unit 4 Flashcards
(122 cards)
1
Q
why do electrons have discrete energies
A
they can only have certain wavelengths
2
Q
why are line spectra produced
A
transforming from one wavelength (waveform) to another
3
Q
general orbital meaning
A
cloudlike region of space around the nucleus where an electron has a high probability of being found
4
Q
energy level
A
identifies how close the electrons are to the nucleus
5
Q
sublevel
A
share the same energy level but have slightly different energies when compared to one another
6
Q
orbital in context to energy and sublevel
A
exists within sublevels and define the approximate boundaries of the electron’s orbit
7
Q
each electron is described by a set of ___, and what do they do
A
4 quantum numbers and these numbers provide info about the 3D region of space the electron is found in as well as the movement of the electron
8
Q
principle quantum number (n)
A
specifies the energy level of an atomic orbital and its relative size
9
Q
what is n determined by
A
the period in which an element is found on the periodic table
10
Q
the higher the number associated with a given energy level:
A
- the higher its energy
- the greater its size
- the more electrons it can hold
- the further away from the nucleus it extends
11
Q
secondary quantum number(l)
A
describes the shape(type) or orbital within each sublevel and its associated energy.
12
Q
l can be any integer ranging from
A
0 to (n-l)
13
Q
the number of sublevels in each principle energy level equals
A
the principle quantum number (n) for that energy level
14
Q
what’s the letter and name for sublevel with l=0
A
s - sharp
15
Q
what’s the letter and name for sublevel with l=1
A
p- principal
16
Q
what’s the letter and name for sublevel with l=2
A
d- diffuse
17
Q
what’s the letter and name for sublevel with l=3
A
f-fundamental
18
Q
correct sublevel order
A
0-s
1-p
2-d
3-f
19
Q
magnetic quantum number (m)
A
describes the orientation of an atomic orbital In space within a sublevel relative to the other orbitals within that same sublevel
20
Q
(m) can have any integer from
A
-l to +l
21
Q
spin quantum number(s)
A
the direction and electron spins in, either +1/2 or -1/2
22
Q
shapes of orbitals
A
s
p
d
f
23
Q
s shaped orbitals
A
spherical and have only 1 possible orientation
24
Q
p shaped orbitals
A
dumbbell shaped and have 3 possible orientations ( along x,y, or z plane)
25
d shaped orbitals
"clover leaf" shaped and have 5 possible orientations
26
f shaped orbitals
7 possible orientations
27
s block
group 1 and 2
28
p block
groups 13-18
29
d block
groups 3-12: transition metals
30
f block
lanthanides, actinides
31
key contributers to the quantum mechanical model of the atom
- de broglie
- heisenberg
- schrodinger
32
De Broglie contribution
- suggested that if light waves have particle like properties(photons) than why could it not be reasoned that electrons have wave like properties.
- his hypothesis was supported by the double slit experiment which demonstrated that moving electrons produce diffraction patterns similar to those that are produced by waves under certain conditions
33
hesienbergs contributions
contributed the Uncertainty principle
- due to the dual nature of electrons (particle and wave) it would be impossible to determine both an electrons position and momentum (speed) at the same time since any attempt to measure one property would interfere with or alter the other
34
Schrodingers contribution
-used mathematical wave equations from quantum physics to describe the behaviour of the electron as a wave
-solutions to this equation indicate where the probability of finding an electron is high to where it is comparitively low for any energy level
- electron probability density graphs provide a visual representation of these regions
35
energy level diagrams
method of showing how electrons are distributed amongst the orbitals of each energy level
36
electron configuration
shorthand notation used to indicate the number of electrons in the sublevels of each energy level
37
what are the 4 rules for filling atomic orbitals
1. AUFBAU principle
2. Pauli exclusion principle
3. Hund's rule
4. Energy Level Maximum Electron Count
38
AUFBAU principle
a) electrons fill orbitals of lowest energy first
b) orbitals within different sublevels have different energies: order of increasing sublevel energy: (lowest) s,p,d,f(highest)
c)orbitals within a given sublevel of an energy level have the same energy
d) filling order can be determined by the periodic table or by using the AUFBAU order (diagonal rule)
39
pauli exclusion principle
an orbital can hold a maximum of 2 electrons though they must have opposite spins
40
hund's principle
within a sublevel, each orbital must contain one electron before the second one can be added to the orbital
41
maximum electron count
each principal energy level can hold a max of 2n^2 electrons
42
for anions and cations how do they differ for energy level diagrams
anions- ex if 3- you add 3 electron
cations- ex if 3+ u draw it as it and u remove 3 but in a way that it stays stable
43
coordinate covalent bond
- Situation where one atom contributes both of the electrons to the shared pair
Ex. PO4 3-
44
resonance bond structures
- Common in molecules with double bonds within the structure
- Used to illustrate when experimental measurements of bond energies are not discrete single or double bond energies but intermediates(ie. Greater energy than of a single bond but less energy than that of a double bond)
- LEDD resonance structures maintain the relative position of the atoms in relation to one another but alter the location of bonding pairs.
