correct ordering occurs and ice forms, releasing energy and bringing the temperature back up to the melting point, where the remainder of the water freezes

69
Q

Superheat

A

When a liquid is raised to temperatures above its boiling point

70
Q

Phase diagram

A

A convenient way of representing the phases of a substance as a function of temperature and pressure.

71
Q

Exp 1: what happens to ice at -20 degrees, p=1 ATM, as it is heated

A

Solid to liquid at 0 degrees and then at 100 degrees becomes vapor

72
Q

Exp 2: when ice is heated from -20 degrees at a very low pressure of 2 torr

A

Sublimation because vapor pressure of ice is equal to external pressure. Because the vapor pressure of liquid water is greater than 2 torr, it cannot exist at this pressure

73
Q

Exp 3: low pressure of 4.58 torr, ice heated from -20 degrees

A

No new phase until T = .01 Celsius

Triple point–all 3 states of matter co exist in a closed system. Solid and liquid states have same VP of 4.85 torr

74
Q

Exp 4: liquid water at 300 degrees Celsius and pressure at 225 atm

A

High ester all pressure, water can be a liquid. As temperature increases, liquid changes into a vapor thru a fluid phase – neither liquid nor vapor

75
Q

Critical temperature

A

The temperature above which the vapor cannot be liquefied, regardless of pressure

76
Q

Critical pressure

A

The pressure required to produce liquefaction at the critical temperature

77
Q

Critical point

A

Cortical temperature and critical pressure
Liquid/vapor line ends at this point. Because transition from one state to another involves the intermediate fluid region above this point

78
Q

N-type semiconductor

A

Negative

Increase conductivity by doping it with atoms with more valence electrons than the host

79
Q

Why does the solid/liquid line for water have a negative slope?

A

The melting point of ice decreases as the external pressure increases. This occurs because the density of ice is less than that of liquid water at the melting point. The maximum density of water occurs at 4 degrees Celsius, when liquid water freezes, its volume increases.
Normally, this line is positive because solid density greater than liquid

80
Q

How does altitude influence cooking and boiling?

A

At higher altitude, water boils at a lower temperature, so takes longer to cook because not as much energy in the system

81
Q

What happens with water when not in a closed system?

A

In dry climates (low humidity), snow and ice seem to sublime
Ice in the natural environment is not a closed system. So the pressure is produced by the atmosphere, not the solid piston. So, the vapor produced over the ice can escape from the immediate region as soon as its formed. The vapor does not come to eq with the solid, and the ice slowly dissapeads. Sublimation, which seems forbidden by the phase diagram, does in fact occur under these conditions, although it is not the sublimation under eq conditions described by the phase diagram

82
Q

Carbon dioxide phase diagram

A

Solid/liquid line has a positive slope since solid CO2 is more dense than liquid.
Triple point at 5.1 atm and -56.6 Celsius
Critical point at 72.8 atm and 31 degrees Celsius.
At 1 atm, solid carbon dioxide sublimes at -78 degrees Celsius, a property that leads to its name, dry ice.

83
Q

Carbon dioxide in fire extinguisher

A

Exists as liquid at 25 clecius under high pressure. Liquid released from extinguisher changes into a vapor immediately. Heavier than air, smothers the fire by keeping oxygen away from the flame. Endothermic reaction, so cooling also results, helping put out the fire.

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