Week 8 Flashcards
(124 cards)
1
Q
constructive interference
A
occurs in regions where peaks or troughs for the 2 waves coincide
2
Q
destructive interference
A
occurs in regions where the peak of one wave coincides with trough
3
Q
amplitudes of interfering waves add together and produce a…
A
resultant wave
4
Q
diffraction
A
wen a wave encounters an obstacle, the wave appears to bend around a small obstacle or spread out in semicircles
5
Q
light shows properties of ____ and ____
A
particles- through photoelectric effect
waves- through light diffractions and interferences
6
Q
energy levels of the hydrogen atom are given by equation
A
En= -2.179*10^-18(J/n^2)
7
Q
De Broglie Wavelength (predicted that a particle with mass m and velocity v should also exhibit the behavior of a wave given by..)
A
wavelength=h/mv
8
Q
wavefunction
A
mathematical description of an atomic orbital that describes the shape of the orbital, it can used to calculate the probability of finding the electron at any given location in the orbital as well as dynamical variables such as the energy and the angular momentum
9
Q
interpretation of the wavefuntion
A
electrons are still particles, and so the wave represented by wavefunction variable aren’t physical waves, when you square them you obtain probability density which describes probability of the quantum particle being present near a certain location in space
10
Q
wavefunction can be used to determine
A
the distribution of the electrons density with respect to the nucleus in an atom (but cannot be used to pinpoint exact location of the electron at any given time)
11
Q
electrons can exists…
A
only on discrete energy levels but not in-between them, meaning the energy of an electron in an atom in quantized
12
Q
principle quantum number variable
A
energy levels are labeled with an n value where n= 1 to infinity (energy levels of an atoms are greater with the greater value of n)
13
Q
principle quantum number
A
defines the location of the energy level, similar concept as n in Bohr’s model of shell number the further from the nucleus the higher the shell number the higher the energy level
14
Q
deltaE= Efinal-Einitial= -2.179*10^-18(1/nf^2 - 1/ni^2)
A
deltaE= Efinal-Einitial= -2.179*10^-18(1/nf^2 - 1/ni^2)
15
Q
atomic orbital
A
a general region in an atom within which an electron is most probable to reside
16
Q
angular momentum quantum number
A
(l) integer defines the shape of the orbital, l=0,1,2.. n-1 (if n=1 only one value of l=0; if n=2 whereas l=1, 0)
17
Q
orbitals with the same value of l
A
will form a subshell
18
Q
angular momentum is
A
a vector, electrons with this can have this momentum oriented in different directions
19
Q
magnetic quantum number
A
(ml), specifies the z component of the angular momentum, the orbital orientation (-l to l, if l=1 ml= -1,0,1)
20
Q
orbital abbreviations
A
l=0 s orbital
l=1 p orbital
l=2 d orbital
l=3 f orbital
21
Q
orbitals
A
mathematically derived regions of space with different probabilities of electrons in them
22
Q
interpretation of wavefunctions (orbitals)
A
probability density of finding an electron at a given point in space [wavefunction(r)]^2
23
Q
dot density diagram
A
higher density of black dots, higher probability of finding electrons
24
Q
radical probability
A
probability of finding a 1s electron at a distance r from the radius
-calc by adding probability of an electron being at all points ona series of x spherical shells of radius r1, r2, r3, rx-1, rx
25
electron probability density greatest at
r=0
26
surface area of each spherical is equal to
4(pi)r^2
27
boundary surface plot
spherical shaped plot constructed by drawing a circle or sphere around a large percentage (75-90ish) of the dots
28
s orbitals in n=2, 3... however look a little different
-the electron probability density doesn't fall off smoothly with increasing r
-series of minima and maxima are observed (corresponds to radical nodes)
29
nodes
points with 0 amplitude
-number of radical nodes in an orbital is n-l-1
30
3 things happen to s orbitals as n increases
1. they become larger extending farther from the nucleus
2. they contain more nodes, similar to standing wave that has regions of significant amplitude separated by nodes
3. for given atom, s orbitals also become higher in energy as n increases because of the increased distance from the nucleus
31
p orbital
-3D model with x, y and axis
-because 2p subshell has l=1, with ml(-1,0,1) there are 3 2p orbitals
32
only ___ orbitals are symmetrical
s
33
as l increases, the number of orbitals in given subshell
increases and shapes become more complex
34
d orbitals
subshells with l=2 have 5d orbitals, the first shell to have a d subshell is n=3
35
f orbitals
principle shells with n=4 can have subshells with l=3 ad ml values of -3,-2,-1,0,1,2,3; subshells consist of 7 f orbitals
36
orbital energies
depend only on principle quantum number(n)- energies of 2s, 2p, and 2d are all equal
37
spin quantum number
(ms) complete quantum phenomenon with no reaction to other quantum numbers and can't be derived from solving Schrodinger's equations, describes an intrinsic electron "rotation" or "spinning" [alpha state- spin up, beta state- spin down] ms value = +-1/2
38
magnet has lower energy if its magnetic momentum is aligned witht he external magnet field and higher is opposite to applied field
this is why ms +1/2 has slightly lower energy in an external field in positive z direction than ms+-1/2
39
Pauli Exclusion Principle