example : SO2
45
name 2 exceptions to the octet rule
- Central atoms with expanded valence shells
- Central atoms with reduced valence shells
46
Central atoms with expanded valence shells
Elements in the second period never exceed the octet simply because the second energy level can only hold a max of 8 electrons
Elements in the third period or higher on the other hand, do have the potential to hold more than 8 electrons since d and potentially f orbitals can be occupied.
47
Central atoms with reduced valence shells
boron (a metalloid) and beryllium(a metal) often form covalent bonds and can exist in a stable union with less than 8 valence electrons
- Boron is stable with 6
- Beryllium is stable with 4
- B and Be do not form ionic bonds readily
- Due to their small atomic radius, energy required to remove the valence electrons is too high.
48
what does vsepr stand for
valence shell electron pair repulsion theory
49
what does vsepr theory provide and what is the theory about
a basis for predicting the 3D shape of a molecule. and is a theory based on the concept that valence shell electron pairs, being negatively charged, orient themselves as far apart from one another as possible.
50
bonding pair
make up a covalent bond
51
lone pair
nonbonding pair
52
order of decreasing repulsion between electron pairs
lonepair- lone pair
lone pair- bonding pair
bonding- pair bonding pair
53
in vsepr theory multiple bonds are treated as _
single bonds
54
vsepr notation: what does axe represent
A- central atom
X- number of bonding pairs around central atom
E- number of lone pairs
55
if a molecule has 2 bonding pair and 0 lone pair whats the name of the structure
linear(180degrees)
56
if a molecule has 3 bonding pairs and 0 lone pairs whats the name of the structure
trigonal planar (120degrees)
57
if a molecule has 2 bonding pairs and 1 lone pairs whats the name of the structure
bent (120 degrees )
58
if a molecule has 4 bonding pairs and 0 lone pairs whats the name of the structure
tetrahydral (109.5degrees)
59
if a molecule has 3 bonding pairs and 1 lone pairs whats the name of the structure
trigonal pyramidal (107degrees)
60
if a molecule has 2 bonding pairs and 2 lone pairs whats the name of the structure
v-shaped (104.5 degrees)
61
when determining the overall polarity of a molecule
both the shape and the polarity of the individual bonds
62
bond dipole
the charge seperation that exists between the bonded atoms and is determined by calculating the electronegativity difference
63
molecular dipole
the overall charge seperation for the entire molecule and is determined by considering both the individual bond dipoles and the 3D shape
64
does containing a polar bond mean polar
not always
65
what does fully symetrical molecules mean
non polar
66
what makes a molecule symmetrical
if all bond angles and bond dipoles are equal
67
what does valence bond theory attempt to explain
how orbitals interact to form bonds
68
non polar covalent bond
equally sharing of electrons
69
polar covalent bond
unequally sharing of electrons
70
sigma bond
single covalent bond formed by the end to end overlap of 2 atomic orbitals
71
pi bond
second bond in a double covalent bond or second and third in a triple covalent bond
formed by the side to side overlap of 2 parallel orbitals
72
what are hybrid orbitals
atomic orbitals obtained when 2 or more non-equivalent orbitals from the same atom combine in preparation for bond formation
73
does the shape of hybrid orbitals differ from the original
yes, hybrid orbitals have a different shape and energy compared to the original that have merged
74
what happens to orbitals during hybridization
orbits are neither created or destroyed but merely transformed
75
shape of hybrid orbital
teardrop
76
how to get steric number
add the bonding pair value to the lone pair value
77
what does the steric number do
identifies the hybridization of the central atoms atomc orbital in the structure
78
if a steric number of 2 what is the atoms hybrid orbital for bonding
sp
79
if a steric number of 3 what is the atoms hybrid orbital for bonding
sp^2
80
if a steric number of 4 what is the atoms hybrid orbital for bonding
sp^3
81
name 4 crystalline solids
ionic solids,
molecular solids,
metallic solids,
covalent network solids
82
examples of ionic solids
NaCl(s), CuSO4(s)
82
particle composition of ionic solids
ions
83
principle force or bond of ionic solids
ionic bond and strong directional bond
84
strong directional bond
established by the electrostatic attraction between a cation and anion
