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INSTRUCTIONAL OBJECTIVES:
1)
DESCRIBE THE ATOMIC THEORY
2)
DETERMINE ATOMIC MASS, ATOMIC NUMBER, MASS NUMBER, NAME, AND SYMBOL OF ISOTOPES
3) STATE
THE RELATIONSHIP BETWEEN ATOMIC SPECTRA, QUANTUM NO., ATOMIC ORBITALS, AND
ELECTRON ENERGY LEVELS
4) DESCRIBE PERIODIC RELATIONSHIPS OF ELEMENTS INCLUDING
ATOMIC RADII, IONIZATION ENERGY, ELECTRON AFFINITY, AND
ATOMIC STRUCTURE AND PERIODIC TABLE NOTES
I. Early Theories
A)
Continuous Theory of Matter(~400B.C.)
1)
matter is infinitely divisible
2)
matter is various mixtures of only 4 elements: earth,
fire, water, and air
B) Atomic Theory of Matter(Democritus,
~400B.C.)
1)
matter is made of indivisible tiny particles, or “atomos”
2)
different matter was made of different atoms
3)
motion of atoms gave matter all its properties
4)
not based on experiment, but on philosophical argument
5)
rejected for relig. reasons
II. Modern Atomic Theory
A)
Law of Conservation of mass- “Mass is neither created nor destroyed in chemical
change”-(Antoine Lavoisier)
B) Law of
Definite Composition- “The proportion by weight of elements in a chemical
compound is constant”
-
(Joseph Proust)
Example: in carbon dioxide, the mass proportion of C
to O is always 3g:8g
C)
Dalton’s Atomic Theory - explained previous 2 laws
1)
matter is made of tiny, indivisible atoms that cannot
be created or destroyed
2)
each element has atoms identical to each other in all
properties; different elements are made of different atoms with different properties
3)
chemical reactions are simple rearrangements of atoms
from one combination to another, which are always in small whole number ratios
4)
He used his theory to make a prediction(below)
D.
Law of Multiple Proportions- “If a pair of elements can form more than one
compound, then the ratio of the mass of the first element that combines with a
fixed mass of the second element form simple, whole-number ratios”
-
(John Dalton)
Example:
In carbon monoxide, the proportion of C to O is 3g:4g
This means that the ratio of Oxygen in carbon dioxide to Oxygen in carbon
monoxide is
E)
Discovery of Cathode Rays- (William Crookes, 1870)
His cathode
ray tube consisted of a vacuum tube with two metal plates hooked to electric
current
Another experiment showed the following results
Crookes’ Tube
Experiment showed that:
1)
Rays originate from the (-) terminal(cathode) and
travel to (+) terminal(anode)
2)
All matter can emit these cathode rays
3)
Cathode rays carry (-) charge
F)
Discovery of electrons- (J.J. Thompson, 1897)
1)
Determined that cathode rays are made of (-) particles he called electrons,
symbol e
2)
Calculated the charge to mass ratio, e/m:
e/m = -1.76 x
108 Coulombs/g by measuring the deflection of
the electrons in electric and magnetic fields
3)
He proposed the “Plum-Pudding” model of the atom:
(-)
charged electrons float in a “sea” of positively charged fluid
, composition unknown
G)
Mass of Electron determined-(Robert Millikan, 1909)
His
“oil-drop” experiment:
Millikan’s Oil-Drop Experiment showed:
1)
x-rays produced electrons on the oil drops
2)
the mass of the drops could be balanced by the
electric force from the plates
3)
-1.6 x 10-19 C was the smallest charge he calculated
could be on any drop. He concluded this was value for the
charge on one electron, e
4)
Using Thompson’s e/m ratio, he
calculated the mass of the electron:
e =-1.76x108C =
-1.6x10-19 C
m g me
so, me
=
H)
Discovery of Nucleus and Proton -(Ernest Rutherford,
1909-1919)
The
gold-foil experiment showed:
1)
over 99.9% of the alpha particles passed through the
gold atoms without change
2)
of the other .1%, most showed only small deflections
from their original path
3)
about 1 in 10,000 alphas
were totally reflected by the gold foil
4)
5)
In 1919,
Charge
was = but opposite to charge on e, and mass was over 1800 times larger than me. He called the new particle the proton, symbol p
6)
Major problem with the “Planetary” model of the atom- Instability. All atoms would emit radiation and quickly
collapse according
to the known laws of physics.
