Exam 2 Flashcards
1
Q
Electromagnetic radiation
A
(EMR) the set of waves of an electromagnetic field that propagates through space and carries momentum and electromagnetic radiant energy
2
Q
Examples of Electromagnetic radiation
A
Radio waves, microwaves, infrared, visible, ultraviolet, x-rays, gamma rays
3
Q
characterization of electromagnetic radiation
A
-wavelength (how long a wave is)
-Frequency (how often a wave repeats)
4
Q
what is different about radio waves and x-rays?
A
wavelength and frequency -> different energies because of these differing values
5
Q
Wave-particle duality
A
the usage of both the wave and particle models to describe the behavior of light
6
Q
frequency of wave units
A
S^-1 or Hz 1S^-1=1Hz
7
Q
relationship between wavelength and frequency
A
lower wavelength = lower frequency
8
Q
Amplitude
A
height of peaks
9
Q
Wavelength
A
weird A (m) : distance between two identical points
10
Q
frequency
A
v (Hz or S^-1) : number of waves per sec
11
Q
c=a*v
A
the speed of light related wavelength to frequency
12
Q
c
A
speed of light, constant, 3.00x10^8 m S^-1
13
Q
energy of light
A
increases as frequency increases (and wavelength decreases)
-Not related to amplitude!
14
Q
Electromagnetic spectrum trends
A
left side (gamma rays) have shorter wavelength, higher frequency and higher energy
right side (Long radio waves) have longer wavelengths, lower frequency and lower energy
15
Q
to calculate a frequency given wavelength
A
c=A*v v=c/A
16
Q
converting meters ot nanometers
A
1nm = 10^-9 m
17
Q
properties of waves
A
diffraction and interference
18
Q
diffraction
A
when a wave crest encounters a barrier with a slit, the wave bends (diffreacts) to fit the entire wave through the slit
-particles will only pass through the exact opening in the slit, leaving some particles on the other side of the barrier
19
Q
interference
A
constructive interference and destructive interference
20
Q
constructive interference
A
when waves are in phase with one another ( their crests and troughs match perfectly- they are at the same point in their cycle simultaneously) the crests and troughs reinforce, doubling the amplitude **but wavelength stays the same
21
Q
Interference patterns
A
when a wave is diffracted first by a boundary with one slit, then by a boundary with two slits, there is a pattern of alternating constructive and destructive interference
21
Q
Destructive interference
A
when waves are out of phase with one another (different points in cycle) the crests and troughs cancel one another
22
Q
evidence that light is a wave
A
light diffracts and interferes like a wave
23
Q
what is the problem with the claim that light is a wave
A
wave nature of light does not explain some phenomena of light- the energy of light wave does not depend on the amplitude of the wave
24
Amplitude and energy of waves
With matter, we know the amplitude of the wave is related to the energy of that wave- the energy of a light wave does not depend on the amplitude of the wave
25
evidence that light is a particle
photoelectric effect
26
Photoelectric effect
many metals emit electrons when the electromagnetic radiation shies on the surface
27
how are these metal electrons emitted?
-light transfers energy to the electrons at the metal surface
-energy is transformed into kinetic energy that give the electrons enough energy to "leave" the atom in the metal
28
why does adding energy in the form of light allow electrons to "escape"
breaking interactions requires energy
29
Data vs Wave Model
Experimental Observations:
* If the intensity (amplitude) of
light is increased, more
electrons are ejected.
* If the frequency of light is
increased, the speed of the
electrons increases.
* No electrons are emitted
below a certain frequency of
light (threshold frequency)
regardless of intensity.
Wave Model Predictions:
* If intensity increases, then more
energy is available, causing
more electrons to be ejected.
* If frequency increases, then
more energy is available,
increasing speed of electrons.
* If the intensity of light is high
enough, electrons should be
emitted at any frequency
30
amplitude
determines the amount of electrons emitted
31
increased frequency
more energy is available, increasing speed of electrons, kinetic energy
32
threshold frequency
the certain frequency needed for electrons to be emitted, regardless of intensity (if this is not met, no electrons are emitted) (this goes against the wave model that if the intensity of light is high enough, electrons should be emitted at any frequency)
33
the difference between the wave model and the photoelectric effect
The wave model held that a high amplitude outweighed frequency, but the photoelectric effect holds that a certain frequency needed to be met regardless of amplitude
34
What part of the light model was revised
light as a particle instead of a wave
35
Einstein's idea
light must come in packets of energy (or particles or quanta) calles photons
36
photons
packets of energy
-Each photon has a certain energy
- The energy of a photon is quantized (can only have certain values)
-energy of a photon only depends on (is representative of) frequency of light, not intensity/amplitude
37
photon equation
E(J)=h(Js)v (Hz or S^-1)
38
h in the photon equation
planck's constant (6.626 x 10^-34 J s)
39
Data vs Particle Model
Experimental Observations:
* If the intensity (amplitude) of
light is increased, more
electrons are ejected.
