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The Atom
Chapter 4
I. History of the Atomic Theory
IV. Mass of Atoms
A. Democritus
A. Atomic Mass
B. Aristotle
B. Average Atomic Mass
C. Lavoisier
C. Mole and Molar Mass
D. Proust
E. Dalton
F. Modern Atomic Theory
II. History of Atomic Structure
V. Radioactive Decay
A. Thomson
A. Process
B. Millikan
B. Comparison
C. Rutherford
C. Types
D. Bohr
E. Chadwick
F. Quantum Atom
III. Subatomic Particles
A. Comparing Particles
B. Atomic Number and Mass Number
C. Ions
D. Changing Number of Particles
E. Nuclear Notation
F. Hyphen Notation
I. History of the Atomic Theory
Remember: a scientific theory explains
behaviors and the ‘nature’ of things
 Theories can be revised when new
discoveries are made
 The theory describing the composition of
matter has been revised many times

A. Democritus (460-370 BC)
1.Matter is made up of “atoms” that
are solid, indivisible and
indestructible
2.Atoms constantly move in space
3.Different atoms have different size
and shape
4.Changes in matter result from
changes in the grouping of atoms
5. Properties of matter result from
size, shape and movement
B. Aristotle (384-322 BC )
1. Four kinds of matter
a. Fire – Earth – Water – Air
2. One kind of matter can transform
into another
3. Rejected idea of the “atom” (idea then
ignored for almost 2000 years
4. This theory was more popular and
it was easier to accept
Aristotle’s Theory of Matter
C. Antoine Lavoisier (1770s)
Experiment:
2 Sn + O2

2 SnO
tin
oxygen
tin (II) oxide
mass before reaction = mass after reaction
Law of Conservation of Mass
Matter cannot be created or destroyed (in a
chemical or physical change)
1788-1799
D. Joseph Proust (1779)
Develops Law of Definite Composition- all
samples of a specific substance contain the
same mass ratio of the same elements
a. ex: all samples of CO2 contains 27.3%
carbon and 72.7% oxygen
b. therefore ‘elements’ are combining
in a whole number ratio
D. John Dalton (1803)
Dalton became a school teacher at the
age of 12 (he left school at age 11)
Loved meteorology - pioneer in this field
Studied works of Democritus, Boyle and
Proust
Wrote New System of Chemical Philosophy in
1808
Develops Law of Multiple Proportions
a. describes the ratio of elements by mass in
two different compounds composed of the
same elements
Example:
carbon monoxide vs. carbon dioxide
1 part oxygen : 2 parts oxygen
*when compared to the same amount of
carbon in each compound
Dalton collects data and develops
Atomic Theory in 1803
a. Matter is made of small particles-atoms
b. Atoms of a given element are identical in size,
mass, but differ from those of other
elements*.
c. Atoms cannot be subdivided or destroyed*.
d. Atoms combine in small whole number ratios to
form compounds.
e. Atoms combine, separate, or rearrange in
chemical reactions.
* Modified in Modern Atomic Theory
F. Modern Atomic Theory (1)
1. All matter is made up of small particles
called atoms.
2. Atoms of the same element have the same
chemical properties while atoms of different
elements have different properties
3. Not all atoms of an element have the same
mass, but they all have a definite average
mass which is characteristic. (isotopes)
F. Modern Atomic Theory (2)
4.
5.
Atoms of different elements combine to
form compounds and each element in
the compound loses its characteristic
properties.
Atoms cannot be subdivided by chemical
or physical changes – only by nuclear
changes
II.
History of the Atomic Structure
A. J.J. Thomson (1887)
Experiments with cathode ray tubes
Voltage source
-
+
Vacuum tube
Metal Disks
Voltage source
-
+
Voltage source
-
+
Voltage source
-
+
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Voltage source

+
Passing an electric current makes a beam
appear to move from the negative to the
positive end
Voltage source

By adding an electric field
Voltage source
+
 By adding an electric field
Voltage source
+
 By adding an electric field
Voltage source
+
 By adding an electric field
Voltage source
+
 By adding an electric field
Voltage source
+
 By adding an electric field
Voltage source
+
 By adding an electric field he found that the
moving pieces were negative
Thomson’s Model of the Atom
a. electrons present (-)
 b. atom is like plum
pudding - bunch of
positive stuff (pudding),
with the electrons
suspended (plums)
 Calculated the ratio
between the charge of the
electron and its mass:

e/m
“Chocolate Chip Cookie” or “Plum Pudding Model”
Millikan’s Oil Drop Experiments
Robert Milikan (1909)
– Oil Drop Experiment
– Measured the electrical charge on the
electron
– Mass can be calculated (Thomson
determined the e/m ratio)
– Mass is 1/1840 the mass of a hydrogen
atom
– electron has a mass of 9.11 x 10-28 g
B. Millikan’s Oil Drop Experiment
(1909)
Oil-Drop Experiment: Animation

http://wps.prenhall.com/wps/media/object
s/602/616761/MillikanOilDropExperiment.h
tml

http://www.britannica.com/nobel/cap/omil
lik001a4.html

http://www.daedalon.com/oildrop.html
II.

