Honors Chemistry
Chapter 5 Notes – Electrons in Atoms
(Student edition)
Chapter 5 problem set:29, 35, 37, 41, 44, 50, 52, 55, 63, 68, 71,73
Useful diagrams: 5.2, 5.5, 5.7, 5.9, 5.10, 5.12, 5.14
5.1Models of the Atom
The Development of Atomic Models
Rutherford’s atomic model could not explain the chemical properties of elements.
In 1913, Neils Bohr (Danish), stated that electrons occupy fixed paths or
without giving off energy.
Electrons far from the nucleus have energy.
This theory was deduced from and .
5.3 Physics and the Quantum Mechanical Model
Spectroscopy – the study of substances from the they give off.
– the instrument used to break light into its component .
Different levels: 1 through 7, K - Q
The Emission and Absorption of Radiation:
Electromagnetic Radiation: energy that travels in waves. Light is an example.
The electromagnetic spectrum:
1 x 10-16 m 1 x 10-6 m 1 x 104 m
The speed of light in a vacuum: m/sec ( miles/sec). It is a little slower in air, but that measurement is still accurate to 3 sig figs.
ROYGBIV: colors of the spectrum.
continuous spectrum
Red versus Violet:
Bright line spectrum (bls): frequencies of light give off by certain substances
when is added to them. Elements can appear to give off the same color light, but each will have its own .
bls is used to determine the composition of and .
bls validates Bohr’s idea that electrons jump to different energy levels and give
off different wavelengths of .
Demonstration of flame tests and gas tubes
The Nature of Light:
Sir Isaac Newton: (English - 1600’s) suggests light is .
Christian Huygens: (Dutch - same time) suggests light is waves of .
Max Plank: (German - early 1900’s) revives the particle theory. Plank studied
light given off by hot bodies. He proposed light is given off in bundles of energy which he called ( ). The amount of energy depends on the of light.
The modern theory is that light behaves both as and .
Frequency (f): the number of peaks that pass per second
measured in cycles per second
note: cycles per second = the Hertz (Hz) = 1/sec (sec-1)
Wave velocity (v): the distance a wave moves in one second
formula: v = f
=
So, red light has a wavelength and a frequency compared to violet.
Light as Energy: Planck derived a formula that expresses the energy of a photon
at any given frequency
Formula 1: E = hf E is the energy of a photon measured in joules
h is Plank’s constant (6.6 x 10-34 J/Hz)
f is frequency measured in Hertz
Formula 2: v = ff is frequency measured in Hertz
= wavelength measured in meters
v = velocity (the speed of light (c) = 3.00 x 108 m/sec)
Combine formulas 1 and 2: E = ch
5.1 Continued . . .
The Quantum Mechanical Model (AKA Charge Cloud Model ):
In the 1930’s - 1940’s more work leads to the charge cloud model which is also known as the quantum mechanical model
This model does not show the path of electrons - just the most location
5.3 Continued . . .
Quantum Mechanics:
Classical mechanics: laws of motion developed by Isaac Newton. These laws apply to the world brilliantly. These laws also apply to the motion of atoms and molecules.
Einstein: theory of relativity shows Newton’s equations don’t work when objects approach the speed of light.
Bohr: electron “jumps” don’t fit into Newtonian physics as well. It only worked
for . Wave mechanics was developed to help Bohr’s theory.
de Broglie (France) and Planck advance the theories of wave mechanics.
In 1927, Heisenberg (German) creates the Heisenberg uncertainty principle.
HUP: It is impossible to know both the and of an electron at the . When light hits a particle, when you observe it, it changes the speed and location.
Quantum mechanics (equations) allows scientists to determine the of finding particles in certain places. This leads to the charge cloud model.
Scientific Models: the following is a brief history of the model of the atom
1803 Dalton1911Rutherford1930 to present
Plank, Shrodinger, Heisenberg
1897 Thomson1913 Bohr
5.2 Electron Arrangement in Atoms
Energy levels of the wave-mechanical model:
Energy levels:principle: what shell, level, etc. n = 1,2,3...7
energy sub level - s, p, d, f
theoretical - g, h, i...
specific energy sublevels – px, py, pz, etc.
Atomic Orbitals andShapes of orbitals:
Sublevel / # of orbitals / e- per orbital / Max # of e- per sublevel / Shape of orbitalEnergy levels, sublevels, and total number of electrons per shell:
Energy Level (shell) / Sublevel(s) and Electrons / Total # of Electrons per Shell1 / s =
2 / s =
p =
3 / s =
p =
d =
4 / s =
p =
d =
f =
Another point to consider besides location of electrons is
spin – can be clockwise or counterclockwise
Pauli exclusion principle - 2 electrons in the same orbital must have opposite spins.
NIB - Quantum numbers
Quantum numbers: used to indicate the location of an electron.
Quantum # / Symbol / Identifies the . . . / ValuePrinciple / n / Energy Level / 1 to 7
Azimuthal / l / Energy Sublevel
(s,p,d,f) / s = 0, p = 1, d = 2, f = 3
Magnetic / m / Specific Sublevel / s = 0
p = +1, 0, -1
d = +2, +1, 0, -1, -2
f = +3, +2, +1, 0, -1, -2, -3
Spin / s / Direction of Spin / +½ or -½
Sample problem: What are the possible quantum numbers for the electrons of Carbon?
Level / n / l / m / s / notesSample problem: Write the possible quantum numbers for one of the d electrons of Iron.
nlm s
Sample problem: Which of the following are wrong quantum # combinations and why?
nlm sWhy?
110+1/2
33-1-1/2
21+2+1/2
Notation for electron configurations
Electron configuration: a representation of the of electrons in an atom.
Example: 2py2
2 =
p =
y =
2 stands for the number of in that orbital
example - oxygen or
Electron configurations :
Electrons enter orbitals in a set pattern. For the most part, they follow the Aufbau
Principle, the Pauli Exclusion Principle, and Hund’s Rule:
Aufbau Principle: electrons must fill energy levels before entering higher levels.
Draw diagonal diagram here:Draw energy level diagram here:
Pauli Exclusion Principle - electrons occupying the same orbital must have spin. It can also be stated as, “
Electron Spin: clockwise or counterclockwise
Hund’s rule ( better known as the ): before any second electron can be placed in a sub level, all the orbitals of that sub level mustcontain at least one electron.
Electron Configurations for higher atomic #’s:
At atomic # 19 (z = 19), a break in the pattern ensues. One would expect that the orbital to fill after 3p would be 3d, but alas, it is not.. 4s is the next level we fill as it has lower energy than 3d. Look at potassium and calcium.
Exceptions to the Aufbau Principle:
Cr is not
Other exceptions:
Students should be able to identify these elements simply based on total number of electrons
Significance of electron configurations
Valence shell electrons - outermost electrons involved with bonding
for n = 5, pattern is very complicated - no atom has more than 8 valence electrons
Noble gases - 8 valence electrons - least reactive of all elements
kernel - part of the atom exclusive of valence electrons (includes the nucleus)
example - sodium kernel =
electron configurations for elements in the excited state
all previous examples have been for ground state electrons
oxygen ( ground state) - 1s22s22p4
oxygen (excited state) - or , etc....
Other topics to be familiar with....
orbital notation, shorthand orbital notation, shorthand electron configuration, Lewis Dot atoms and ions, Bohr model drawings (atoms and ions)
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