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The Rutherford Model (aka the Planetary Model) was an improvement over the previous models, but it was still incomplete. It did not include the distribution of the negatively charged electrons in the atom.
We know that negative and positive particles (that is e - and p + )
attract each other, so the big question became:
Why don’t the electrons crash into the nucleus ?
If + and – charges attract, why don’t e - collapse into the nucleus?
In 1913, a student of Rutherford’s created a new model for the atom; he proposed the e - ’s were arranged in concentric circles around the nucleus (patterned after the movement of planets around the sun):
The Planetary Model
Along with this, he stated that the e - ’s have fixed energy that allows them to avoid falling into the nucleus, analogous to the rungs of a ladder. More on this later.
When an electric current is passed through a gas sample at low
pressure, the potential energy of the gas changes.
The ground state of an electron, the energy level it normally occupies, is the state of lowest energy for that electron. There is also a maximum energy that each electron can have and still be part of its atom. Beyond that energy, the electron is no longer bound to the nucleus of the atom and it is considered to be ionized. When an electron temporarily occupies an energy state greater than its ground state, it is in an excited state . An electron can become excited if it is given extra energy, such as if it absorbs a photon , or packet of light , or collides with a nearby atom or particle.
Niels Bohr , scientist extraordinaire, solved the puzzle of why different atoms give off different color spectra. He linked the atom’s electrons to photon (color spectra) emission. According to his new model, electrons can only circle the nucleus in allowed paths, or orbits .
It would be inaccurate to say that the electrons orbit the nucleus in the same way the planets orbit the sun, i.e., in a fixed and set path. The Heisenberg Uncertainty Principle states that you can know the position and velocity of an electrons at any given point, but never both at the same time. So if you were to plot the position of an electron many, many times, you would begin to build a picture of where it occupies space 90% of the time. This is called an orbital .
Orbital : the probable location of an electron around the nucleus.
As n increases, the number of different types of orbitals increases as well. At n = 1 , there is one type of orbital; at n = 2 , there are two types of orbitals; and so on. The number of orbitals at any given energy level is equal to the principal quantum number ( n ). These are known as sublevels .
3. d-orbitals : after the s and p orbitals, there is another set of orbitals
which becomes available for electrons to inhabit at higher energy
levels. At the third level, there is a set of five d orbitals (with more
complex shapes names) as well as the 3s and 3p orbitals (3px, 3py,
3pz). At the third level there are a total of nine orbitals altogether.
3d yz 3d xz 3d xy 3d z 2 3d x 2 -y 2
“Rungs of a ladder” N Energy of e- increases as you travel further away from the nucleus. e- can jump from energy levels when they gain/lose energy Quantum = amount of energy req’d to move an e- from its present energy level to the next highest; “quantum leap” Unlike a ladder, levels are not evenly spaced; closer further away thus easier to move b/t or leave.
QMM = probability of finding an e- within a certain volume surrounding the nucleus; represented by an electron cloud
The > probability of finding an e- is within these areas surrounding the nucleus ( represent where the e- is 90% of the time ). N The “fatter” the area of the e- cloud, the greater the chance of finding an e- and vice versa.
Within each energy level there are sublevels ; the # of sublevels equals the principal energy level (n)
The sublevels are also arranged from lowest to highest energy
These sublevels have orbitals within them; each orbital can hold a max of 2 e-
3 sublevels n = 3 2 sublevels n = 2 1 sublevel n = 1 # of sublevels in that level Principal energy level (n) 7 orbitals 4 th = f 5 orbitals 3 rd = d 3 orbitals 2 nd = p 1 orbital 1 st = s # of orbitals within each sublevel Sublevels (lowest highest energy)
Principal energy level (n) Energy sublevels Orbitals in sublevels
n = 1, 2, 3, 4… s, p, d, f, g … s =1; p = 3; d = 5; f = 7
(2 e-; 6 e-; 10 e-; 14 e-)
QMM describes an e - position within an e - probability cloud; e - don’t travel in fixed circular paths, therefore we cannot call them orbits. Rather, we call them atomic orbitals ( s, p, d, f, g …) SHAPES OF ATOMIC ORBITALS DICTATE PROBABILITY!!!
s orbital p orbital (x 3) d orbital (perpendicular orbital coming at you; x 5) Fig 13.4, 5 in book Low to High
The 2 nd energy level ( n = 2 ) has 2 sublevels, s and p.
N P P P P P Coming @ you Going away from you 3.) Spaces represent what? P S 2.) How many total orbitals are there? What are the max # of e- that can be held in n= 2? 1.) P orbitals stick out further therefore they have > ____?
1) Aufbau principle : Electrons enter orbitals of lowest energy first . The various sublevels of a principle energy level are always of equal energy. Furthermore, within a principle energy level the s sublevel is always the lowest-energy sublevel. Each box represents an atomic orbital.
2) Pauli exclusion principle : An atomic orbital may describe at most two electrons. For example, either one or two electrons may occupy an s orbital or p orbital. A vertical arrow represents an electron and its direction of spin ( ↑ or ↓). An orbital containing paired electrons is written as ↑↓ .
3) Hund’s Rule : When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins. For example, three electrons would occupy three orbitals of equal energy as follows: ↑ ↑ ↑ Second electrons then add to each orbital so their spins are paired with the first electrons.
According to the “wave model”, light consists of electromagnetic waves
All waves travel in a vacuum at 3.0 x 10^10 cm/s (or 3.0 x 10^8 m/s) = ? I’m smarter than he is? How’d he measure that?
Anatomy of a Wavelength origin amplitude Λ = “lambda” Frequency ( ν ) = “nu” = # of wave cycles that that pass through a point in a given time = Hertz (Hz) or s^-1 Wavelength and frequency are inversely related! Which leads us to…
Take 3 minutes only for quiz – hand in when finished.
Give the basic anatomy of a wavelength.
What do we broad term describes all forms of light? Which portion makes up the smallest portion of this “spectrum”?
How are wavelength and frequency related? Do they relate to anything else?
Have essays and homework questions ready!
Massive quiz on Monday (in lab) on all ch. 13
Remember to bring notebooks to class.
Tuesday – Print out a PT and after reading chapter 14, create a “map” of how to interpret the periodic trends