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Chapter 3 – Atoms: The Building Blocks of Matter
I.
The Atom: From Philosophical Idea to Scientific Theory
A.
The Philosophers
1.
Democritus (~400 BC) –
2.
Another Greek philosopher of the time, Aristotle, believed
*
B.
Foundations of Atomic Theory
Law of Conservation of Mass –
Law of Definite Proportions –
Law of Multiple Proportions-
C.
John Dalton’s Atomic Theory (1808)
1.
Basic Points
a.
b.
c.
d.
e.
D.
2.
The Law of Conservation of Mass is explained by point e
3.
The Law of Multiple Proportions is explained by point d
Modern Atomic Theory
1.
Some of Dalton’s points were not exactly accurate and have
been modified throughout the years
2.
Two important concepts remain unchanged:
a.
b.
II.
The Structure of the Atom
A.
Discovery of the Electron
1.
Cathode Rays and Electrons (Late 1800s)
a.
b.
When current was passed through the cathode ray
tube, the surface of the tube opposite the cathode
glowed
c.
1.
Hypothesized that the glow was caused by a
stream of particles called “cathode rays”
2.
Rays seemed to travel from the cathode to the
anode
Further experiments supported the hypothesis and
added information:
1.
Included a paddle wheel which, when placed
on rails between the electrodes, rolled along
the rails from the cathode to the anode
2.
Cathode rays, when exposed to electricallycharged plates, were deflected by the negative
plate and attracted to the positive plate
* This information led to the hypothesis that
2.
3.
Two Important Scientists
a.
JJ Thomson (1897) –
b.
Robert Milliken (1909) –
Two Other Inferences
a.
Because atoms are electrically neutral, they must
contain
b.
Because electrons have such a small mass, atoms
must contain
B.
Discovery of the Atomic Nucleus
1.
2.
3.
The Gold Foil Experiment (1911)
a.
Carried out by Ernest Rutherford, Hans Geiger, and
Ernest Marsden (among others)
b.
The Experimental Design
1.
A thin, gold foil was bombarded with alpha
particles which were
2.
The foil was surrounded by a screen that would
detect where the alpha particles hit after
passing through the foil
The Results
a.
Most of the alpha particles went straight through the
foil or were only slightly deflected (this was expected)
b.
HOWEVER,
Rutherford’s Explanation (1913)
a.
b.
c.
C.
Composition of the Atomic Nucleus
1.
Except for the simplest type of hydrogen atom, all atomic
nuclei are made up of two kinds of particles,
Particle
2.
Electrons travel
3.
Summary of Subatomic Particles
Symbols
Relative
Electric
Charge
Mass
Number
(A)
Relative
Mass
(amu)
Actual
Mass
(g)
Electron
Proton
Neutron
D.
The Sizes of Atoms
1.
The region of the atom occupied by the electron can be
thought of as
2.
The radius of an atom is defined as
3.
Atomic radii are expressed in units of picometers (pm)…or,
in other words,
III.
Counting Atoms
A.
Atomic Number (Z) –
1.
Every element has a different atomic number…this identifies
the element!
2.
Found on the periodic table…in each element block near the
symbol of the element
3.
B.
In a neutral atom, there will be
Mass Number (A) –
1.
MUST BE GIVEN TO YOU!!! NOT on the periodic table!
2.
Atoms of the same element may have different numbers of
neutrons and, therefore, different mass numbers
a.
These atoms are referred to as isotopes…
b.
Example - Hydrogen
Protium (Most
Deuterium
common form of
“Heavy
H…
Hydrogen”
99.985%)
(0.015%)
Tritium
“Radioactive
Hydrogen”
(trace)
#
protons
#
neutrons
Mass
Number
c.
C.
Although isotopes have different masses, they do not
significantly differ in chemical behavior
Designating Isotopes
1.
Nuclear Symbol
2.
Hyphen Notation
3.
Using either designation, one can determine the number of
protons, neutrons, and electrons…
4.
Nuclear
Symbol
a.
# protons = atomic number (Z) = # electrons
b.
# neutrons =
Examples
Hyphen
Notation
#
Protons
#
Neutrons
#
Electrons
20
21
20
35Cl
131I
Uranium-235
5.
D.
Nuclide –
Relative Atomic Masses of the Elements
1.
Because atoms are so small, using mass in units of grams is
inconvenient and hard to understand
1 atom hydrogen-1 = 1.673 x 10-24 grams
2.
Scientists set up a relative scale of atomic masses…known
as the atomic mass scale
a.
Carbon-12 was arbitrarily assigned a mass of 12
atomic mass units (amus)
b.
One atomic mass unit (amu), or 1 amu, is
c.
E.
The mass of all other nuclides are determined by
comparing each one with the mass of a carbon-12
atom
Average Atomic Masses of the Elements
1.
The mass found in each element block on the periodic table
must account for all isotopes of that element that are in
existence and how abundant they are in nature
2.
Average atomic mass -
a.
To do this, use the following formula for each isotope:
(% occurrence  100) (mass)
b.
3.
Add all values together to get the average atomic
mass
Example
a.
In a box containing 2 sizes of marbles, 25% of the
marbles have a mass of 2.00 g each and 75% of the
marbles have a mass of 3.00 g each. What is the
weighted average of the marbles?
b.
Find the average atomic mass for oxygen using the
following information:
Percentage
Natural
Abundance
99.762
0.038
0.200
Isotope
Oxygen-16
Oxygen-17
Oxygen-18
4.
Atomic
Mass
(amu)
15.994915
16.999131
17.999160
In our textbook, an element’s atomic mass is usually
rounded to two decimal places before it’s used in a
calculation
IV.
Basic Bonding
A.
Atoms are held together (forming compounds) by chemical bonds;
there are two types
1.
Ionic Bonding -
2.
Covalent Bonding -
B.
Determining the Type of Bonding Based on the Chemical Formula
1.
If the compound is ionically bonded…
2.
If the compound is covalently bonded…