History of the Atomic Structure –
Summary thus far
So, at this point we know:
- Atoms are indivisible particles
– Electrons are negatively charged
– The mass of an electron is very small
HOWEVER
– Atoms should have a (+) portion to balance
the negative part
- Electrons are so small that some other
particles must account for mass
C. Ernest Rutherford (1871-1937)
Experiment (1909)
Gold Foil Experiment (Expectations)
a. Shot alpha particles at atoms of gold
b. expected them to pass straight
through
Lead
block
Uranium
Florescent
Screen
Gold Foil
He thought this would happen:
According to Thomson Model
He thought the mass of the positive charge was
evenly distributed in the atom
Here is what he observed:
The positive region accounts for deflection
Gold Foil Experiment Results
a. Most positive alpha particles pass right
through
b. However, a few were deflected
c. Rutherford reasoned that the positive
alpha particle was deflected or repelled
by a concentration of positive charge
Gold Foil Experiment Conclusions
a. the atom is mostly empty space
b. the atom has a small, dense positive center
surrounded by electrons
II. History of the Atomic Structure

At this point in 1909, we know:
– p+ = 1.67 x 10-24 g
– e- = 9.11 x 10-28 g
– The charges balance!

But,
– How are the electrons arranged?
– There is still mass that is unaccounted for
Rutherford Model of the Atom
D. Niels Bohr (1913)
Electrons orbit nucleus in
predictable paths
E. Chadwick (1891 – 1974)
In 1935
1. Discovers neutron in
nucleus
2. Neutron is neutral - does
not have a charge n0
3. Mass is 1.67 x 10-24 g
slightly greater than the
mass of a proton
History of the Atomic Theory
1803
1897
1909
1913
1935
Today
solid
particle
electron
proton
e- orbit
nucleus
neutron
Quantum
Atom
theory
Dalton
Thomson
Rutherford
Bohr
Chadwick
Schrodinger
and others
II. History of the Atomic Structure
Charges balanced
 Mass accounted for
 However –
what about the
behavior of the
electrons?

