Test 3 Study Guide Flashcards
(95 cards)
1
Q
What is wavelength?
A
the distance between one peak and the next
*the measurement of wavelength is between peak to peak or trough to trough
2
Q
As wavelength increases, the frequency?
A
decreases
3
Q
What is the speed of light?
A
c=wavelengthXfrequency=300x10^8 m/s
4
Q
What is the unit of frequency?
A
per second
5
Q
1 hertz=
A
1 sec^-1
6
Q
100.3 MHz=
A
wavelength
7
Q
What is the principle quantum number?
A
N
8
Q
What symbol is the angular momentum quantum number?
A
L (cursive)
9
Q
What is the symbol for the magnetic quantum number?
A
M sub e
10
Q
What is the symbol for the spin quantum number?
A
M sub s
11
Q
What does the principle quantum number tell us?
A
energy or distance from nucleus
12
Q
What numbers is the principle quantum number (n) allowed to be?
A
1,2,3,4
13
Q
And N (principle quantum number) increases, what does E do?
A
increase
14
Q
What does L (angular momentum quantum number decribe?
A
Tells the shape
15
Q
What is the angular momentum quantum number allowed to be?
A
n-1 but 0 is the least it can be
16
Q
If you know n and L, then what are you describing?
A
a subshell
17
Q
What letters go along with L for the various subshells? 0 1 2 3
A
0-s
1-p
2-d
3-f
18
Q
What is the highest subshell?
A
f or 3
19
Q
Give the protocol for naming a subshell
A
use value->nl
20
Q
What does the magnetic quantum number tell?
A
Orientation in 3D space
21
Q
If we know n, L, and M sub l we are describing?
A
an atomic orbital
22
Q
Give the notation for each of the following orbitals. Determine if it is allowed, if not why?
a) n=4, L=1, m sub L=0
A
a)allowed 4p
23
Q
Give the notation for each of the following orbitals. Determine if it is allowed, if not why?
b) n=2, L=2, m sub L = -1
A
Not allowed, because n & l cannot be equal
24
Q
Give the notation for each of the following orbitals. Determine if it is allowed, if not why?
n=5, L=3, m sub L= +3
A
Allowed, 5f
25
What does the spin quantum number tell us?
about magnetic properties
26
What is # are allowed in the spin quantum number?
+1/2
| -1/2
27
If we know n, L, m sub L, and m sub s, then we are describing?
an individual electron
28
The s orbital makes what shape?
spherical
29
The p orbital makes what shape?
a infinity symbol
30
The d orbital makes what shape?
a flower w/4 petals shape
31
The f orbital makes what shape?
a flower with many petals
32
Nature looks for?
the lowest energy arrangement possible
33
What is effective nuclear charge?
weighted average of nuclear charge that affects an electron in the atom, after correction for the shielding by inner electrons and interelectronic repulsions
34
The _________ the affected nuclear charge, the more able the nucleus is able to control the electrons and vice versa
the greater the affected nuclear charge
35
Give the order of the subshells-
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<5d<6p etc.
36
What is the Pauli exclusion principle?
No two electrons in the same atom can have the same set of four quantum numbers
37
What is the Aufbau Principle?
as electrons are added to an atom one at a time they are assigned the quantum numbers of the lowest energy orbital that is available
38
The ground state is?
the lowest energy state possinle
39
What is hund's rule?
One electron occupies earch degenerate orbital with the same spin before a second electron is placed in an orbital
40
Group 1A atoms always form?
1+ cations
41
Which group of atoms is particularly stable?
Noble gases
42
What are valence electrons?
those with the highest principle quantum number and any electrons in a unfilled subshell from a lower shell
43
What is an ion?
charged particle
44
A cation is
forms a positive charge from loosing an electron
45
An anion is
negative charge from an atom gaining an electron
46
What is an isoelectronic series?
is a group of atoms and ions that contain the same number of electrons
47
What is an example of an isoelectronic series?
