Atomic Theory
&
Atomic Structure
Tro, Chapter 4 & 9
Sections 4.1 – 4.4, 4.8, 4.9; 9.2 – 9.9
Document BIG IDEAS about:
• Atomic structure
– Electrons (mass, size, position)
– Protons and neutrons (mass, position)
– Isotopes
• Changes in (MODERN) thought
– Dalton
– Thomson, Rutherford and Bohr
• Quantum theory (CONTEMPORARY)
Early Atomic Theories
Democritis
(400 BCE)
• First to propose idea of atom
• Atom = “a” + “tomos” = cannot be cut
• Based solely on logic; not supported
by experiments
Alchemy
(12-1500 CE)
• Modern word ‘chemistry’ came from
Arabic ‘alkimiya’
• recognized importance of
experimentation
• Responsible for developing lab
equipment & procedures still used
today
NOTE: Alchemy is a field, NOT a person…
Galileo
(~1600 CE)
• Birth of modern science - combining
logic, experimenting, publishing
results
Lavosier & Priestly
(1700’s)
• Quantitative
analysis of
chemicals
Law of Conservation of
Mass:
Matter can neither be
created nor destroyed
Proust
(1700’s)
• Developed Law of Definite Proportions
Law of Definite Proportions:
Different samples of the same
compound always contain its
constituent elements in the same
proportions by mass
Law of Definite Proportions
• Copper carbonate always contains
– 5.3 parts copper
– 4 parts oxygen
– 1 part carbon
by mass
Dalton
(1800’s)
• School teacher that proposed the first
modern-day idea of atoms
Law of Multiple Proportions:
If 2 elements combine to form more than one
compound, the masses of one element that
combine with a fixed mass of the other element
are in small whole # ratios
Law of Multiple Proportions
Dalton’s Atomic Theory - 1808
• All matter is composed of atoms which
cannot be subdivided
• Atoms of same element are identical (size,
mass, reactivity)
• Atoms combine to form compounds in
simple, whole # ratios
• Chemical reactions involve the separation,
combination, or rearrangement of atoms; it
does not result in their creation or
destruction
Modern Atomic
Theories
General Principle #1
Electric Charges
+
–
positive
negative
Objects with an equal amount of positive
and negative charge are said to be
electrically neutral
General Principle #2
Forces between Charges
• Objects with like charge repel
+ +
+ +
• Objects with opposite charge attract
+ +
– –
Forces between Charges
• Electrostatic force becomes greater
with more charge
• Electrostatic force becomes smaller
the greater the distance between the
charges
Thomson’s Atomic Model (1904)
Cathode Ray Experiments
• Any metal worked for
anode
• Negative electric field
repelled beam
• Object placed in path
of glow blocked beam
J.J. Thomson’s Contribution
• Discovered the electron (1897)
• Plum Pudding model
• Determined the charge-to-mass ratio of
an electron using data from cathode
ray tube experiments
Evidence & Conclusions
• cathode rays consisted of subatomic
particles from atoms of anode
• cathode rays are negatively charged
•
\ must also be positive charge
• Millikan (oil drop experiment, 1909) calculated
electron’s mass to be 9.11 x 10-31 kg
Modern View of Atomic Structure
Particle
Symbol
Relative
Charge
Mass (kg)
proton
+
p+
+1
nucleons
neutron
0
n0
0
1.67510 x 10-27
e-
-1
9.1096 x 10-31
electron
1.6726 x 10-27
Modern View of Atomic Structure
Particle
Relative
Charge
Mass (kg)
Relative
mass (amu)
p+
+
+1
1.6726 x 10-27
~1
n0
0
0
1.67510 x 10-27
~1
-1
9.1096 x 10-31
~0
e-
Rutherford’s Problems
• How is nucleus held together?
• Why don’t electrons collapse into
nucleus?
• H atom has 1 proton & He atom has 2
protons, \ mass ratio should be 2:1;
instead the ratio is 4:1
…there must be another particle
The Gold Foil Experiment: Hypothesis
• The α-particles will
pass straight
through the atoms
What is an () alpha particle?
4
2
He
It is a positively charged Helium nucleus
Rutherford’s Gold Foil Experiment
The Gold Foil Experiment: Outcome
What’s happening?
The Gold Foil Experiment: Conclusions
Atoms :
• must be mostly space
• must have a very
small, dense area of +
charge
• Protons have same
charge as e-, but
almost 2000x more
mass!