85
overall crystalline structure of ionic solids
- crystal lattice created by the attraction of an ion by the surrounding ions of opposite charge
- regular repeating pattern of ions of which the anions, being larger, generally determine the packing order
86
name 5 properties of ionic solids
hard, brittle, high mp, electrical conductivity, soluble in polar solvents
87
explain the hardness of ionic solids
network of ionic bonds (strong directional bonds) which molds ions in a fixed position
88
explain the brittleness of ionic solvents
- they have strong directional bonds and if stressed, like charged ions come too close together causing significant repulsion forces resulting in the Crystal fracturing
89
explain ionic solids high melting point
ionic bonds(strong directional bonds) require significant input of heat to break
90
explain the high conductivity of ionic solids
ions dissociate from 1 another and thus can pass electrons
91
explain the solubility of ionic solids in polar solvenets
- waters attraction to the ions can override the attraction between the cations and anions
- degree of solubility is determined by charges /size of ions
92
if something has a larger radius is it more or less soluble
more
93
if something has a higher charge is It more or less soluble
less soluble because they are attracted to other ions in compounds more strongly
94
give examples of molecular solids
I2(s), H2O(s)
95
particle composition of molecular solids
molecules
96
principal force or bond of molecular solids
- intermolecular forces
- London dispersion
- dipole-dipole force
- hydrogen bond
97
overall crystaline structure of molecular solids
molecules held together by intermolecular forces
98
name 4 properties of molecular solids
softer, lower melting point, no electrical conductivity, varying degree of solubility
99
explain the softer and lower melting point of molecular solids
- weak intermolecular forces hold molecules together
- varies depending on the type of intermolecular forces present
100
why is there no conductivity in molecular solids
electrons are locked in covalent bonds and thus are not free to move
101
explain the varying degree of solubility of molecular solids
like dissolves like
102
metallic solids examples
Pb(s), Fe(s), Cu(s), Al(s)
103
particle composition of metallic solids
cations
104
principale force or bond of metallic solids
Metallic bond
- Force of attraction between positive cations and “sea” of valence electrons that move freely through the structure
-Non directional bonds
105
Overall crystalline structure of metallic solids
3D array or metal cations held together by metallic bonds
106
107
explain the hardness of metallic solids
Electron sea surrounding cations form strong non directional bonds
108
109
explain how metallic solids are malleable and ductile
If stressed the sea of valence e- can redistribute themselves to absorb the stress and stabilize the structure
110
explain the Electrical conductivity of metallic solids
Delocalized electron are mobilized
111
explain how metallic solids are a conductor of heat
Heat is carried through metal by colliding electrons
112
explain metallic solids insolubility in solvents
Solvents are unable to disrupt the attraction of the e- to the nuclei of the atoms.
113
examples of covalent network solids
C(s), SiO2(s), SiC(s)
114
particle composition of covalent network solids
atoms
115
principal force or bond of covalent network solids
covalent bond
116
overall crystalline structure of covalent network solids
Contains a 3D arrangement of atoms held together by strong directional covalent bonds
Atoms are covalently bonded together to form one giant molecule where the combined structure is much stronger than any one unit due to an interlocking structure
117
5 properties of covalent network solids
very hard, very high melting point, brittle, no electrical conductivity, and insoluble in solvents
118
explain how covalent network solids are very hard and have very high melting points
Directional covalent bonds
Combined structure is very strong compared to any one component due to interlocking structure
119
explain how covalent network solids are brittle
Due to rigid interlocking structure
However they are so hard they rarely break
120
explain how covlaent network solids have no electrical conductivity
Electrons are locked in directional covalent bonds so thus are not free to move
121
explain why covalent network solids are insoluble in solvents
Generally insoluble since millions of covalent bonds would have to be broken