* If the frequency of light is
increased, the speed of the
electrons increases linearly.
* No electrons are emitted
below a certain frequency of
light (threshold frequency)
regardless of intensity.
Particle Model Predictions:
* If intensity increases, more
photons are emitted, resulting
in more electrons ejected.
* If frequency increases, each
photon has more energy that
is transferred to the electron
as kinetic energy, increasing
speed.
* Only photons with enough
energy (high enough
frequency) can eject electrons.
40
necessary equations for light
E=hv c=Av E=(hc)/A
41
where is wave-particle duality most important
very small scale- microscopic of microscopic
42
why doesn't matter and energy behave like they do in our macroscopic world as the do in the microscopic world
light travels like a wave and interacts with matter like a particle
43
Visible spectrum
visible light is only a very small part of the full electromagnetic spectrum
-comes in different wavelengths and frequencies
-energy of visible light photons depends of frequencies
-visible light from the sun (white light) can be separated by a prism
44
Atomic Emission spectra
-atoms can emit light
-one particular element does not emit all the colors of the spectrum- only a few wavelengths
-each element has its own unique emission spectrum
45
Atomic Absorption spectra
-atoms can absorb light
-one particular element does not absorb all the colors of the spectrum-only a few wavelengths
-each element has its own unique spectrum absorption spectrum
46
how are absorption spectrums found
atomic absorption spectrometer
- a hollow cathode that emits white light through a lense followed by an atomized sample and then another lense which flows into a monochromator and a detector then an amplifier and a readout
47
spectroscopy
branch of science that investigates and measures spectra
48
how to create a continuous spectrum
sine white light through a prism- a rainbow line is emitted
49
how to create an emission spectrum
shime hot gas of an atom through a prism- separate lines of color are displayed on a background of black
50
how to create an absorption spectrum
shine white light through cold gas of atoms, a rainbow line with cutouts of what is not absorbed will be displayed
51
relation between lines displayed on emission and absorption spectrums
what is emitted is void in the absorption spectrum, what is absorbed if void in the emission spectrum (note; elements are known for emitting specific light, so they absorb a greater variety than what they emit)
52
does the spectrum of an element differ if not on earth?
no, each element has a unique singular spectrum weather on earth, the sun, or in a galaxy light years away
53
does the Rutherford model of the element explain atomic absorption/emission and its stagnence?
no
54
what is the issue of the Rutherford model of the atom
-electrons would be attracted to nucleus and eventually collide, causing the atom to implode
-as the electrons fall into the nucleus, energy would be emitted CONTINUOUSLY
-- we would observe a continuous emission spectrum
55
Bohr Model of the atom
-electrons move in orbits around the nucleus with definite energies and a definite distance from the nucleus
--energies of electrons in atoms are quantized
-Explained emission and absorption spectra by invoking discrete energy levels- characterized by quantum numbers
--photons of electromagnetic energy are emitted or absorbed by atoms as electrons move from one energy level to another
--the energy of the photons corresponds to the difference in energy between the orbits
56
quantized
a system can not have any possible energy values, but is rather limited to a certain specific energy value
57
electrons of an orbital
have the same energy / need the same amount of energy to be removed
58
energy relates to
wavelength and frequency
59
Atomic excitation
An electron moves to a higher energy orbit when a photon is ABSORBED (a photon of the exact wavelength/frequency representative of the amount of energy needed for that electron to move shells)
60
Atomic emission
A photon is emitted when an electron moves to a lower energy orbit (the photon emitted is of the exact amount of energy released when the electron moves and corresponds to a specific wavelength/frequency)
61
Energy diagrams
-each energy level has a quantum number (n)
-higher number = higher energy
-energy levels are NOT ORBITS
--represent energy only, NOT distance from the electron to the nucleus
-Electrons transition between energy levels by absorbing or emitting photons with energies equal to the exact difference in energy between the two levels
62
Bohr's model works for
only the Hydrogen model- describes energy of electrons and distance from the nucleus
63
Energy diagrams work for
all elements
-the energies of the electrons can be represented by the diagram but, we can not determine precisely where they are
64
why can we not determine exactly where the electrons are?
