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Chapter 5
Early Atomic Models

Atoms: the smallest particle of an element
that retains the properties of that element.
 (Greek: atomos = indivisible)

Democritus (Greek teacher in the 4th
century BC)
 First suggested the idea that atoms
existed
The Atom
1700’s – chemists were able to relate
changes to individual atoms
 Average atom size:
 Mass = 1 x 10 –23 g
 Diameter = 1 x 10-8 cm
 How small is that?100,000,000 copper
atoms in a row would = 1 cm in length!

Law of Conservation of Mass
 Definition: mass cannot be created or destroyed
only transformed
Law of Definite Proportions
 Definition: a chemical compound contains the
same elements in exactly the same proportions
by mass regardless of the size of the sample or
source
 Example:
▪ Sodium chloride: NaCl always consists of exactly
39.34% sodium & 60.66%chlorine by mass
▪ Water: H2O always consists of exactly 11.18%
hydrogen & 88.82% oxygen by mass
Law of Multiple Proportions
 Definition: if two or more different compounds
are composed of the same two elements, then
the ratio of the masses of the second element
combined with a certain mass of the first
element is always a ratio of small whole
numbers
 Examples:
 CO & CO2 : 1:1 ratio & a 1:2 ratio
 H2O & H2O2: 2:1 ratio & a 2:2 ratio
John Dalton
 English school teacher
 Proposed an explanation for the 3 laws
 Established in 1808
Dalton’s Atomic 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.
3. Atoms of different elements can combine
with one another in simple whole number
ratios to form compounds.
 H2O
C12H22O11
NOT
H2.5O¾
4. Chemical reactions occur when atoms are
separated, joined or rearranged. Atoms of
one element are not changed into atoms of
another by chemical means!
5. Atoms can not be subdivided.
The Structure of the Atom
Most of Dalton’s Atomic Theory is accepted
 One major revision includes that idea that
atoms are indivisible….
There are 3 parts to an atom….
1. electrons
2. protons
3. neutrons
Discovery of the Electron
Negatively charged subatomic particles
 J.J. Thomson discovered in 1897
 Passed a electric current through gases at low pressures
called a “Cathode Ray Tube”
 Noticed the surface of the tube directly opposite the
cathode glowed.
 Why? Opposites attract and the electrons were attracted
to the positive ends and lights up!
 Cathode Ray Tube
Cathode Ray Tube
Cathode rays are identical regardless of the
element
 Therefore all elements must have electrons!
Other important findings:
 Atoms are electrically neutral, so they must
contain a positive charge to cancel it out
 Since electrons are so small, atoms must contain
other particles that account for their mass
Robert Millikan (1868-1953)
 Found quantity of charge in 1
electron (e-)
 Also determined the ratio of the
charge to the mass of 1 e Calculated the mass of 1 eElectrons weigh 9.109 x10-31 kg
J.J. Thomson – plum-pudding model
 e- are spread evenly though out the positive
charge of the rest of the atom
 Ms. Agostine’s “mint chocolate chip ice
cream model”
Ernest Rutherford (1911)
 nucleus of the atom is positively
charged
 Gold Foil Experiment

Most  particles go straight through

Positively charged  particles deflect off of the
positively charged nucleus(~1/8,000)

Gold Foil Experiment

“…it was as if you fired a 15-inch [artillery]
shell at a piece of tissue paper and it came
back and hit you.”

Nucleus was very small

If a nucleus were a marble
 the atom would be a football field




Protons (p+)
Positively charged
particles
Mass = 1.673 x 10-27
kg
1,836 times heavier
then an electron
Neutrons (no)
Subatomic particles
with no charge
 Discovered by Sir
James Chadwick
 Mass is nearly the
same as a proton
 Mass = 1.675x10-27 kg


Particle Symbol Relative
Charge
Electron
e-
1-
Relative
Mass
(amu)
1/1836
Actual
Mass (kg)
Proton
p+
1+
1
1.67x10-27
neutron
no
0
1
1.68x10-27
9.11x10-31
Atomic Number : the number of protons in the
nucleus of an atom of an element
Atoms are electrically neutral
 Tells how many electrons there are also!
Periodic Table
#1 – Hydrogen: has 1 p+ and 1 e#6 – Carbon: has 6 p+ and 6 e-
Mass Number – total number of protons and
neutrons in a nucleus
# of neutrons = mass #
- atomic #
= (# p+ + # no) - (# p+)
Ex) Beryllium – 9
 Hyphen notation: The number “9” is the
mass number
# of p+?
# of no?
# of e-?
Definition – atoms that have the same number
of protons but different numbers of neutrons
 Different types of the same element
 Ex) Carbon – has 3 isotopes
 Carbon – 12
 Carbon – 13
 Carbon – 14
 Differ by # of no
All have the same # of p+
 If not, it would be a different element
 All have 6 protons
Carbon – 12
 Has 6 neutrons
Carbon – 13
 Has 7 neutrons
Carbon – 14
 Has 8 neutrons
Hydrogen-1: 1 p+ and 0 no
 Relative abundance = 99.985 %
 Commonly called normal “hydrogen”
Hydrogen-2: 1 p+ and 1 no
 Relative abundance = 0.015%
 Commonly called heavy hydrogen or “deuterium”
Hydrogen-3: 1 p+ and 2 no
 Relative abundance = ~0.00%
 Commonly called “tritium”
Definition – weighted average mass of the
atoms in a naturally occurring sample of the
element
Carbon-12 = 98.89 % abundant
Carbon-13 = 1.11% abundant
Carbon-14 = ~0.0000001% abundant
Formula:
Atomic = relative • mass # + relative • mass #
mass
abund.
#
abund.
+
Repeats for as many isotopes as exist for that element….
Units: atomic mass unit (amu): defined as
exactly 1/12 the mass of a carbon-12 atom
 1 amu = approximately the mass of 1 proton
 amu’s are used so you don’t have to use scientific notation
when talking about such small masses
Sample Problem:
 Chlorine has 2 isotopes:
chlorine-35 which is 75.77%
abundant and chlorine-37
which is 24.33% abundant.
What is the atomic mass of
chlorine?
35 Cl
= 75.77% abundant = 0.7577 rel. abund.
37 Cl = 24.33% abundant = 0.2433 rel. abund.
Atomic mass =
= (35 amu x 0.7577) + (37 amu x 0.2433)
= (26.5195 amu)
+ (9.0021 amu)
= 35.5 amu
Compare to value on Periodic Table = 35.45 amu
which rounds to 35.5 amu