I)
Atomic Number and Mass Number
1)
Henry Mosely(1913) found that the number of protons in an element in an
x-ray tube(like a Crookes tube)
#
protons =
k (x-ray frequency)1/2
The
number of protons in an atom’s nucleus is called its atomic number,
symbol Z.
Z
is with element in per. table
2)
James Chadwick(1932) discovered the neutron (symbol n)
also in nucleus. Its mass is ~ mass of
p, but carries 0 charge.
All the known subatomic particles
discovered by the time of Chadwick:
Particle Charge Mass
electron
proton
neutron
A
“u” or “amu” is 1 atomic mass unit.
i) The sum of proton + neutron mass is nearly the entire
mass of the atom. This mass is called
the mass number, symbol A.
A
= #p + #n
ii)
Units are amu, or u(for
atomic mass units).
1
amu = 1.66 x 10-24 g
3)
All atoms of the same element have same #p, #e, and atomic number Z.
i) Isotopes, with different # neutrons are possible.
Example: hydrogen has 3
isotopes, called-
protium:
deuterium:
tritium:
All three isotopes have the same atomic
number, but different masses.
ii)
Nuclear symbols of isotopes use Z and A numbers
So,
symbols of 3 hydrogen isotopes are:
protium deuterium tritium
iii)
Names of isotopes are based on their mass no., A
isotope name = (element name) - (mass no.)
For
the 3 hydrogen isotopes,
names are
iv) To find the # of neutrons in an isotope, subtract the
atomic no. Z(#p) from mass no. A(#p
+ #n)
#
neutrons = A - Z
Example:
Compare the no. of neutrons in carbon-12 and carbon-14:
Z
= 6 for element carbon.
A
= 12 for carbon-12
#n
= A - Z =
A
= 14 for carbon-14
#n
= A - Z =
Symbols
for carbon-12 and 14:
4)
Atomic Masses of elements in the periodic table
i) If there are 2 or more naturally occurring isotopes of
an element, then the average mass, called the atomic mass, is displayed
ii)
The atomic mass is calculated as follows:
We
must know the masses and abundances of each isotope of the element.
Example:
Boron is composed of 2 isotopes, boron-10 and boron-11.
isotope abundance mass
boron-10 19.8% 10.0
boron-11 80.2% 11.0
The
average, or atomic mass is calculated by the formula
at.mass = sum of (dec.%)(mass)
at.mass = S (dec. %) x
(mass)
at.mass = ( )( )+(
)( )
=
In the periodic table, the atomic mass
of boron is reported 10.8
J. Quark Theory
1)
particle physicists using colliders
have shown that protons and neutrons are not fundamental particles
2)
There are actually >100 known particles + antiparticles
Example: p, n, p(pion), L(lambda),
S(sigma), X(xi), etc.
3)
To account for all of these, Murray Gell-Mann of Cal-Tech. proposed the Quark
Model of subatomic particles.
Quark
Theory states that particles like p, n, etc. are made of various combinations
of quarks, which total 6 in number:
Quark charge
mass (amu)
up(u) +2/3 1/3
down(d) -1/3 1/3
charm(c) +2/3 75
strange(s) -1/3 10
top(t) +2/3 ~ 9000
bottom(b) -1/3 215
4)
All of the known particles except leptons (electrons) and bosons(photons,
or light) are made of either 2 or 3 quarks.
Example: protons
are made of 3 quarks; 2 up, and 1 down
neutrons are made of 3 quarks: 2 down, and 1 up
The
quark model now explains the existence of all the over 100 known subatom. particles
5)
The Strong Nuclear Force holds quarks together, and protons and neutrons
in the nuclei of atoms. Other forces are:
Weak
Nuclear Force- responsible for nuclear decay;
Electric
Force- responsible for electricity, chemical bonding;
Gravity- responsible for weight, orbits