F. The Quantum Atom Theory
1. The atom is mostly empty
space
2. Two regions:
a. Nucleus- protons and neutrons
b. Electron cloud- region where
you have a 90% chance of
finding an electron
III. Subatomic Particles
Comparing Particles
Relative Actual
mass (g)
Name Symbol Charge mass
Electron
e-
-1
Proton
p+
+1
1amu 1.67 x 10-24
Neutron
n0
0
1amu
0
9.11 x 10-28
1.67 x 10-24
III.
Subatomic Particles
B. Atomic Number and Mass Number
1. Atomic number
1. the number of protons in the nucleus
of an atom
a. identifies the element
b. no two elements have the same
atomic number
2. Ex. C is 6, N is 7 and O is 8
carbon
nitrogen
oxygen
B. Atomic Number and Mass Number
2. Mass number
a. the number of protons plus neutrons in the
nucleus of an atom
b. mass number is very close to the mass of an
atom in amu (atomic mass units)
c. two atoms with the same atomic number but
different mass number are called isotopes
1) (mass #) – (atomic #) = #n 0
C. Formation of Ions
1. Electrons and Ions
a. For neutral atoms, #e- = #p+
b. When there are more electrons than
protons, a negative ion forms (anion)
c. When there are less electrons than protons,
a positive ion forms (cation)
For now, we will work only with neutral atoms
C. Formation of Ions
Examples of Ions
Atom loses electrons Atom gain electrons
and form cations
and forms anions
Cations (+ ions)
Anions (- ions)
K+
BrCa2+
O2Al3+
N3-
C. Formation of Ions From Atoms
Na loses an electron
and forms a cation
Na0 – e- --> Na+
Cl gains an electron
and forms an anion
Cl0 + e- --> Cl-
D. Changing Number of Particles
1. You can never change the number of
protons and have the same element
2. If you change the number of neutrons in
an atom, you get an isotope
3. If you change the number of electrons in
an atom, you get an ion
E. Nuclear Notation
1. Nuclear Notation is one method for
depicting isotopes of an element
2. contains the symbol of the element,
the mass number, and the atomic
number
Mass
number
Atomic
number
X
Na
23
11
 How many protons?
 How many neutrons?
 How many electrons?
F. Hyphen Notation
1. Element symbol or name – mass #
2. EXAMPLES
a. Fluorine-19
b. C-14
c. U-238
IV.
Mass of Atoms
A. Atomic Mass
1. Mass of an atom
a. too small to measure in grams
b. use relative mass (amu)
1) atomic mass unit
2) 1 amu is defined as 1/12 the mass
of one C-12 atom
B. Average Atomic Mass
1. weighted average mass of all known
isotopes
a. weighted means that the frequency of
an isotope is considered
b. mass of each isotope is multiplied by
its percent occurrence in nature – then
product of all isotopes is added to get
the average atomic mass
Calculating Average Atomic Mass
Process
1. Mulitply % occurrence x mass of isotope
2. Add products for each isotope
isotope occurrence
isotope occurrence
Ex. X- 40 (30.0% )
X-30 (70.0%)
(40 x .300) + (30 x .700) =
12 +
21
= 33 amu
C. The Mole and Molar Mass
1. measures the amount of substance
a. 1 mole = 6.02x1023 (Avogdro’s #) of
particles (atoms, molecules, ions, electrons)
b. standard – 1mole is the number of atoms in
12g of C-12 isotope
2. Molar mass – mass in grams of one mole
(mol) of any substance
a. numerically equal to atomic mass in amu
b. unit is grams/mol
V.
Radioactive Decay
A. The Process
What is radioactive decay?
spontaneous nuclear change in which
unstable nuclei emit radiation and
lose energy
radiation – rays and particles
emitted by radioactive materials
Why do atoms undergo decay?
produce a nucleus that is more
stable- p/n ratio
B. Comparison of alpha, beta, gamma
Alpha
Form particle
Beta
particle
Gamma ____
electromagnetic
radiation
Symbol
Mass
Charge
Notation
4 amu
+2
no mass
-1
no mass
none
Alpha Particle
V.
1.
2.
Radioactive Decay
C. Types of Decay
Beta Decay (neutronproton + electron)
a. beta particle (electron) is given off
b. atomic number increases by one
c. mass number stays the same
Alpha Decay
a. alpha particle (2p++2n0) is given off
b. atomic number decreases by 2
c. mass number decreases by 4
V.
Radioactive Decay
D. Examples of Decay
Beta Decay (n0 p+ + e-) e- released
Parent Nuclei
Daughter Nuclei
Co-60 --------------> Ni-60 + e(z = 27)
(z = 28)
C-14 ---------------> N-14 + e(z = 6)
(z = 7)
Beta Decay
Beta Decay
Beta Decay
+ beta particle
+ beta particle
+ beta particle
+ beta particle
+ beta particle
Radioactive Decay
D. Examples of Decay
Alpha Decay (alpha particle (2n0 + 2p+)
released)
Parent Nuclei
Daughter Nuclei
Th-232 ----------------> Ra-228 + alpha
(z = 90)
(z = 88)
Ra-226 ---------------> Rn-222 + alpha
(z = 88)
(z = 86)
Alpha Decay
Alpha Decay
Alpha Decay
Alpha Decay
+ alpha particle
+ alpha particle
+ alpha particle
+ alpha particle
+ alpha particle
Radioactive Decay of Uranium
Radioactive Decay of Uranium
Radiation
Radiation
Radiation
Mole Calculations
Using Conversion Factors
1 mole
6.02 x 1023
molar mass
1 mole
PARTICLES <----> MOLES <----> MASS
6.02 x 1023
1 mole
1 mole
molar mass
Mole Conversions [mass-mole-atoms]
Type
Equality Used
1. MOLES  MASS
2. MASS  MOLES
1 mole= molar mass (g)
3. MOLES  ATOMS
4. ATOMS  MOLES
1 mole = 6.02 x 1023 atoms
5. MASS  ATOMS
6. ATOMS MASS
1 mole = 6.02 x 1023 atoms
1 mole = molar mass (g)
Mole Calculations
(Mass of Helium is 4.00 )
1. 1.00 mole of helium = 4.00g
2. 2.00 mole of helium = 8.00g
3. 1.00 mole of helium = 6.02 x 1023 atoms
4. 2.00 mole of helium = 1.20 x 1024 atoms
5. 16.0g of helium = 4.00 mol
6. 3.0l x 1023 atoms of helium = .500 moles
7. 8.00g of helium = 1.20 x 1024 atoms.