Download Atomic Structure - Marlington Local Schools

October 03, 2014
Atomic Structure
Chapter 4
Democritus's Atomic Philosophy
1. Democritus
2. Dalton
3. Thompson
4. Millikan
5. Moseley
6. Bohr
7. Rutherford
8. Schrodinger
Aristotle
-Rejected the idea of the atom
-Said matter could be cut continually
*most people sided with Aristotle
Dalton's Atomic Theory
Dalton transformed Democritus's ideas on atoms into
a scientific theory. He studied the ratios in which
elements combine in chemical reactions.
(Concrete Evidence)
There are 4 parts to this theory:
1. All elements are composed of tiny indivisible
particles called atoms.
2. Atoms of the same element are identical. The
atoms of any one element are different from those
of any other element.
Atom:
the smallest particle of an element that retains its
identity in a chemical reaction.
3. Atoms of different elements can physically mix
together or can chemically combine in simple wholenumber ratios to form compounds.
4. Chemical reactions occur when atoms are separated,
joined, or rearranged. Atoms of one element, however,
are never changed into atoms of another element as a
result of a chemical reaction.
October 03, 2014
Atomic Size
Copper atoms are very small. A pure copper coin the size
of a penny contains about 2.4 x 1022 atoms.
Despite their small size, individual atoms are observable
with instruments such as scanning tunneling
microscopes.
Iron Atoms
Subatomic Particles
Electrons
The atoms are now broken down into even smaller, more
fundamental particles, called subatomic particles.
J. J. Thomson discovered the electron. He thought some
smaller particle had to exist.
There are 3 subatomic particles
Electrons : negatively charged subatomic particle.
1. electrons (e-)
He used the cathode ray tube
to make this discovery.
2. protons (p+)
3. neutrons (no)
How did J. J. Thomson make this discovery?
The cathode ray tube has a cathode and anode at each
side. The cathode originally is positively charged and the
anode negatively charged. When the electric current
was based through, the cathode became negatively
charged and the anode became a positive charge.
http://www.youtube.com/watch?v=XU8nMKkzbT8
Robert A. Millikan
*He was then able to calculate the mass of the electron.
Protons and Neutrons
How did we determine the presence of protons and neutrons?
Four ideas about matter and electric charges:
1. Atoms have no net electric charge.
2. Electric charges are carried by particles of matter.
3. Electric charges always exist in whole-number multiples of a
single basic unit. (There are no fractions of charges.)
4. When a given number of negatively charged particles
combines with an equal number of positively charged particles,
an electrically neutral particle is formed.
October 03, 2014
Protons and Neutrons
The Atomic Nucleus
Proton: Positively charged subatomic particles.
Confirmed by Eugen Goldstein
Neutron: Subatomic particles with no charge but with a
mass nearly equal to that of a proton
Confirmed by James Chadwick
Scientists wondered how these particles were put together in an
atom.
J. J. Thomson came up with a model known as the
"plum-pudding" model.
This model stated how electrons were stuck into a lump
of positive charge, similar to raisins stuck in dough.
This model did not last very long when Ernest Rutherford entered
the picture.
The Atomic Nucleus
Rutherford Atomic Model
http://www.youtube.com/watch?v=5pZj0u_XMbc
Ernest Rutherford: Gold-Foil Experiment
He proposed that the atom is mostly empty space, explaining
the lack of deflection of most of the particles.
All the positive charge and almost all the mass are concentrated
in a small region called the nucleus.
Nucleus: Tiny central core of an atom and is composed of
protons and neutrons.
Atomic Number:
The number of protons in the nucleus.
Atomic Number
Element
Atomic
Number
K
19
Protons
Electrons
19
5
S
V
K+
S 2-
16
23
October 03, 2014
Mass Number
Shorthand Notation
197
79
Au
Au-197
Gold-197
Sample Problem 4.1
How many protons, electrons, and neutrons are in
each atom?
