Download Big Idea 1 LO: 12, 13 Introduction: Greek philosopher Democritus

Big Idea 1
LO: 12, 13
Introduction: Greek philosopher Democritus first suggested the existence of the atom. As he was a philosopher, not a scientist as he had no evidence to
support claims, his theories were not readily accepted. It took almost two millennia before the atom was placed on a solid foothold as a fundamental
chemical object by John Dalton (1766-1844). Although two centuries old, tenants of Dalton's atomic theory remains valid in modern chemical thought.
Notes:
Dalton’s ATOMIC THEORY OF MATTER- (based on knowledge at that time): Note that #ing is not important
1. All matter is made of atoms. These indivisible and indestructible objects are the ultimate chemical particles.
2. All the atoms of a given element are identical, in both weight and chemical properties. However, atoms of different elements have different weights and
different chemical properties.
3. Compounds are formed by the combination of different atoms in the ratio of small whole numbers.
4. A chemical reaction involves only the combination, separation, or rearrangement of atoms; atoms are neither created nor destroyed in the course of
ordinary chemical reactions.
Atomic Models: Plum Pudding
**TWO MODIFICATIONS HAVE BEEN MADE TO DALTON' S THEORY
Model
1. Atoms are divisible, subatomic particles were discovered.
e–
=
JJ Thomson
=
Cathode Ray Tube
p+
=
Rutherford
=
Gold Foil Experiment
n0
=
Chadwick
=
Scattering data from small particle bombardment
2. Isotopes were discovered.
a. Mass Spectroscopy (NMR)
Electron “Plum Pudding Model” = J.J. Thomson, English (1898-1903) = electron and its charge discovered —
found that when high voltage was applied to an evacuated tube, a “ray” he called a cathode ray
[since it emanated from the (-) electrode or cathode when YOU apply a voltage across it] was
produced.
Key Postulates:
- The ray was produced at the (-) electrode (cathode)
Experiments: Cathode Ray Tube
- Repelled by the (-) pole of an applied electric field, E
- Postulated the ray was a stream of NEGATIVE particles now called electrons
- Measured the deflection of beams of e- to determine the charge-to-mass
ratio (Basis of all types of Spectroscopy)
em
=
1.76 x 108 C/g
-e is charge on electron in coulombs (C), and m is its mass.
-Electron mass,
m=
9.1094 ×10 -28 grams
-Electron charge, e =
−1.602 × 10−19 coulombs (C)
-Thomson discovered that he could repeat this deflection and calculation using
electrodes of different metals ∴ all metals contained electrons and ALL
ATOMS contained electrons
-Furthermore, all atoms were neutral = there must be some (+) charge
within the atom and the “plum pudding” model was born. [the British call
every dessert pudding—we’d call it raisin bread where the raisins were the
electrons randomly distributed throughout the + bread]
“Oil Drop Experiment” = Robert Milken, American 1909 = electron mass
discovered – sprayed charged oil drops into a chamber. Next, he halted their
fall due to gravity by adjusting the voltage across 2 charged plates. Now the
voltage needed to halt the fall and the mass of the oil drop can be used to
calculate the charge on the oil drop which is a whole number multiple of the
electron charge. Mass of e- = 9.1094 × 10-31 kg.
Experiments: Milken’s Oil Drop
Nucleus “Nuclear Model” = Ernest Rutherford, English 1911 = nucleus with
(+) core discovered and theorized the existence of n0 — carried out
experiments to test/verify Thomson’s plum pudding model.
- Directed α particles at a thin sheet of gold foil. He thought that if
Experiments: Gold Foil Experiment
Thomson was correct, then the massive α particles would blast
through the thin foil like “cannonballs through gauze”. [He actually
had a pair of graduate students Geiger (inventor of Geiger counter) &
Marsden do the first rounds of experiments.] He expected the α
particles to pass through with minor and occasional deflections.
-The results were astounding [poor Geiger and Marsden first suffered
Rutherford’s wrath and were told to try again—this couldn’t be!].
-Most of the α particles did pass straight through, BUT many were
deflected at LARGE angles and some even REFLECTED!
-Rutherford stated that was like “shooting a howitzer at a piece of tissue
paper and having the shell reflected back”.
-He knew the plumb pudding model could not be correct!
-Those particles with large deflection angles had a “close encounter” with
the dense positive center of the atom. Those that were reflected had
a “direct hit”
-He conceived the nuclear atom; that with a dense (+) core or nucleus
-This center contains most of the mass of the atom while the remainder of
the atom is empty space!
Atomic Models: Nuclear Model
Observation
Most αparticles travel through
the foil undeflected.
Some αparticles are deflected
by small angles
Occasionally, on αparticle
travels back from the foil.
=not supported by evidence
Interpretation
The atom is mostly empty space.
The nucleus is positively charged as
is the αparticle.
The nucleus carries most of the
atom’s mass.