S2-, Cl-, Ar, K+, and Ca2+
48
How is the size of an atom determined?
by measuring its radius
49
What type of atom has the smallest size?
Species with the highest number of protons
50
Atoms are always larger than their _______.
Atoms are always larger than their cations
51
Atoms are always smaller than their _______.
Atoms are always larger than their anions.
52
Effective nuclear charge increases?
up and to the right
53
As zeff increases, size?
decreases
54
Atomic radi increase?
down and to the left of the periodic table
55
What is ionization energy?
the energy required to remove an electron
56
1st ionization Energy increases?
1st ionization Energy increases to the right and from the bottom to top
57
What are the exceptions to 1st ionization energies?
Group IIIA and VIA exceptions
58
An isoelectronic species with the greatest charge in the nucleus will have?
the largest ionization energy
59
Successive Ionizations are always?
larger than 1st ionization energies
60
What produces a much larger ionization energy?
when a non-valence electron is removed
61
What is Electron affinity?
energy change when a electron is added to an atom to a gaseous atom to form an anion
62
What is the equation for electron affinity?
A(g) + e- ->A-(g)
63
The reactivity of the Group 1A metals increases?
The reactivity of the Group 1A metals increases down the group
64
What are alloys?
mixture of metals
65
The halogens all exist as?
diatomics
66
Halogen reactivity increases as you go?
up the group
67
What is ionic bonding?
+/- attracted to form bonds
68
What is covalent bonding?
The sharing of electrons
69
What is lattice energy?
the energy required to seperate one mole of ionic solid into its gaseous ions
70
*********
| As charge goes up, lattice energy goes?
up
71
```
*************
as r (distance between ions) goes up, then lattice energy goes?
```
down
72
To arrange by increasing lattice energy you?
assess charge first and then distance
73
Ionic bonding is?
a metal and a nonmetal cations and anions electron transfer
74
What is bond order?
the number of electron pairs shared between two atoms
75
What is the bond order below?
a) N-N
b) N=N
a) 1
| b) 2
76
As bond order increases, bond length?
decreases
77
As bond order increases, bond strength?
increases
78
What are the steps to writing Lewis Dot structure
1) Skeleton structure
2) Sum valence electrons
3) Subtract 2e- for every bond in skeleton structure
4) Count electrons needed to complete (need) structure
5) Add extra electron as base pai
79
If the electron sharing is equal, what type of bond is it?
a nonpolar convalent bonding
80
If electron sharing is not equal, what type of bond is it?
polar covalent bonding
81
What is electronegativity?
a measure of the ability of an atom to attract the shared electrons in a chemical bond
82
Electronegativity increases?
up and to the right
83
As the difference in electronegativity increases, the polarity of the bond?
increases
84
The most polar bond will have?
the biggest difference in electronegativity
85
The central atom is?
the least electronegative element
86
How do you calculate formal charge?
The # of valence electrons - # of bonds - # of lone electrons = formal charge
87
What Lewis structures are favored?
the structures of the smallest formal charge
88
What are the rules for evaluating the stability of structures?
Smallest formal charge favored
Structures with adjacent atoms with formal charges of the same sign are much less favored
Lewis structures that place negative formal charges on more electronegative atoms are favored
Formal charges of opposite signs are usually on adjacent atoms
89
What is a resonance structure?
when a molecule has more than 1 formal charge
90
No resonance structure is correct by itself, how is the correct structure formed?
By the average between all the resonance structures
91
What molecules disobey the octet rule?
```
H needs 2e- to be stable
He needs 2e- to be stable
2A&3A central atoms don't need all electrons
Group IIA central atoms need 4e-
Group IIIA central atoms ned 6e-
Expanded valence shell z>15
```
92
____ formal charge is more favored than ___ formal charge.
No formal charge is more favored than any formal charge.
93
An expanded valence shell molecules have?
more than eight electrons about an atom in a Lewis structure
94
bond energy is?
the energy required to break bond
95
Bond energies are always?
endothermic because it takes energy to break a bond
