Transcript History of the Atomic Model

History of the Atomic Model
Early Greek Theories
Aristotle (350 B.C.)
• 4 Elements
Democritus (400 B.C)
• Atoms and a void (empty space)
• Atoms are indivisible
• Based on philosophical arguments,
NOT experimental evidence
Dalton’s Billiard Ball Model (1805)
• All matter is made of tiny, indivisible particles
called atoms.
• Each element is made up of atoms that are
unique and unlike atoms of any other
element.
• Atoms of an element are identical.
• Matter is composed of
indestructible, indivisible
atoms
Thomson’s Raisin Bun Model
(1897)
• Materials, when rubbed, can develop a charge difference.
• This electricity was called “cathode rays”
• These rays have a small mass and are negatively
charged.
• Thomson noted that these negative subatomic particles
(electrons) were a fundamental part of all atoms.
Rutherford’s Nuclear Model (1911)
• Rutherford shot alpha () particles at gold foil.
Zinc sulfide screen
Thin gold foil
Lead block
Radioactive
substance path of invisible
-particles
Most particles passed through. So,
atoms are mostly empty space.
Some positive -particles deflected
or bounced back!
Thus, a “nucleus” is positive
(protons) & holds most of an atom’s
mass.
Table 1, p. 26
Practice
• P. 26 #1-5
Atomic numbers, Mass numbers
Elements are often symbolized with their mass
number (A) and atomic number (Z)
16
E.g. Oxygen: 8
O
Z = # of protons = # of electrons
A - Z = # of neutrons
Calculate # of e–, n0, p+ for Ca, Ar, and Br
Atomic
Mass
p+
e-
no
Ca
20
40
20
20
20
Ar
18
40
18
18
22
Br
35
80
35
35
45
Bohr - Rutherford diagrams
• Putting all this together, we get B-R diagrams
• To draw them you must know the # of protons, neutrons,
and electrons (2,8,8,18 filling order)
• Draw protons (p+), (n0) in circle (i.e. “nucleus”)
• Draw electrons around in shells
He
p+
2
2 n0
Li
Li shorthand
3 p+
4 n0
3 p+ 2e–
4 n0
1e–
Draw Be, B, Al and shorthand diagrams for O, Na
Be
B
Al
4 p+
5 n°
O
5 p+
6 n°
13 p+
14 n°
Na
8 p+ 2e– 6e–
8 n°
11 p+ 2e– 8e– 1e–
12 n°
Isotopes and Radioisotopes
Isotopes: Atoms of the same element that have different
numbers of neutrons
– Due to isotopes, mass #s are not round #s.
– E.g. Li (6.9) is made up of both 6Li and 7Li.
– Often, at least one isotope is unstable.It breaks down, releasing
radioactivity.These types of isotopes are called radioisotopes
Q- Sometimes an isotope is written without its atomic number
- e.g. 35S (or S-35). Why?
Q- Draw B-R diagrams for the two Li isotopes.
A- The atomic # of an element doesn’t change Although the
number of neutrons can vary, atoms have definite numbers
of protons.
6Li
7Li
3 p+
3 n0
2e– 1e–
3 p+
4 n0
2e– 1e–
Practice
• P. 29 #1-7
Limitations to Rutherford’s
Model
• Orbiting electrons should emit light, losing
energy in the process
• This energy loss should cause the
electrons to collapse into the nucleus
• However, matter is very stable, this does
not happen
Bohr’s Planetary Model
• Electrons orbit the nucleus in energy “shells”
• An electron can travel indefinitely within a shell without
losing energy
• The greater the distance between the nucleus and the shell,
the greater the energy level
• An electron cannot exist between shells, but can move to a
higher, unfilled shell if it absorbs a specific quantity of
energy, or to a lower, unfilled shell if it loses energy
(quantized)
• When all the electrons in an atom are in the lowest possible
energy levels, it is in its ground state.
• An atom becomes excited when one of its electrons absorb
energy
• If enough energy is absorbed then the electron can make a
quantum leap to the next energy level, if there is room
• When the electron returns to a lower energy state the energy
is released in the form of a photon, which we see as visible
light
• The energy of the photon determines its wavelength
or color
• Each element has its own frequencies of color, so it
emits its own distinctive glow
Summary of Atomic Models
1) Dalton’s “Billiard ball” model (1800-1900)
Atoms are solid and indivisible.
2) Thomson’s “Raisin bun” model (1900)
Negative electrons in a positive framework.
3) Rutherford’s “Nuclear” model (~1910)
Atoms are mostly empty space.
Negative electrons orbit a positive nucleus.
4) Bohr’s “Planetary” model (~1920)
Negative electrons orbit a positive nucleus.
Quantized energy shells
5) Quantum Mechanical model (~1930)
Electron probabilities (orbitals)
Practice
• Pre-lab: Atomic Spectra (p. 40)
• p. 42 #1-3
• p. 45 #6-8