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CHEM > Exam 2 > Flashcards

Exam 2 Flashcards

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1
Q

Frequency

A

V

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2
Q

Wavelength

A

λ

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3
Q

Frequency measured in

A

hertz

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4
Q

Wavelength measured in

A

nm

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5
Q

Speed of light (on formula sheet)

A

C

3.00 * 108 m/s (

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6
Q

Energy and Frequency

A

Directly proportional

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7
Q

Energy and Wavelength

A

Inversely proportional

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8
Q

Wavelength and Frequency

A

Inversely Proportional

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9
Q

Energy of a photon equations (2 equations)

A

E= h * v (on formula sheet)

E= h*c/λ

E=J/Photon

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10
Q

What is h?

A

Planck’s constant

6.26*10-34J/s

(on formula sheet)

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11
Q

convert m to nm

convert m to um

A

1m= 1*109nm

1m=1*106um

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12
Q

ROYGBIV

A

Red low energy, longest wavelength. (infrared next to red)

Violet is highest energy, shortest wavelength. (UV next to violet)

ROYGBIV

Energy ——>

Frequency ——>

<————wavelength

Remember: Wavelength is inversely proportional!

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13
Q

Avagadros for light problems!

A

1 mol of photons/6.022*10^23 photons

Take something like #KJ/mol. Convert KJ to J then multiply by 1mol/6.022*1023photons so you get J/photon that can be used in E equations. E=hc/λ

Or if you have number of J/photon and want it in moles, multiply by 6.022*1023photons/1mole.

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14
Q

E=hc/λ

λ should be what unit?

A

METERS

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15
Q

q=mc∆T

A

c=4.184 J/gC heat capacity of water

m is mass in grams

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16
Q

Electron moving to a different energy level

A

only if it absorbs or emits a photon that has the same energy as the difference between the two energy levels

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17
Q

ground state

A

lowest energy oribtal n=1 for hydrogen

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18
Q

Excited State

A

When the electron is at a higher energy orbital n=2+ for hydrogen

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19
Q

Ionize

A

When an electron is ejected from an atom. Nfinal=infinity.

Losing an electron.

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20
Q

Exothermic and endothermic electrons

A

Electrons go from a high level (7) to a low level (2). ∆H= neg. Emission!

Endothermic go from low (2) to a high level (7). Absorption!

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21
Q

Lyman series

A

UV 90-100nm

ends at n=1

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22
Q

Balmer Series

A

400 to 700

Ends at n=2

Visible Light

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23
Q

IR series

A

IR 1000 nms

ends at n=3

all the series are exothermic BTW. High to low level.

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24
Q

Rydberg Equation Wavelength of electron transition in a Hydrogen atom

(all given on formula sheet)

A

1/λ =R(1/n12-n22)

n2>n1!!

R=1.097*107

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25
Q

Rydberg for energy transition (other one was to calculate wavelength)

A

RHC*Z2(1/nfinal2-1/ninitial2)

z=# of protons element has – so 1 for hydrogen.

RHC= -2.18*10-18

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26
Q

wavelength, mass, and speed– de broglie

A

λ =h/mu

mass in kg, u is speed in m/s and h is planck’s.

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27
Q

emissions at n=1

A

anything going or coming to n=1 has highest E. Passes through lots of levels. Or could use the rhc equation.

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28
Q

n

A

level (look at periodic table row).

Indicates relative distance

higher th enumber, greater the distance of the orbital from the nuclues, larger the shell, higher energy.

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29
Q

l

A

subshell

max value: of n-1. Max number of possible l values=n

Indicates shape

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30
Q

s, p, and d l values and shape

A

s =0; sphere

p=1; dumbbell

d=2; double dumbbell

f=3; flower

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31
Q

ml

A

-1 to +1 range.

Orientation.

labeling the slotsof the orbitals e.g. if p: -1, 0, 1 for the 3 slots.p l value is 1.

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32
Q

Pauli Exclusion principle

A

max number of e per orbital is 2

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33
Q

Hund’s rule

A

Fill with one e in each orbital, then go back and fill in the rest.

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34
Q

ms

A

-1/2 or +1/2

if up arrow, positive.

represents magnetic field

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35
Q

testing quantum validity

A

n>l> or equal to |ml|.

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36
Q

pauli exclusion principle

A

no 2 e in the same atom can have the same 4 quantum numbers.

2 e in an orbital must have opposing signs

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37
Q

aufbau principle

A

filling orbitals with the lowest energy

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38
Q

Z or nuclear charge

A

electrong is attracted to teh positive charge of the nucleus

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39
Q

Electron Shielding

A

If there are several electrons in the same orbital, they repel each other and shield one another from the attractive force of the nucleus (prevents feeling of full nuclear charge-Z). Experience an effective nucleuar charge (Zeff) due to shiledings. Effective shielding done by inner electrons loweres the effect of Zeff on outer electrons.

There is more shielding in the p orbital- easier to remove a p orbital electron.

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40
Q

electron removal order

A

remove from highest n first, then from f, d, p, s.

41
Q

Inner/core electrons

A

noble gas electrons and any completed transition series (d10 or f14)

42
Q

valence electrons equation for transition metals

A

number of electrons-core

43
Q

para vs. diamagnetic

A

paramagnetic- unpaired

diamagnetic- all electrons are paired.

44
Q

Transition orbital filling exceptions

A

most important: copper and chromium and palladium and silver

45
Q

atomic size

A

left and down.

46
Q

Zeff and shielding

A

zeff: up and right.

(more protons)

Shielding is the opposite. Down and left increases shielding.

47
Q

ion size

A

look at what they are isoelectric with, compare the size of that.