The Neutron
• Discovered by James Chadwick in 1932.
• Neutron is electrically neutral & has
slightly greater mass than a proton
Mystery solved.
Atomic theory timeline
Updating Dalton’s Atomic Theory
3 major differences between modern atomic
theory & Dalton’s atomic theory:
• Atoms are NOT indivisible – they are made
up of protons, neutrons, and electrons
• Atoms of the same element are NOT exactly
alike – they can have different masses
(isotopes)
• Atoms CAN be changed from one element to
another, but not by chemical reactions
(nuclear reactions)
Atomic Structure &
Isotopes
Atomic Mass Unit (amu)
• defined as a more convenient unit for
reporting mass of small numbers of
atoms
• 12C is used as the reference
• 1 amu is defined as exactly 1/12 of a
12C atom
Getting Information from the
Periodic Table
6
C
12.0111
Atomic # = # p+ in nucleus
Elemental symbol
Atomic mass
(more on this later)
Isotopic Notation
• Atomic number (Z) = # of p+ in the nucleus
• Mass number (A) = sum of # p+ & n0 in
nucleus
• For a neutral atom, # e- = # p+
Examples
1
H
1
4
2 He
12
C
6
16
8O
Mass number (A)
Atomic number (Z)
63
Zn
30
Isotopes
• All atoms in an element have the same
atomic number
• However, 2 atoms of the same element
can have different mass numbers –
called isotopes
• Isotopes have:
– Same # of p+
– Different # of no
Some Common Isotopes
1
H
1
12
C
6
235
92 U
2
1H
13
C
6
238
92 U
3
1H
14
C
6
Relative
Abundance
Mass Spectrometry
• Technique used to determine atomic
mass
e
e- collides with
atom, “bounces”
off, but transfers
some energy to it
e-
e-
+
Atom dissipates
excess energy by
expelling an electron
Atom bombarded by stream of high energy electrons
Mass Spectrometry, cont.
• Ions are accelerated through a magnetic field
• Amount of deflection depends on the ion’s mass
• Highest mass deflected least
• Lowest mass deflected most
+
+
N
+
+
S
++
+
+
Mass Spectrometry, cont.
Sample mass spec for chlorine
Relative abundance of each
isotope can be determined
from relative peak heights
Mass (amu)
35
37
Relative Abundance & Atomic Mass
• Relative isotopic abundance is then
used to calculate atomic mass
• Atomic mass is the weighted average of
the mixture of isotopes
Example
Calculate the atomic mass of Cl given the
relative abundances of its isotopes:
35Cl – 75.77%
37Cl – 24.23%
average atomic mass
= (atomic mass 35Cl)(fraction 35Cl) + (atomic
mass 37Cl)(fraction 37Cl)
= (34.968 amu)(0.7577) + (36.965
amu)(0.2423)
= 35.45 amu
Practice
Copper, a metal known since ancient times, is
used in electrical cables & pennies, among
other things. The atomic masses of its 2 stable
isotopes, 63Cu (69.09%) and 65Cu (30.91%) are
62.93 amu and 64.9278 amu, respectively.
Calculate the average atomic mass of copper –
the relative abundances of each ion is given in
parentheses.
Answer: 63.54 amu
The Bohr Model
Electromagnetic Spectrum
Light
c = lu
c = speed of light (3.0 x 108 m)
l = wavelength
u = frequency
Frequency vs. Wavelength
Light
c = lu
• Energy  as frequency 
• Energy  as wavelength 
• Light behaves like a particle (photon)
as well as a wave
Emission Spectrums
• When electricity is run through a
sample of hydrogen gas, hydrogen
atoms gain energy
• H atoms loose that energy by emitting
photons
• Resulting spectrum is discontinuous
continuous
discontinuous
What’s happening?
Bohr Model
• Electrons move in
circular orbits around
the nucleus
• Only certain energy
levels are “permitted”
(this explains the
discrete lines for the
emission spectrum of
hydrogen)
Schroedinger/Heisenburg
• Experiments used mathematics
(probability) to predict behavior of
electrons
– Schroedinger equation approximated the
probability of finding a single electron
for H within a region close to the nucleus
– Heisenburg [Uncertainty Principle]
reinforces the idea that we just don’t
know!
Math in Context: Blackbody
Experiments