no 2 electrons in the same atom can have exactly the same set of all 4 quantum numbers (can share n,l,ml only if ms have different values and because ms can only be +-1/2 any atomic orbital can be populated by only 0, 1 or 2 electrons
40
black body radiation
radiation(light) given off by hot bodies, the intensity is a function of wavelength, distribution is a function of the temperature
41
quantized
positions that are fixed and can't be in between 2 points
42
continuous
positions that can be anywhere on a spectrum
43
is energy absorbed or released when an electron goes from n=4 energy level to n=2 energy level
released, down arrow=energy releasing (up arrow is absorbing energy)
44
rydberg equation energy of a photon
2.179(or J given from atom being used)*10^-18((1/nf^2)-(1/ni^2))
45
what is the color of this photon equation
E=hc/wavelength
45
should electromagnetic radiation be described as a wave or a particle
evidence that its a wave: undergoes diffraction
evidence that its a particle: photoelectric effect, blackbody radiation
46
in a spectrum summary (chart with lines pointing up and down showing electrons going from one state to another)
the smallest energy change is the longest wavelength
46
wave-particle duality equation
wavelength= plancks constant/ mass * velocity
46
electrons are described by wave functions... Schrodinger's equation allows for the calculation of
-the energy levels available for the electrons in an atom
-the probability of finding an electron at a particular place in an atom
results of the equation
-electrons are found in atomic orbitals
-the exact location of an electron can never be pinpointed, but the probability of an electron being in a particular location can be calculated
47
a group of atomic orbitals in a subshell constitutes a
subshell in a shell
48
how many atomic orbitals constitute the n=3 shell
9, n^2
48
a group of subshells constitutes a
shell or principle of energy level
49
shell
principle quantum number n+=1, 2, 3; describes orbital size and energy
50
subshell
Azimuthal Quantum Number l=0, 1, 2...(n-1); s p d ; orbital shape
51
orbitals
magnetic quantum number ml= -1, 0, 1; orbital direction
52
what are quantum numbers associated with the 1s subshell
n=1, l=0, ml=0
53
Application of Paulis Expulsion Principle and the number of electrons allowed in specified orbitals; how many electrons can there be in the 1s orbitals and what are the corresponding 4 quantum numbers for both of these electrons?
2 electrons in 1s orbital
1,0,0,1/2
1,0,0,-1/2
54
orbital energies
differing from orbital diagram in previous sections cause now we have more than one electron
55
repulsion between electrons make the energies of subshells with different values of l differ
the energies of the orbitals increase within a shell in the order s
56
electrons in successive atoms on the periodic table
tend to fill low energy orbitals first (this is why 4s, 3d, 4p)
57
as n increases the size of the orbital increase sand electrons spend more time farther from the nucleus therefore
attention to nucleus is weaker and energy is associated with orbital is higher (less stabilized)
58
electron configuration
arrangement of electrons in the orbitals of an atom describe with symbol containing 3 pieces of info
-# of principle quantum numbers
-letter that designates orbital type
-superscript number that designates the number of electrons in particular subshell
59
aufbau principle
-to determine the electron configuration for any particular atom we can build the structures in the order of atomic numbers
-beginning with hydrogen and continuing across the periods er add one proton at a time to the nucleus and one electron to the proper subshell until we have described the electron configurations of all the elements
60
writing orbital diagrams
pictorial representations of the electron configuration, showing individual orbitals and pairing arrangement of electrons
61
noble gas configuration
abbreviation that consists of the elemental symbol of the last noble gas prior to that atom followed by configuration of remaining electrons
62
exceptions for electron configuration
-subshells with similar energy and small effects can lead to changes in order of filling
-in cases of Cr and Cu we find half-filled and completely filled subshells represent conditions of preferer stability
63
ionic compounds
complete transfer of electrons
64
covalent compounds
sharing of electrons
65
valence electrons
electrons don't participate in chemical behavior
66
core electrons
electrons don't participate in chemical behavior
67
for main group elements
number of valence electrons is given by group number
68
coulombs law
says the force between 2 charged particles is dependent on the distance between the charges; the further the electron from the nucleus the less pull itll feel from the positively charged protons
69
potential U
constant * (Q1Q2)/d Q1charge on proton Q2 charge on electron
70
Z
number of protons in the nucleus
71
Z effective
estimate of the effective nuclear charge felt by an electron in an atom
72
EX hydrogen atom
the outer most electron is in the 1s shell, experiments confirm that both Z and Zeff are H
73
due to shielding the electrons closest to the nucleus
decreases the amount of nuclear charge affecting the other electrons caused by the partial screening of nuclear charge by core electrons
74
equation to estimate the effective nuclear charge for any given charge for any given atom or ion
Zeff= Z-S (protons(atomic number) -core electrons
75
the total energy of the atom or ion, when more than one electron is present, depends on the