Atomic Number
Mass Number
4
9
A. Beryllium (Be)
10
20
B. Neon (Ne)
11
23
C. Sodium (Na)
A. Beryllium (Be)
B. Neon (Ne)
C. Sodium (Na)
Protons
4
10
11
Electrons
4
10
11
Neutrons
5
10
12
Practice:
Express the following in shorthand notation
a) carbon-12
b) fluorine-19
c) beryllium-9
How would you then determine the number of neutrons in an atom?
Number of neutrons = mass number - atomic number
X= Element Symbol
A= Mass Number
Z= Atomic Number
Answer
Mass Number:
The total number of protons and neutrons in an atom.
Practice:
How many neutrons are in each atom?
16
a) 8 O
b)
32
16
c)
108
47
Ag
d)
80
35
Br
e)
207
82
S
Pb
Isotopes: Atoms that have the same number of protons
but different numbers of neutrons
* Because isotopes of an element have different numbers
of neutrons, they also have different mass numbers.
October 03, 2014
Uses of Isotopes
1. Americum-241: Used in many smoke detectors
2. Carbon-14: Used in biological research, agriculture,
pollution control and archeology.
3. Cesium-137: Used to treat cancerous tumors, control
liquid flow in oil pipelines.
4. Iodine-123: Used to diagnose thyroid disorders.
5. Iodine-129: Used to check radioactivity counters in in-bitro
diagnostic testing labs.
6. Iodine-131: Used to treat thyroid disorders. (Graves's
disease)
7. Plutonium-238: Has powered more than 20 NASA
spacecraft since 1972.
8. Uranium-234: Used in dental fixtures like crowns and
dentures to provide a natural color and brightness.
9. Uranium-235: Fuel for nuclear power plants
In nature, most elements occur as a mixture of two or
more isotopes. Each isotope of an element has a fixed
mass and a natural percent abundance.
We are then able to determine the isotope that is in
more abundance due to the average (atomic mass) of
the element.
Example:
The two stable isotopes of chlorine are chlorine-35 and
chlorine-37.
The average SHOULD be 35.969 amu. However,
this value is higher than the actual value of 35.453.
So, chlorine-35 is in higher abundance than
chlorine-37.
Atomic Mass
Practice: Isotopes
1) Three isotopes of oxygen are oxygen-16, oxygen-17
and oxygen-18. Write the symbol for each, including the
atomic number and mass number.
It is more useful to compare the relative masses of atoms using a
reference isotope as a standard. This standard is Carbon-12.
Atomic mass unit (amu):
One twelfth of the mass of a carbon-12 atom.
Example: A helium-4 atom, with a mass of 4.0026 amu has about
one-third the mass of a carbon-12 atom.
2) Three isotopes of chromium are chromium-50,
chromium-52 and chromium-53. How many neutrons
are in each isotope, given that chromium has an atomic
number of 24?
Practice Problems
1) Boron has two isotopes: boron-10 and boron-11. Which
is more abundant given that the atomic mass of boron is
10.81?
Table 4.3
Average Atomic Mass:
a weighted average mass of the atoms in a naturally occurring
sample of the element. This average reflects both the mass and
the relative abundance of the isotopes as they occur in nature.
Calculating Atomic Mass of an Element
Multiply the mass of each isotope by its natural abundance,
expressed as a decimal, and then add the products.
Atomic Mass of Carbon:
2) There are three isotopes of silicon; they have mass
numbers of 28, 29, and 30. The atomic mass of silicon is
28.086 amu. Comment on the relative abundance of these
three isotopes.
C-12: 98.89%, 12.00 amu
C-13: 1.11%, 13.003 amu
Atomic mass= (12.000 amu x 0.9889) + (13.003 amu x 0.0111)
Atomic mass= 12.011 amu
October 03, 2014
Calculating Atomic Mass of an Element
Calculating Atomic Mass of an Element
1) The element copper has naturally occurring isotopes with
mass numbers of 63 and 65. The relative abundance and
atomic masses are 69.2% for mass-62.93 amu, and 30.8%
for mass=64.93 amu. Calculate the average atomic mass
of copper.
2) Calculate the atomic mass of bromine. The two isotopes
of bromine have atomic masses and relative abundance of
78.92 amu (50.69%) and 80.92 amu (49.31%).