Light/Quantum Numbers
Viable wavelengths
400-700 nm
momentum
Mass x Velocity
what does a graph of the wave function look like?
Orbital shapes
What is total energy proportional to?
amplitude and frequency of wave The larger the wave amplitude, the more force it has The more frequently the waves strike, the more total force there is
Wave behavior of electrons
de Brogile proposed that particles could have wave-like characteristics Because electrons are so small, the wave character is significant He predicted that the wavelength of a particle was inversely proportional to its momentum On eq sheet: wavelength = planck's constant / (mass * velocity)
energy of orbitals in a multi-electron atom...
depend on n and l
Color of light is determined by...
its wavelength or frequency
subshell
l (angular momentum quantum number) corresponds to a subshell
subshells
l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital
pauli exclusion principle
no two electrons in the same atom can have the same set of four quantum numbers which means An atomic orbital may describe at most two electrons, each with opposite spin direction
full d or f sub shells in main group elements are...
not considered valance electrons
full f sub shells in main group elements are...
not considered valance electrons
energy of orbitals in a single electron atom...
only depend on n
nodes
point where electron probability is 0
valance electrons
the electrons in all the subshells with the highest principal quantum number
m sub l corresponds to...
the number of different orbitals in a subshell
Energy of photon using binding energy and kinetic energy (on eq. sheet)
1 photon at the threshold frequency has just enough energy for an electron to escape the atom. This is the binding energy, BE. For higher frequencies, the electron absorbs more energy than is necessary to escape. This excess energy becomes kinetic energy of the escaped electron, KE KE = energy of photon - binding energy So, hv = KE + BE
Rules for electrons and atomic emission spectrums
1. Electrons only emit light at discrete wavelengths 2. Each specific wavelength corresponds to a specific energy. For hydrogen (On eq sheet): Energy = [(Planck's constant)*(Speed of light)] / wavelength 3. Electrons must only make transitions between discrete energy levels 4. The difference between energy levels must equal the energy of the photon emitted 5. Energy levels in an atom are quantized
Bohr's model of the atom
1. e- can only have specific (quantized) energy values, OR only orbits of certain radii, corresponding to certain specific energy levels, are permitted for the electron in an element's atom. 2. Light is emitted as e- moves from one energy level to a lower energy level Energy of energy level = -Rh (1/n^2) where n is the principal quantum number (allowed orbital number) and Rh is the Rydberg constant, 2.18 X 10^-18 J (energy needed to remove electron from hydrogen atom)
what does Schrodinger's wave equation describe?
1. energy of an electron with given wave function value 2. probability of finding an electron in a volume of space
Aufbau building up priciples
1. lower energy orbitals fill first 2. each orbital holds two electrons; each with different m sub s (spin) 3. half fill degenerate orbitals (orbitals that have the same energy) before pairing electrons
At what speed do all electromagnetic waves move?
3.00 X 10 ^8 m/s in a vacuum = the speed of light
What is white light?
A mixture of all the colors of visible light, so it is not a line spectrum. Red orange yellow green blue violet
node
A point of zero amplitude on a standing wave
Is red light (closer to infrared) or blue light (closer to ultra violet) higher energy?
Blue light
anomalous electrons configurations
Chromium should be [Ar] 4s^23d^4 but is actually [Ar] 4s^1 3d^5 Copper should be [Ar] 4s^2 3d^9 but is [Ar] 4s^1 3d^10 It makes sense that these are happening in the 4th and 9th columns This occurs because it is more stable to have half-filled and full-filled subshells Mo in column 4 and silver and gold in the 9th column follow the same pattern.
What was the mystery surrounding the classic wave theory?
Classic wave theory attributed the photoelectric effect to light energy being transferred to the electron. According to this theory, if the wavelength of light is made shorter or the light waves intensity made brighter, more electrons should be ejected. However, it was observed that high intensity lights had a maximum wavelength for electrons to be emitted (regardless of the intensity). This was called the threshold frequency. It was also observed that high frequency light with a dim source caused electron emission without any lag time.
What does wavelength determine?
Color
Toroid
Donut shape in electron orbitals d and f
Photons (explanation of photoelectric effect)
Einstein proposed: that light energy is delivered to atoms in packets called quanta or photons The energy of a photon of light is directly proportional to its frequency and inversely proportional to its wavelength The proportionality constant is called planck's constant The photon's energy must exceed a minimum threshold ( threshold frequency) for electrons to be ejected from the metal. The energy of the photon only relies on the frequency
Energy of photon using planck's constant (on eq sheet)
Energy = planck's constant * frequency = [(planck's constant) * (speed of light)] / wavelength
Energy of photon using principal quantum numbers and rydberg's constant (only works for hydrogen!)