If still tied,

Then, look at number of protons in the original ion to decide. (more protons, larger Zeff, smaller size, large Ionization energy)

Remember: cations (+) are smaller than anions (-).

48
Q

Ionization energy trend

A

up and to the right increases IE (due to Zeff holding them tighter)

49
Q

successive ionization energies

A

IE greatly increases when removing an inner (core) electron.

50
Q

metal and nonmetallic character

A

think about location of metals and nonmetals on periodic table. Follow that trend.

51
Q

covalents make acids or bases

A

acids

so acidity increases with nonmetallic character (up and right)

52
Q

metals make acids or bases?

A

Bases and basciity increases as you go down and to the left on periodic table (with metal trend)

53
Q

Redox behavior

A

Group 1 and 2 are losing electrons (LEO), oxidizing, so strong reducing agents.

Group 6 are gaining electrons, so strong oxidzing agents.

54
Q

Electron Affinity

A

Energy released when an electron is added to a neutral gas atom. EA is -.

Increase in negativity as you go up and to the right.

55
Q

1st Ea vs. 2nd EA

A

1st is exothermic and -, second is endothermic and a + ∆H

56
Q

Lattice energy

A

energy absorved when 1 mol of an ionic compound breaks up into gas phase ions

57
Q

What impacts lattice energy (strength of ionic bond)

A

1) ionic charge 2) ionic size

greater LE with greater charger

LE decreases as the ionic size increases

Look at anion size first. you want the smaller anion– greater LE

58
Q

Steps of born haber cycle

A

sublimation, IE, bond dissociation, EA1, Lattic energy

59
Q

sublimation

A

Li(s)—> Li(g)

+

60
Q

IE1

A

Li(g)—>Li+ + e

+

61
Q

bond dissociation

A

1/2 F2(g)—>F(g)

+

62
Q

EA

A

F(g)+e—>F-1

-

63
Q

lattice energy

A

Li+(g)+F1-(g)—> LiF(s)

-

64
Q

Formation

∆H of f

A

Li(s)+1/2F2(g)—>LiF(s)

-

add up steps to get this

65
Q

Electronegativity

A

up and to the right

ability to attract electrons.

Or drop by .5 going tot he left and 1/2 it going down.

66
Q

H electronegativity

A

2.1

67
Q

I electronegativity

A

2.5

68
Q

Br electronegativity

A

2.8

69
Q

Cl electronegativity

A

3.0

70
Q

N electronegativity

A

3.0

71
Q

O electronegativity

A

3.5

72
Q

F electronegativity

A

4.0

73
Q

Be and B and H

A

Be only needs 4 e in octet

B only 6

H 2

74
Q

Formal charge

A

Valence electrons-dots-sticks

negative charge on most EN element, and positive on least EN.

75
Q

Bond orders for resonance

A

single- 1

double- 2

resonance bond order= # of bonds/# of elements around central atom

76
Q

what is the longest bond?

A

single bond

77
Q

what is the weakest bond?

A

single bond

78
Q

The shortest and strongest bond?

A

triple

79
Q

Calculating ∆H from bond energies

A

Bonds broken (reactants)- bonds formed(products)

80
Q

VSPER Shortcut

A

subtract highest multiple of 8— divide that by 8 to get the X

take the remainder from subtraction and divide by 2 to get lone pairs E.

81
Q

2 groups

A

sp

linear

AX2

82
Q

3 groups

A

Group: trigonal planar sp2

AX3-trigonal planar

AX2E1-bent

83
Q

4 groups

A

tetrahedral-sp3

AX4-tetrahedral

AX3E1-trigonal pyramidal

AX2E2-Bent

84
Q

5 groups

A

trigonal bipyramidal sp3d

AX5 trigonal bipyramidal

AX4E1-see saw

AX3E2- t shaped

AX2E3-linear

85
Q

6 groups

A

sp3d2 Octahedral

AX6 octahedral

AX5E1 square pyramidal

AX4E2 square planar

86
Q

bond angles for linear trigonal planar and tetrahedral

A

linear- 180

trigonal planar- 120

tetrahedral-109.5

87
Q

small elements vs. large elements and bond size

A

small elements make short/strong bonds and large elements make long/weak bonds

88
Q

most polar

A

greatest difference in electronegativity

89
Q

most polar type of bond

A

ionic bond!

metal and nonmetal have biggest differences in EN.

90
Q

molecular polarity test

A

is the geometry symmetric?

are the outer atoms the same?

91
Q

symmetric molecular shapes

A

all the non E pair geometries and the Ax2E3 linear and the Ax4E2 square planar

92
Q

multiple center compound polarity

A
  • any lone pairs? automatically polar.
  • Different configurations on central atoms makes it polar
  • same configurations and same elements around it- nonpolar
93
Q

sigma bond

A

single localized bond. produced by overlap of two s’s, 2 sp’s, or an sp with an s.

94
Q

delocalize

A

sp orbital is created to delocalize the electron density around the central atom.

to move the electron density away from the middle of the lement- creates space for bonding and lowers energy of the element.

95
Q

Pi bonds

A

double or triple localized.

overlap of 2 unhybridized p creates a pi bond and 4 p orbitals creates 2 pi bonds

These both have sigma bonds.

pi bonds weaker than sigma bonds

96
Q

MO theory

A

delocalized. orbiatals are spread out, eliminates resonance and predicts stability.

97
Q

MO theory bond order

A

1/2 (nonstars-stars)

bonding–antidbonding

bond order>0 the molecule forms and exists

98
Q

greater mo theory bond order

A

greater the bond energy it has’more stable than separate atoms.

CHEM (5 decks)

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