=supported by new evidence
“Fixed” Energy Levels/“Fixed” Shells/“Fixed” Orbits ”Planetary Model” = Niles Bohr, English 1913 = e- may only exist in certain distinct regions around the nucleus based upon the energy they
possess — came to work in the laboratory of Ernest Rutherford. Rutherford, theorized the planetary model of the atom asked Bohr to work on it
because there were some problems with the model: According to the physics of the time, Rutherford's planetary atom should have an extremely short
lifetime. Bohr thought about the problem and knew of the emission spectrum of hydrogen. He quickly realized that the two problems were connected
and after some thought came up with the Bohr planetary model of the atom.
Atomic Models: Bohr Model
Key Postulates:
-Electrons can circle the nucleus only in allowed paths or orbits
-The energy of the electron is greater when it is in orbits farther from the nucleus
-The atom achieves the ground state when atoms occupy the closest possible
positions around the nucleus
-Electromagnetic radiation is emitted when electrons move closer to the nucleus
Regarding Energy Transitions
-Energies of atoms are fixed and definite quantities
-Energy transitions occur in jumps of discrete amounts of energy
-Electrons only lose energy (by emission) when they move to a lower energy state
With these conditions Bohr was able to explain the stability of atoms as well as the emission
spectrum of hydrogen. According to Bohr's model only certain orbits were allowed
which means only certain energies are possible. These energies naturally lead to the
explanation of the hydrogen atom spectrum
E(light) = Efinal orbit – Einital orbital
ν = E(light) / h(Planck’s Constant)
Shortcoming of the Bohr Model
-Doesn't work for atoms larger than
hydrogen (more than one electron)
-Doesn't explain chemical behavior
Quantum Interpretation of the Atom / Orbitals - Bohr's model of the atom is important because it introduced the concept of the quantum in explaining
atomic properties. However, Bohr's model ultimately needed revision because it failed to explain the nature of atoms more complicated than hydrogen. It
took roughly another decade before a new more complete atomic theory was developed - the modern atomic theory – Quantum Model.
“Wave/Particle Duality of Matter” - Lous de Brogllie French (1921) = Traditional (classical) physics had assumed that particles were particles and waves
were waves and that is that. However, de Broglie suggested that particles could sometimes behave as waves and waves could sometimes behave as
particles - the wave/particle duality of nature.
Experiments: Wave Particle Duality
-He suggested a simple equation that would relate the
two: Particles have momentum (p), waves have
wavelengths (λ) and the two are related by the
equation:
λ=h/p
h = Planck's constant = 6.626 x 10-34 J•s
p = (mass) (velocity)
-This wave/particle duality of nature turned out to be a
key to modern atomic theory as it demonstrates that
wavelength of a particle is miniscule compared to the
size of the particle.
Heisenberg’s Uncertainty Principle
“Heisenberg’s Uncertainty Principle” - Werner Heisenberg German (1923) = Walt’s nickname from Breaking
Bad = cannot know both the relative energy and relative position of an e – — Classical physics had always
assumed that precise location and velocity of objects was always possible. Heisenberg, however
discovered that this was not necessarily the case at the atomic level. In particular, he stated that the act of
observation interfered with the location and velocity (we define more broadly as energy) of small particles
such as electrons. This is the case because observation requires light and light has momentum. When light
bounces off an electron momentum exchange can occur between light and the electron which means the
electrons location and velocity have been altered by the act of measurement. This idea is pictured right.
“Schrodinger Equation” – Erwin Schrodinger German (1926) = equation describes the “orbitals” (regions of
space) an e– occupies — took the ideas developed by de Broglie, Heisenberg and others and put them
together in a single equation that is named after him. Solving this equation can in principle predict the
properties and reactivates of all atoms and molecules:
Planck’s Constant
Called the “del-squared operator”, this
quantity describes how the wavefunction,
Ψ, changes from one place to another
A
mathematical quantity called an
“imaginary number”. It is equal to the
square root minus one.
Describes how the wavefunction, Ψ,
changes its shape with time.
Mass of particle being described
Describes the forces acting on the particle
Unfortunately, it is extremely difficult to solve, hence the development of quantum computing, for any but the
most simple atoms and molecules.
“Quantum Mechanical Model” – Collaboration of Scientists = the orbitals and energies of them are what determine the properties of atoms and
molecules — though too difficult to solve for any but the simplest atom/molecules, we can nevertheless extract some essential conclusions from it
that form the basis of the Quantum Mechanical Model.
-Energies are quantized: that is, atoms and molecules cannot have any energy. Rather they are limited to certain energies that can be predicted
by use of the Schrodinger equation.
-Energies associated with each orbital determine where e– are located: Each orbital is determined by a quantum number call the angular
momentum quantum number "l". This quantum number can take on the values l=0 (s-orbital), l=1 (p-orbital), l=2 (d-orbital), l=3 (f-orbital) etc.
Note: Quantum Numbers are not part of the content for this course.
-The result is the use of electron configurations to describe and predict these properties. Hence the following chart also found in your ECN Notes
(refer to them)