Energy of final level = -Rh (1/n(final)^2) Energy of initial level = -Rh (1/n(initial)^2) On eq sheet: Energy of photon = (delta)E = Rh [ (1/n(initial)^2) - (1/n(final)^2) ] When the electron falls to a lower energy level, the value will be negative because n(final) is smaller than n(initial). Thus, when (delta)E is negative, a photon is being emitted. The energy of the photon is always positive, so when (delta)E is negative, the energy of photon = -(delta)E When the electron jumps to a higher energy level, (delta)E is positive because n(final) is larger than n(initial). Thus, when (delta)E is positive, a photon is absorbed.
quantum-mechanical model
Explains the manner in which electrons exist and behave in atoms; helps understand and predict atom properties as they relate to electron behavior
Relationship between frequency and wavelength
For waves traveling at the same speed, the shorter the wavelength, the more frequently they pass This means that wavelength and frequency of electromagnetic waves are inversely proportional Since the speed of light is constant, if we know wavelength we can find the frequency and vice versa
Eq. for frequency and wavelength
Frequency = speed of light/wavelength
crest
Highest point of a wave
What is amplitude a measure of?
How intense the light is, the larger the amplitude, the brighter the light
trough
Lowest point of a wave
what is light composed of?
Perpendicular oscillating waves, one for the electric field and one for the magnetic field. An electric field is where an electrically charged particle experiences force. A magnetic field is where a magnitized particle experiences a force.
variables in wave function (quantum numbers)
Principle quantum number (n): distance from nucleus Angular-momentum quantum number (l): shape of orbital Magnetic quantum number (m sub l) : orientation in space Spin quantum number (m sub s) (not really a variable): direction of electron spin
Photoelectric effect
The emission of electrons from a metal when light shines on the metal
amplitude
The height of a wave or the distance from node to crest or node to trough
Hund's rule and what it means for orbital filling diagrams
The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. This mean orbitals degenerate orbitals will fill with arrows going in the same direction first
Frequency
The number of waves that pass a point in a given period of time. Number of waves = number of cycles Units are hertz (Hz) or cycles/s = s^-1 1 Hz = 1 s^-1 Signified by nu (looks like v)
Uncertainty principle
The product of the uncertainties in both the position and speed of a particle are inversely proportional to its mass. This means the more accurately you know the position of a small particle, like an electron, the less you know about its speed and vice versa
What is observed color of objects?
When an object absorbs some of the wave lengths of white light while reflecting others, it appears colored. The observed color is predominantly the color reflected.
Wavelength
a measure of the distance covered by a wave, or the distance from one trough to the next, or one crest to the next, or the distance between alternate nodes. Signified by lambda
core notation
a method of writing an electron configuration in which core electrons are represented by a noble gas symbol in brackets followed by the valence electrons for example (NE) 3s2
Just like all matter is made of atoms...
all electromagnetic energy is made of photons
s block, p block, etc on periodic table (with number reminders for d and f)
draw and refer to lecture slides
d orbital
each each principal energy state above n = 2 has 5 d orbitals (m sub l = -2, -1, 0, 1, 2) 4 of the 5 orbitals are aligned in a different plane while the fifth is aligned with the z axis 3rd lowest energy orbitals in a principal energy state mainly 4-lobed planar nodes (entire planes that don't have electrons) 10 electrons
p orbital
each principal energy level above n = 1 has 3 p orbitals (m sub l = -1, 0, 1) each of the 3 orbitals point along a different axis 2nd lowest energy orbitals in the principal energy state two-lobed nodes at the nucleus, total of n nodes 6 electrons
S orbital
each principal energy state (n) has 1 S orbital S orbital is the lowest energy orbital in a principal energy state spherical number of nodes = (n-1) 2 electrons
f orbital
each principal energy state above n = 3 has 7 d orbitals (m sub l = -3, -2, -1, 0, 1, 2, 3) 4th lowest energy orbitals in a principal energy state mainly 8-lobed planar nodes 14 electrons
shell
each value of n (principle quantum number) is called a shell.
orbital filling diagram
electrons are arrows (arrows are different directions to represent different spin)
electrons and line spectrums
electrons in atoms of elements have certain energy levels in which they can exist. when they become excited, they jump to higher energy levels. Because they can't stay excited forever, they fall back down to certain energy levels, and the energy they lose is given off as light (or photons). The spectrum of light an element gives off (called line spectrum or electromagnetic emission spectrum) when excited is due to the energy levels that its electrons can exist in.
core electrons
electrons in energy shells lower than the valance electrons
diamagnetic
elements which have all electrons paired and are unaffected by magnetic fields
paramagnetic
elements which have unpaired electrons and are affected by magnetic fields. The more unpaired electrons, the more magnetic.
electromagnetic radiation
form of energy, travel as waves, e g radio waves, visible light, UV, gamma rays.
planck's constant
h on equation sheet
periodic table and valance electrons, quantum number, and electrons in subshells
in groups 1-8A, the number 1-8 indicates the number of valance electrons in the atom of an element in that column the row number indicates the quantum number (with d being 1 less than s and f being 2 less than s) the number of columns over you are in the s, p, d ,or f block indicate the number of electrons in that subshell.
atomic emission spectrum or...
line spectrum
quantum
little individual units
electron configuration configuration
n(letter l corresponds to)^number of electrons in orbital or subshell
Allowed quantum numbers
n: 1, 2, 3, ... l: 0,1, 2, 3 .... n-1 (number of values = n) m sub l: -l,....0.....,+l (number of values = 2l+l) m sub s: -1/2 or +1/2 (number of values = 2)
electromagnetic spectrum from least energetic to most
radio waves, microwaves, infrared, ultraviolet, x-rays, gamma rays
orbital
region in space where electron is more likely to be found
psi greek
symbol used for the wave equation
