Download Atomic Theories

Atomic Theories
CHEMISTRY REVIEW: SECTION 1.4
PG. 18-26
Atomic Theories
  Theories – are comprehensive sets of ideas based on general
principles that explain a large number of observations
 
i.e. Atomic Theory – matter is composed of atoms
  Theories are dynamic; they are continually undergoing
refinement and change
 
Geocentric Model vs. Heliocentric Model of the solar system
  Theories must:
  Describe observations
  Explain observations
  Predict future results and future experiments
  Be simple in concept and application
It is never possible to prove a theory!
  Predictions are made based on a theory...
  Then experimental evidence is collected...
  If
verified – evidence agrees with the prediction and the theory is
trusted more
  If
falsified – evidence contradicts the prediction and the theory is
unacceptable (falsified)
 
The theory must then be restricted as to where it can be used,
revised (some details are changed), or replaced by another theory
The Evolution of Atomic Theory
Ancient Greeks
Four elements
J.J. Thomson
Raisin Bun
John Dalton
Billard Ball
Neils Bohr
Bohr Model
Ernest Rutherford
Nuclear Model
1) Ancient Greeks
  One of the earliest beliefs about matter came from the Greek
Philosopher, Aristotle.
  He believed very firmly that all matter was composed of
combinations of Earth, Air, Fire and Water.
  Another lesser known Greek philosopher, Democritus,
believed that matter was made up of tiny particles that could
not be divided into smaller pieces
  Democritus' idea of matter is much closer to what we
currently know is true about matter but Aristotle was better
known and well respected at the time, so his idea was accepted
for over 2000 years.
2) John Dalton (1766-1844)
  Matter is composed of indestructible, indivisible atoms, which
are identical for one element, but different from other elements
(“Billard Ball Model”)
  This simple model is still used today to represent the
arrangement of atoms in molecules
  See Table 1 on pg. 19
3) Joseph John Thomson (1856-1940)
  The atom is a lump of positive protons with negative electrons
embedded into it (“Raisin Bun Model” / “Plum Pudding”)
 
The kind of element is characterized by the number of electrons in the atom
  The discovery of cathode rays (streams of electrons observed in vacuum tubes)
made this model possible because it showed the presence of electrons
  This was the first “divisible” model of the atom
  Raisin bun Model: An atom is a sphere with a uniformly
distributed positive charge and embedded within it are
enough electrons to neutralize the positive charge.
  See Table 2 pg. 19
4) Ernest Rutherford (1871-1937)
  Developed a more refined picture of the atom in 1911 with the
“Gold Foil Experiment”
 
In this experiment, a beam of positively charged helium nuclei (alpha particles = +2
charge) was passed through a thin sheet of gold foil
 
Most particles passed through the foil but a few particles were deflected back
 
From these observations, Rutherford concluded that:
 
The mass of an atom and its positive charge are concentrated in a
small region he called the nucleus (positive part)
 
The rest of the atom was mostly empty space where the electrons
(negative part) would be found
 
See Table 3 pg. 20
Rutherford Model
aka Nuclear Model
Alpha particles are very dense
and strongly positive particles
Expected: Because α particles
were more massive than
electrons they would not be
significantly deflected by the
electrons in the gold foil
Found:
Over 99% went straight through,
so atoms are mostly empty
space.
BUT, a few alpha particles were
highly deflected
- So there was something very
dense and very positive in the
center = NUCLEUS
- Nucleus must be small though
b/c most alpha particles went
straight through
Rutherford’s Gold Foil Experiment
5) Neils Bohr (1885- 1962)
  In 1913, he proposed that electrons are arranged in
concentric circular paths (orbits) around the
nucleus in fixed energy levels. These energy levels
have a very specific amount of energy; “quantized
energy”
  Electrons can “jump” from one energy level to
another but can only go into specific orbitals;
shifting gears in a car; there’s no 2½ gear)
(like
 
Electrons jump to a higher level when a photon of
energy is absorbed; “excited state”
 
As the electrons fall back down, and leave the excited
state, energy is re-emitted, the wavelength of which
refers to the discrete lines of the emission spectrum.
BOHR MODEL
5) Neils Bohr (1895- 1962)
  Bohr ‘s theory, for the first time, could
explain periodic law.
 
 
 
Periods = the number of energy levels
that are filled
A period comes to an end when the
maximum number of electrons are in
the valence (outermost) energy level
(max for each energy level = 2, 8, 8, 18
etc)
Note: You may recall that the last digit
in the group number = the number of
electrons in the valence energy level
See Table 6 pg. 23
Example: F (fluorine atom)
Period 2= 2 filled energy levels
Group 17 = 7 valence electrons
7 e2 e9 p+
Bohr’s energy level (orbit) is a region where the electron is likely to be
moving. But the most important property of the electron is its energy, not
its motion or location. Because of this, we will no longer use the circular
orbit when drawing energy level diagrams.
Bohr Energy Level Diagram
Energy Level Diagram
7 e−
2 e−
9 p+
F
fluorine atom
8 e−
2 e−
9 p+
Ffluorine ion
Review of Science 10 Terms
  Bohr also suggested that noble gases are unreactive because they have full outer
electron orbits. (remember: The Octet Rule)
  So what would sodium want to do:
Sodium atom
1 e−
8 e−
2 e−
11 p+
Na
Lose 1 electron
8 e−
2 e−
11 p+
Na+
Sodium ion
  Monatomic ions – single atoms that have gained or lost electrons
 
Cations – positively charged ions (Na+ = sodium ion, K+, Ca2+, Mg2+)
 
Anions – negatively charged ions (Cl-,= chloride ion, Br-,O2-, S2-)
Review of Science 10 Terms
  Atoms are composed of protons, neutrons and electrons
  Atomic number = number of protons in the nucleus of an atom
  Mass number = the number of protons plus neutrons in the nucleus
  Remember: [Mass number – Atomic number = # of neutrons]
OR
[Atomic mass (rounded to nearest whole #) – Atomic # = # of neutrons ]
  Isotopes – have a fixed number of protons but the number of neutrons may vary
(so stability and mass also vary)
mass number
  Two isotopes of carbon.
  Carbon-12 is stable with 6 protons and 6 neutrons
  Carbon -14 is radioactive with 6 protons and 8 neutrons
FYI pg. 24 DID YOU KNOW - Useful Isotopes
atomic number