Download Sec 4-1 Chapter 4 Notes

Chapter 4 Notes
Sec 4-1
Sec. 4-1
The Rutherford model (or the planetary model) of the atom was an
improvement over the previous models, but it was still incomplete.
It didn’t explain where the e- were located in the space surrounding the
nucleus.
Before we can understand what came next, we must first
learn about light.
Electromagnetic Spectrum
Electromagnetic radiation is a form of energy that travels through
space. All electromagnetic radiation (or waves) travel at
3.0 x 108 m/s. This is called the speed of light.
In the early 1900’s, scientists conducted 2 experiments involving light
and matter. The first involved an phenomenon known as the
photoelectric effect, where metals will release e- when a light
shines on them.
Ex. Solar cells, photovoltaic cells, etc.
In 1900, a German physicist, Max Plank, suggested that the object
emits energy in small, specific amounts of quanta.
A quantum is the minimum quantity of energy that can be lost or
gained by an atom. (Quantum Mechanics is now a branch of
physics)
In 1905, Albert Einstein expanded on Plank’s theory and said that
any electromagnetic radiation exhibited properties of waves and
particles. This was called the Wave-Particle Duality. Light
now behaved as a wave and as a particle.
Each particle of light carries a quantum of energy called a photon.
Einstein explained the photoelectric effect by proposing the
electromagnetic radiation is absorbed in whole #’s of photons.
In 1913, Niels Bohr came to work with Rutherford. Rutherford
knew that there were errors in his model. Bohr then proposed the
atom’s e- was linked to the photon emissions. He included that the ecan circle the nucleus in allowed paths or orbits.
E1 is called ground state
E3
E1
E2
nucleus
When 1 photon hits an atom, it becomes excited and moves to a higher
energy state. We call this absorption.
E3
E1
E2
When an electron gets so excited that it “pops” off of the atom it
becomes free and creates electricity.
Emission is when an excited electron falls back down.
E3
E1
E2
Excited neon atoms emit light when falling back to ground state or to a
lower energy level.
Each element gives off a distinct band of light when it is excited.
Sec 4.2
Because the wave – particle theory was still being debated, Bohr’s
model of the atom wasn’t readily accepted.
Investigations into Einstein’s idea of the Wave-Particle Duality led to a
revolution of our basic idea of matter.
In 1924, Louis de Broglie showed that the e- were waves by showing
that they exhibited diffraction and interference.
In 1927, Werner Heisenberg came up with his uncertainty principle.
He said it is impossible to determine simultaneously both the
velocity and position of an e- .
This was very hard to swallow
In 1926, Erwin Schrodinger developed an equation that treated e- as
waves.
Together, the Heisenberg uncertainty principle and the Schrodinger
equation laid the foundation for the Quantum theory.
Heisenberg uncertainty principle gives us a probability of where
to find the e-.
The e- does not travel around the nucleus in neat orbital’s (like Bohr
thought). They exist in certain regions called orbital's. We still
use the planetary model to show two dimensions, but the orbital’s
are actually 3 dimensional.
Quantum #’s – Specify the properties of the atomic orbital's and the
properties of the e- in the orbital's.
4 quantum #’s – n, l, ml, ms
“n” - Principal Quantum #– indicated the main energy level
occupied by the e-.
“l” – Angular Momentum # - indicated the shape of the orbit.
There are 4 shapes – s, p, d, and f.
 “ml” – magnetic quantum # -orientation of the orbital.
Is it located in the x, y, z position, etc.
 “ms”– Spin - shows the spin of the electron
Can be +1/2 or -1/2 for clockwise and counterclockwise
(overheads)
Sec 4.3
Pauli’s Exclusion Principle – e- carry a spin, 1 cw and 1 ccw, which
create a magnetic field.
Aufbau’s Principle – e- must occupy the lowest energy level possible.
Hund’s Rule – e- fill up each orientation (x,y,z, etc) before they pair up.
 Electron Configuration -