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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
Rydberg for energy transition (other one was to calculate wavelength)
RHC\*Z2(1/nfinal2-1/ninitial2) z=# of protons element has -- so 1 for hydrogen. RHC= -2.18\*10-18
26
wavelength, mass, and speed-- de broglie
λ =h/mu mass in kg, u is speed in m/s and h is planck's.
27
emissions at n=1
anything going or coming to n=1 has highest E. Passes through lots of levels. Or could use the rhc equation.
28
n
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.
29
l
subshell max value: of n-1. Max number of possible l values=n Indicates shape
30
s, p, and d l values and shape
s =0; sphere p=1; dumbbell d=2; double dumbbell f=3; flower
31
ml
-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.
32
Pauli Exclusion principle
max number of e per orbital is 2
33
Hund's rule
Fill with one e in each orbital, then go back and fill in the rest.
34
ms
-1/2 or +1/2 if up arrow, positive. represents magnetic field
35
testing quantum validity
n\>l\> or equal to |ml|.
36
pauli exclusion principle
no 2 e in the same atom can have the same 4 quantum numbers. 2 e in an orbital must have opposing signs
37
aufbau principle
filling orbitals with the lowest energy
38
Z or nuclear charge
electrong is attracted to teh positive charge of the nucleus
39
Electron Shielding
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.
40
electron removal order
remove from highest n first, then from f, d, p, s.
41
Inner/core electrons
noble gas electrons and any completed transition series (d10 or f14)
42
valence electrons equation for transition metals
number of electrons-core
43
para vs. diamagnetic
paramagnetic- unpaired diamagnetic- all electrons are paired.
44
Transition orbital filling exceptions
most important: copper and chromium and palladium and silver
45
atomic size
left and down.
46
Zeff and shielding
zeff: up and right. (more protons) Shielding is the opposite. Down and left increases shielding.
47
ion size
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
Ionization energy trend
up and to the right increases IE (due to Zeff holding them tighter)
49
successive ionization energies
IE greatly increases when removing an inner (core) electron.
50
metal and nonmetallic character
think about location of metals and nonmetals on periodic table. Follow that trend.
51
covalents make acids or bases
acids so acidity increases with nonmetallic character (up and right)
52
metals make acids or bases?
Bases and basciity increases as you go down and to the left on periodic table (with metal trend)
53
Redox behavior
Group 1 and 2 are losing electrons (LEO), oxidizing, so strong reducing agents. Group 6 are gaining electrons, so strong oxidzing agents.
54
Electron Affinity
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
1st Ea vs. 2nd EA
1st is exothermic and -, second is endothermic and a + ∆H
56
Lattice energy
energy absorved when 1 mol of an ionic compound breaks up into gas phase ions
57
What impacts lattice energy (strength of ionic bond)
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
Steps of born haber cycle
sublimation, IE, bond dissociation, EA1, Lattic energy
59
sublimation
Li(s)---\> Li(g) +
60
IE1
Li(g)---\>Li+ + e +
61
bond dissociation
1/2 F2(g)---\>F(g) +
62
EA
F(g)+e---\>F-1 -
63
lattice energy
Li+(g)+F1-(g)---\> LiF(s) -
64
Formation ∆H of f
Li(s)+1/2F2(g)---\>LiF(s) - add up steps to get this
65
Electronegativity
up and to the right ability to attract electrons. Or drop by .5 going tot he left and 1/2 it going down.
66
H electronegativity
2.1
67
I electronegativity
2.5
68
Br electronegativity
2.8
69
Cl electronegativity
3.0
70
N electronegativity
3.0
71
O electronegativity
3.5
72
F electronegativity
4.0
73
Be and B and H
Be only needs 4 e in octet B only 6 H 2
74
Formal charge
Valence electrons-dots-sticks negative charge on most EN element, and positive on least EN.
75
Bond orders for resonance
single- 1 double- 2 resonance bond order= # of bonds/# of elements around central atom
76
what is the longest bond?
single bond
77
what is the weakest bond?
single bond
78
The shortest and strongest bond?
triple
79
Calculating ∆H from bond energies
Bonds broken (reactants)- bonds formed(products)
80
VSPER Shortcut
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
2 groups
sp linear AX2
82
3 groups
Group: trigonal planar sp2 AX3-trigonal planar AX2E1-bent
83
4 groups
tetrahedral-sp3 AX4-tetrahedral AX3E1-trigonal pyramidal AX2E2-Bent
84
5 groups
trigonal bipyramidal sp3d AX5 trigonal bipyramidal AX4E1-see saw AX3E2- t shaped AX2E3-linear
85
6 groups
sp3d2 Octahedral AX6 octahedral AX5E1 square pyramidal AX4E2 square planar
86
bond angles for linear trigonal planar and tetrahedral
linear- 180 trigonal planar- 120 tetrahedral-109.5
87
small elements vs. large elements and bond size
small elements make short/strong bonds and large elements make long/weak bonds
88
most polar
greatest difference in electronegativity
89
most polar type of bond
ionic bond! metal and nonmetal have biggest differences in EN.
90
molecular polarity test
is the geometry symmetric? are the outer atoms the same?
91
symmetric molecular shapes
all the non E pair geometries and the Ax2E3 linear and the Ax4E2 square planar
92
multiple center compound polarity
- any lone pairs? automatically polar. - Different configurations on central atoms makes it polar - same configurations and same elements around it- nonpolar
93
sigma bond
single localized bond. produced by overlap of two s's, 2 sp's, or an sp with an s.
94
delocalize
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
Pi bonds
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
MO theory
delocalized. orbiatals are spread out, eliminates resonance and predicts stability.
97
MO theory bond order
1/2 (nonstars-stars) bonding--antidbonding bond order\>0 the molecule forms and exists
98
greater mo theory bond order
greater the bond energy it has'more stable than separate atoms.

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