Ch. 4.2 - Atomic Structure
I. Subatomic Particles
(p.113 - 114)
Subatomic Particle Properties
Particle
Symbol
Location
Charge
electron
e-
Electron
cloud
proton
p+
nucleus
+
1
neutron
n0
nucleus
0
1
–
Relative
Mass
(amu)
Actual
Mass (g)
1/1840
approx 0
9.11 x 10-28
1.67 x 10-24
1.67 x 10-24
Symbols

Elements are listed by their chemical symbols

Symbols are usually either one capital letter
like C for Carbon, or one capital and one
lowercase letter like Ne for Neon
Periodic Table

The periodic table gives much information we
need to learn more about the atom of each
element
Atomic Number



Atomic number = # of protons in an atom
Whole number shown on periodic table
Periodic table is arranged by atomic number
Atomic Mass


The average atomic mass is the
number at the bottom of this square
Found by averaging the natural
abundances of its isotopes
Atom Math
Protons
Protons
Electrons
Neutrons
# n0 = Atomic mass – Atomic
Subatomic Particles
ATOM
ATOM
NUCLEUS
NUCLEUS
ELECTRONS
ELECTRONS
PROTONS
PROTONS
NEUTRONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE
CHARGE
NEGATIVE CHARGE
equal
in a
0 = Atomic
Atomic
Number
#n
mass
Most of the atom‟s mass.
neutral
equals the # of...
- Atomic
# atom
Ch. 4.3 - Atomic Structure
II. How Atoms Differ (p. 114 - 121)
Mass Number
Isotopes
Relative Atomic Mass
Average Atomic Mass
A. Mass Number

mass # = protons + neutrons
always a whole
number
NOT on the
Periodic Table!
© Addison-Wesley Publishing Company, Inc.
B. Isotopes

Atoms of the same element with different
numbers of neutrons
Isotope notation:
12
6
Mass #
Atomic #
Element name
Mass #
Isotope name: carbon-12
C
B. Isotopes

Isotope notation:
Chlorine-37

atomic #:
17

mass #:
37

# of protons:
17

# of electrons:
17

# of neutrons:
20
37
17
Cl
Natural Abundances of
Isotopes


Most elements are found as mixtures of
isotopes
Relative abundance of each isotope is the
same in each source
an atom or molecule in which the total number
of electrons is not equal to the total number of
protons, giving it a net positive or negative
charge.


Cation: has a positive charge, due to the loss
of electrons
Anion: has a negative charge, due to the
gain of electrons
C. Relative Atomic Mass
 12C
atom = 1.992 × 10-23 g
atomic mass unit (amu)
1 amu = 1/12 the mass of a 12C atom
1 p = 1.007276 amu
1 n = 1.008665 amu
1 e- = 0.0005486 amu
© Addison-Wesley Publishing Company, Inc.
D. Average Atomic Mass


weighted average of all isotopes
on the Periodic Table
Avg.
Atomic
Mass
(mass)(%) (mass )(%)
100
D. Average Atomic Mass

EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O,
and 0.20% 18O.
Avg.
Atomic
Mass
(16)(99.76 ) (17)(0.04)
100
(18)(0.20)
16.00
amu
D. Average Atomic Mass

EX: Find chlorine‟s average atomic mass if
approximately 8 of every 10 atoms are chlorine-35
and 2 are chlorine-37.
Avg.
Atomic
Mass
(35)(80) (37)(20)
100
35.40 amu
Atomic Theory
Development of our
understanding
of the atom
Early Models
Battle of Philosophers

Aristotle
vs. Democritus
-Matter is infinite
- Matter is composed of
extremely small particles
-4 Basic „elements‟
- Earth
-Air
-Fire
-Water
- Called these particles „atoms‟
(From „atmos‟ meaning invisible)
*Eventually after many years
Democritus is proven right
Early Models

Dalton‟s model was the “Billiard Ball”

Published an Atomic Theory
1. All matter is composed of atoms that are
indivisible (did not know about protons,
electrons, or neturons)
2. Atoms of a given element have same size, mass and
chemical properties and are different from those of
another element (no longer true, b/c of isotopes)
3. Different atoms combine in whole number ratios to
form compounds and are separated, combined and
rearranged in chemical reactions
4. In chemical reactions atoms can combine or
separate, but are neither created or destroyed
Law of Definite Proportions
Each compound has a specific ratio of
elements
 It is a ratio by mass
 Water is always 8 grams of oxygen for
each gram of hydrogen

Law of Multiple Proportions

if two elements form more than one
compound, the ratio of the second
element that combines with 1 gram of
the first element in each is a simple
whole number.
What?
Water is 8 grams of oxygen per gram of
hydrogen.
 Hydrogen Peroxide is 16 grams of
oxygen per gram of hydrogen.
 16 to 8 is a 2 to 1 ratio
 True because you have to add a whole
atom, you can‟t add a piece of an atom.

