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Chapter 1000A
Spring 2014
Chapter 2 & 25
Atomic structure and Nuclear Chemistry
History of the Atom
– Greek philosophers
– Thales: All matter was water
– Empedocles: All matter was a fusion
of opposites.
– Aristotles: philosophical elements (air,
earth, fire & water and principles (Hot,
cold, moist & dry) ~2000 yrs
– Leucippus and Democritus:
Atomos: The point at which matter can
no longer be subdivided.
a = not & tomos = cut
• Phlogistons (Bechner & Stahl)
• Combustion
• Metal burns
• Wood burns
• Joseph Priestly burned Hg with air in a
sealed container making HgO. Priestly
discovered O2 .
Antoine Lavoisier (1743-1794): observed
no change in mass.
• 1785, Law of Conservation of Mass:
.
• 1794, Joseph Proust (1754 -1826):
– Law of definite proportions:
• In a given chemical compound the
proportions of the elements that compose
it are fixed integers (independent of its
origin or preparation).
– lead sulfide
– Law of multiple proportions:
• The same elements can combine in
different proportions to form different
chemical compounds (independent of its
origin or preparation).
– water & peroxide
– carbon oxides I & II
– (do calc Wt% - chalk talk)
• 1808, John Dalton (1766-1844) adopts
ideas of Democritus, Lavoisier & Proust
to formulate Dalton’s Atomic theory:
1. All matter consist of solid and indivisible
atoms
2. Atoms are indestructible and retain their
identity in all chemical reactions.
3. All atoms of a given element are identical in
mass and other properties.
– *Different elements have different kinds of
atoms and differ in mass from others.
4. Compounds consist of elements combined
in small whole number ratios.
• 1808, John Dalton (1766-1844) adopts
ideas of Democritus, Lavoisier & Proust
to formulate Dalton’s Atomic theory:
1. All matter consist of solid and indivisible
atoms
2. Atoms are indestructible and retain their
identity in all chemical reactions.
3. All atoms of a given element are identical in
mass and other properties.
4. Compounds consist of elements combined
in small whole number ratios.
• Any theory is only as good as the
current data/evidence.
Table 2.2
Properties of the 3 Key Subatomic Particles
Charge
Name
(Symbol)
Relative Absolute (C)*
Proton
(p+)
1+
Neutron
(n0)
Electron
(e-)
Mass
Relative
(u)†
Absolute (g)
Location in
Atom
+1.60218x10-19
1.00727
1.67262x10-24
Nucleus
0
0
1.00866
1.67493x10-24 Nucleus
1-
-1.60218x10-19
0.00054858
9.10939x10-24 Outside
nucleus
Charge =
JJ Thomson –
Eugen Goldstein –
James Chadwick –
* The
†
coulomb (C) is the SI unit of charge.
The atomic mass unit (u) equals 1.66054x10-24 g.
Rutherford’s a-scattering experiment and discovery of the atomic nucleus.
Rutherford proposed that atoms consisted of tiny, massive,
positive nucleus surrounded by electrons
Most of an atom is empty space:
(volume of 5 x 10-46 m3)
• A tiny nucleus (about
1/1,000,000,000,000,000th of
the atom’s volume)
• contains most of the atom’s
mass: protons and neutrons
• bound together in a region of
positive charge
• electrons travel around the
nucleus, balancing the
overall charge of the atom.
• Give analogy:
Defining an element
• The relative mass and charge of a particle is
described by the notation:
• An element
.
A = mass number
Z = charge of the particle
Example:
electron
proton
neutron
0
e
-1
1
p
1
1
n
0
A
Z
X
• Every atom has an atomic number and a
mass number:
– Mass number (A) = # nucleons*
C
– Atomic number (Z) = # protons
12
6
*nucleon is a general term for a proton or neutron (i.e. a particle in the nucleus)
• Atoms of the same element must have
the same atomic number but different
mass numbers are possible.
• These are
. Most elements
have more than one naturally occurring
isotope:
H
H
H
1
2
3
1
1
1
12
13
14
6
C
6
C
6
C
12
Isotopes & applications
• Isotope ratios
• Climatology
• A forensic application1 developed in Sweden
– Atmospheric 14C levels increased sharply
during testing of nuclear bombs
C
C
C
(1955-1963).
– 14C levels decreased since banning of tests
12
13
14
6
6
6
K.L. Spalding, B.A. Buchholz, L.-E. Bergman, H. Druid and J. Frisén
“Age Written in Teeth by Nuclear Tests”
Nature 437, 333 - 334 (14 Sep 2005).
1
13
The amount of 14C incorporated in tooth
enamel can be used to determine the year in
which the enamel was formed and, therefore,
the year of birth (to within 1.6 years).
How Can We Measure Isotopes?
• The average atomic mass for a chlorine atom
is 35.4527 amu. How is the isotopes present
in a sample determined?
– Revisit the main difference
– Take advantage of that difference
15
The end result is a spectrum showing the
proportion of atoms in the sample belonging
to each isotope:
This is mass spectrometry! (the “MS” in GC/MS)
16
• Most elements occur as mixtures of several
isotopes. These are treated as consisting of
“averaged” atoms with “averaged” masses.
• Atomic mass –

Mav  

% abundance of isotope
 Misotope
100%
Where
% abundance 
# atoms of isotope
 100%
total # atoms of element
Exceptions: 19F, 23Na, 31P
17
“Average” Samples
• Silicon has three naturally occurring
isotopes:
– 92.23% 28Si (27.9769 u)
– 4.67% 29Si (28.9765 u)
– 3.10% 30Si (29.9738 u).
First estimate
then calculate the average atomic mass of
silicon.
18
• First, the isotope present > % is
• Ave. atomic weight Si =
• Ave. atomic weight Si =
• Ave. atomic weight Si =
.
Quick revisiting of the “mole” Concept.
• How many atoms of Galium (Ga) are in
2.85 x 10-3 mol Ga?
• 2.85 x 10-3 mol Ga x
• =
.
atoms
• Gallium has two naturally occurring
isotopes and an average atomic mass of
69.723 u. 69Ga has an atomic mass of
68.926 u, and 71Ga has an atomic mass
of 70.925 u.
• First predict which isotope is more
abundant
• then calculate the natural abundance of
each isotope of gallium.
21
69Ga
68.926 u
the average amu Ga is 69.723 u
71Ga 70.925 u.
• First calculate the differences from the
average reported: 0.797u and 1,202
therefore it would be 69Ga.
There are 2 variables , 2 equations are needed.
MWGa =
Or using mole concept & mole fraction
MWGa =
Also have the second equation
x+y=1