not only on the attractive electron nucleus interactions but also on repulsive electron-electron interactions
76
the Zeff experienced by an electron in a given orbital
depends on the spatial distribution of the electron in that orbital and on the distribution of all other electrons present
77
shielding is determined by
the probability of another electron being between the electron of interest and the nucleus
78
Zeff for valence electrons increases
as we move from left to right across a period
79
atomic radius
the radius of a sphere that represents an atom (1/2 the distance between the nuclei of 2 identical atoms when joined by covalent)
80
atomic radii is determined by
measuring bond lengths between 2 bonded atoms
81
Cl-Cl bond length is 1.99 so the atomic radius is
0.99
82
ionization energy
the energy needed to remove an electron from an atom or cation in the gas phase; process can be written as a thermochemical equation(Na(g)->Na2+(g)+e-
83
first ionization energy equation
represents the first electron removed form a atom
84
second ionization energy
removing a second electron from atom
85
electron affinity
the energy change when an electron is added to a gaseous atom to form an ion
86
another definition of ionization energy
the energy needed to move a ground-state electron to the n=infinity energy level
87
valence electrons play biggest role in
chemical equations cause they're in outer most shell, have the highest energy
88
elements in the same group(column)
have same number of valence electrons
89
similarity in chemical properties among elements of the same group occurs because
same number of valence electrons, its the loss, gain or sharing that defines how they react
90
main group elements
last electron added enters s or p orbital in outermost shell
91
transition elements or metals
metallic elements in which the last electron added occupies a d orbital; valence electrons in these elements include the ns and (n-1) d electrons
92
inner transition elements
metallic elements in which the last electron added occupies an f orbital; valence shells consist of the (n-2) f and (n-1)d and ns subshell; two inner transition series: 1. lanthanide series: lanthanum (La) through lutetium(Lu) 2. actihide series: actinium (Ac) through lawrencium (Lr)
93
variation in atomic radius
as we scan down a group n increases by 1 for each element; this shows electrons being added to a region of space that is increasingly distant from the nucleus, the size of atoms increase as you increase the distance of the outermost electrons from the nucleus to the
94
variation in ionization energies
energy is always required to remove electrons from atoms or ions
95
as Zeff for the valence electrons increase from left to right the first ionization energy increases too
because more energy will be required to remove electrons
96
as n level for the valence electrons increases down a group the atoms get
bigger and the coulombic attraction between nucleus and valence electrons decreases; valence electrons become easier to remove and first ionization energy decreases
97
ion
an atom or molecule that has gained or lost one of its valence electrons resulting in a positive or negative net charge
98
cation
a positively charged ion that has more protons than electrons (when electrons are removed from parent function)
99
anion
a negatively charged ion that has more electrons than protons
100
(hw11) the magnitude of the enthalpy of vaporization is a direct indication
of the intermolecular forces of the a liquid
101
electron affinity
the amount of energy released when an electron is added to a neutral form of an atom to form an anion
102
probability density distributions contain
90% probability of the electrons being in that region
103
4(wavefunciton)r^2
the probability of finding the electron at all points in space at distance r
104
in a radial distribution graph function most electrons
are found close to the nucleus and the graph is highest right away, farthest left
105
the highest point of the graph is
where it is most probable to find an electron from the nucleus
106
how to determine the number of spherical nodes for a given orbital
n-L-l
107
number of nodal planes
L
108
electron spin is quantized
ms= 1/2 or -1/2
109
Pauli Exclusion Principle
no two electrons in an atom can have the same set of four quantum numbers
110
paramagnetic
contains 1 or more unpaired electron when looking at the electron configuration boxes
111
diamagnetic
all electrons are paired in orbitals
112
determining Zeff
Zeff=Z(atomic number)-S(core electrons)
113
transition metal electron configuration
electrons in s orbital are easier to remove than from d or f electrons, highest ns electrons are lost first then (n-1)d electrons are removed
114
main group metals
the electrons that were added last are first electrons removed; cation is formed when electrons are removed from parent function
115
main group nonmetals
added electrons fill in order predicted by Aufbau principle; anions form when one or more electrons are added to parent atom
116
ferro magnetism
occurs when certain metals act as permanent magnets, due to presence of unpaired electrons
117
paramagnetism
occurs when an atom or ion is attracted to a magnetic field, due to unpaired electrons
118
diamagnetism
occurs when an atom or ion isn't attracted to a magnetic field; due to all electrons being paired
119
atomic radii of cations
cations have fewer electrons and same number of protons than parent atom so it has a smaller atomic radii than the parent atom
120