electrons are waves (and particles)
65
evidence for wave-like properties of matter
-electrons shot through aluminum foil also exhibit a diffraction pattern just like x-rays (Davisson and Germer Experiment)
-- x-rays and electrons have similar wavelengths
-When electrons are used in the double slit experiment, they show an interference pattern
--constructive and destructive interference are occuring. Diffraction is happening to matter
66
Wave-Particle Duality
Large object:
* large momentum (mv)
* small wavelength (λ)
* definite position
Small object:
* small momentum (mv)
* large wavelength (λ)
* uncertain position
67
wave-particle duality equation
A (m) = h/m(kg)v(m/s)
68
electrons in an atom are treated like a
wave
69
can we ever know the exact position of the electron and its speed?
no, when we try to observe the electron, we change its position and speed
70
how do we determine the position of an electron about the nucleus?
using its probability, which arises from quantum mechanics
71
orbitals
describe where an electron with a given energy is likely to be found
-can be visualized as shapes
-can be represented with quantum numbers
72
quantum numbers
principle quantum number, angular momentum quantum number, magnetic quantum number (and the fourth)
73
Principal quantum number
n
gives size/ energy (relates to energy diagram)
74
Angular momentum quantum number
l
denotes shape (e.g. sphere, dumbell)
75
Magnetic quantum number
m sub l
denotes orientation (e.g. on e-axis, on y-axis, on z-axis
76
Orbial shapes
s, p, d
77
s orbitals
-spheres (l = 0)
-have one orientation (m sub l = 0)
-orbital size increases as n increases (1s < 2s)
--e- sends more time further away from nucleus
78
Boundary surface plots
define the region where 90% probability of finding the electron
-human constraint- not a real boundry
79
p orbitals
- dumbbell shaped (l = 1)
-3 orientations (m sub l = -1, 0, +1)
-- px, py, pz
-orbital size increases as n increases (2p < 3p)
- p orbitals have one node
80
node
regions where the probability of finding an electron os 0%
-occurs where orbital changes phase (often represented by change in color)
-the more nodes an orbital has, the higher the energy of the orbital
81
Nodes on a wave model
- electrons have wavelike properties and can be treated as standing waves
-node occurs where the standing wave has an amplitude of 0
82
d orbitals
- l = 2
- have 5 orientations (m sub l = -2, -1, 0, 1, 2)
--dxy, dxz, dyz, dx^2-y^2, dz^2
- orbital size increases as n increases (3d < 4d)
-d orbitals have two nodes/planer nodes
83
orbitals can hold...
2 electrons
each orbital can "hold" up to 2 electrons
84
Radial nodes
the space between two shells
85
Quantum number possibilities
n = 1, 2, 3, 4, ...
l = 0, 1, 2, 3 ... (s=0, p=1, d=2)
m sub l = 0, +/- l (p -> m sub l = -1, 0, 1 ; d -> m sub l = -2, -1, 0, 1, 2
m sub s = +1/2, -1/2
86
spin magnetic quantum number
electron spin (e.e., up or down)
87
pauli Exclusion Principle
no two electrons in an atom can have the same set of four quantum numbers
-ex. no two up arrows in the same box
88
Box diagrams
visual display of elements based on their electron
89
Aufbau Principle
orbitals fill with e- from lowest energy to highest energy
90
ground state
lowest energy state
91
violation of the Aufbau principle
can mean it is in an excited state i.e one up arrow in 2s and one up arrow in 2p
92
Hund's Rule
maximize the number of unpaired electrons with parallel spins in orbitals of the same energy
-this forces electrons to occupy different orbitals and minimizes e- to e- repulsion
93
core electrons
low in energy and are closest to the nucleus- within the closed shell (n = 1, 2, 3,...)
-these don't participate in reactions
94
valence electrons
those that are higher in energy- outside the closed shell
-electrons that determine reactivity
95
electron configuration breakdown
ex 1s^2 1s is the orbital description, ^2 is the number of electrons in that orbital
96
condensed electron configuration description
noble gas in previous row in brackets ex [He]
97
what is the evidence for that electrons are in orbitals with quantized energy, regarding periodic trends?