Parts of Atoms
J. J. Thomson - English physicist. 1897
 Made a piece of equipment called a
cathode ray tube.
 It is a vacuum tube - all the air has been
pumped out.

Thomson’s Experiment
Voltage source
-
+
Vacuum tube
Metal Disks
Thomson’s Experiment
Voltage source
-
+
Thomson’s Experiment
Voltage source
-
+
Thomson‟s Experiment
Voltage source
-
+
Thomson‟s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the negative
to the positive end
Thomson‟s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the negative
to the positive end
Thomson‟s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the negative
(cathode) to the positive end (anode)
Thomson‟s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the negative
(cathode) to the positive end (anode)
Thomson’s Experiment
Voltage source

By adding an electric field
Thomson‟s Experiment
Voltage source
+
 By adding an electric field
Thomson‟s Experiment
Voltage source
+
 By adding an electric field
Thomson‟s Experiment
Voltage source
+
 By adding an electric field
Thomson‟s Experiment
Voltage source
+
 By adding an electric field
Thomson‟s Experiment
Voltage source
+
 By adding an electric field
Thomson‟s Experiment
Voltage source
+
 By adding an electric field he found the ratio of
electrical charge to mass (e/m) for an electron
 The e/m ratio is (negative) 1.76 x
108coulombs per gram (or C/g in SI units).
Thomsom’s Model
Thomson always found the
same value for the e/m ratio
no matter what the tube
materials or the gas inside.
 Reinforced the notion that the
electrons are a fundamental
component of matter.
 „Plum Pudding‟ model: a thin
positive fluid, which contains
most of the mass, w/ negative
electrons embedded to
balance the charge

Other pieces
Proton - positively charged pieces 1840
times heavier than the electron
 Neutron - no charge but the same mass
as a proton.
 Where are the pieces?

Millikan used oil drop experiment


Would spray a fine mist of oil droplets
above a pair of parallel plates. Some of the
oil drops would pass through the hole in the
top plate.
He then used X-rays to knock electrons off
of the air molecules in the barrel and some
of those electrons attached themselves to
the oil drops. The oil drops, which were
now negative, could now be affected by the
electrical field. He then could now measure
the charge of the oil drops.





Millikan found that all the values he obtained
were whole-number multiples of -1.60 x 10-19
coulomb. This value must be the charge of
an electron.
The electron‟s charge was -1.60 x 10-19 coulombs
Using two values and solving for m
- 1.60 x 10-19 coul = - 1.76 x 108 coul/g
m
m = 9.11 x 10-28 grams
(a negligible mass even in the smallest atom)


Confirmed the negative charge of an electron
Determined mass of the electron
Rutherford’s experiment
Ernest Rutherford English physicist.
(1910)
 Believed in the plum pudding model of
the atom.
 Wanted to see how big they are using
radioactivity
 Alpha particles - positively charged
pieces given off by uranium
 Shot them at gold foil which can be
made a few atoms thick

Lead
block
Uranium
Florescent
Screen
Gold Foil
He Expected
The alpha particles to pass through
without changing direction very much
 Because
 The positive charges were spread out
evenly. Alone they were not enough to
stop the alpha particles

What he expected
Because, he thought
the charge was evenly
distributed in the atom
What he got
+
How he explained it
+
Atom is mostly empty
 Small dense, positive piece
at center (nucleus)
 Refined the concept of the nucleus &
concluded it was composed of positively
charged particles called protons
 James Chadwick: discovered a neutral atomic
particle with a mass close to a proton. Thus
was discovered the neutron.

Moving Forward…

Neils Bohr said electrons move in orbits



Found in energy levels
Explains bright-line spectrum
Called “Solar System Model” where
Electrons move in orbits around the nucleus
What we believe now
Heisenberg/Schrodinger



Heisenber Uncertainty Principle:
 You can know either the eˉ position or velocity
but not both
Schrodinger said the eˉ are located in orbitals,
(regions of probability) around the nucleus… not
orbits
“Electron Cloud” model