Atomic Radius and Effecting Nuclear Charge
98
Atomic radius
half the distance between two nuclei of atoms of the same element when their interaction is at its most stable point
-depends on weather you are measuring the covalent interaction (based on data) ot the Van der Waals interaction
-- covalent interaction is within the same electron cloud, van der waals is not
99
Atomic radius trends down a group
outermost electron is in increasingly larger orbitals further from nucleus -> radius increases
100
Atomic radius trends across a row
radius decreases (a jump is displayed from row to row)
101
what determines the size of an atom
the balance between:
-attractive force between protons and electrons; the pull of electrons closer to nucleus decreases radius
-repulsive force between electrons; the push of electrons by other electrons increases radius
--takes us back to coulomb's law F=q1q2 / r^2
102
Effective nuclear charge interactions
- electrons in same or larger shell (other valence electrons) in relation to the electron of interest (which is a valence electron) have no effect on the charge experienced by the electron of interest
- electrons between nucleus and electron of interest (core electrons) shield some of the nuclear charge
**Every valence electron is attracted by the "effective nuclear charge"
103
Effective Nuclear charge
- core electrons 'cancel out' the positive charge from the same number of protons
- each electron in the valence shell feels the effect of the protons that are left
104
Effective nuclear charge equation
Zeff = Z - S
Net attractive force = total attractive force - total repulsive force
Z=#protons S=#core electrons
105
What happens to the attractive force as the effective nuclear charge increase?
it increases
106
Summary of Atomic Radii trends
Trend #1: Atoms radius increases down a column of the
periodic table.
* Orbitals at higher energy levels are larger and further from
the nucleus → reduces electron-electron repulsion
* Greater distance between nucleus and electrons results in a
weaker electrostatic force <- important for ionization trends
Trend #2: Atoms radius decreases across a row of the periodic
table.
* Number of protons increases but number of core electrons
stays the same → effective nuclear charge increases
* Electrostatic force between nucleus and electrons increases
* Electrons are pulled closer to the nucleus
107
The atomic radius represents
the stat where the forces of attraction between electrons and protons are equal to the forces of repulsion between electrons
108
when electrons have been added or removed
ions
109
cation
positively charged ion
110
anion
negatively charged ion
111
removal of an electron = what charge
positive
112
addition of electron = what charge
negative
113
size of neutral atom vs anion
anions have more electrons meaning a greater electron-electron repulsion and a larger radius
114
comparing sizes of cations and anions
calculate the effective nuclear charge (Zeff = Z-S) protons-core electrons
115
Isoelectronic series
When various cations, anions, and atoms have the same electron configuration but a different number of protons
ex O2- F- Ne Na+ Mg2+
116
Would you expect an atom where the valence electron experiences a large effective nuclear charge to have a large or small radius, with electrons easy or difficult to remove
small radius with electron difficult to remove
117
Ionization energy
energy required to remove an electron from an atom in the gas phase
118
why is energy required to remove an electron?
need energy to break the attractive interaction between the nucleus and the electrons
119
Why does Ionization energy decrease down a group
The outermost electron is further from the nucleus, yielding a weaker attractive force because of less energy to overcome
120
What happens to the force between each valence electron and the nucleus? (aka why does IE decrease down a group
Neither the charge of the electron nor the effective nuclear charge change, but the radius increases, decreasing the force (Coulomb's law)
121
why does ionization energy increase across a row
Effective nuclear charge increases, meaning a larger attractive force and more energy needed to overcome force
122
Exceptions to Ionization trends
B and O
123
why is Boron an exception to IE row trends
it is the first of the row to have an electron in the P orbital, meaning it is further from the nucleus and positive charge, and easier to remove
124
Why is Oxygen an exception to IE row trends
because removing the only opposite spinning electron is easier than removing electrons with the same spin - it reduces electron-electron repulsion
125
sucessie ionization
the removal of an electron after removing an electron and so on
126
when does ionization energy become the largest
when it is for removing core electrons
127
Factors affecting ionization energies
* Effective nuclear charge (larger Zeff → larger IE)
* Size of atom/ion (smaller size → higher IE)
* The shell that the electron is removed from (IE of core e–
much larger than IE of valence e–)
128
do successive ionization energies increase or decrease
increase
