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Unit 3 – Atomic Structure
Bravo – 15,000 kilotons
DO NOW
• How many moles are in 78.54 g Magnesium?
(1 mole = 24.31 g Mg)
• How many atoms are in 45.7 g of sodium?
(1 mole = 6.02 x 1023 atoms;
1 mole = 22.99 g Na)
Agenda
• DO NOW
• Atomic Structure History
Democritus
• 400 BC
• Greek philosopher
• 1st to come up with idea of atoms
John Dalton – 1800’s
•
1)
2)
3)
4)
Major contributor of Atomic Theory
All matter made of atoms
All atoms of an element are alike
Atoms cannot be created or destroyed
Atoms combine in whole-number ratios
to form compounds
JJ Thomson – late 1800’s
• Cathode Ray Experiment – discovery
of electrons
• “Plum Pudding” model of atom
• Measured charge to mass ratio of
e-
Thomson’s Atomic Model
Thomson believed that the electrons were like plums
embedded in a positively charged “pudding,” thus it was
called the “plum pudding” model.
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube
to deduce the presence of a negatively charged
particle.
Cathode ray tubes pass electricity through a gas
that is contained at a very low pressure.
Conclusions from the Study of
the Electron
 Cathode rays have identical properties regardless
of the element used to produce them. All elements
must contain identically charged electrons.
Atoms are neutral, so there must be positive
particles in the atom to balance the negative
charge of the electrons
 Electrons have so little mass that atoms must
contain other particles that account for most of
the mass
Millikan - 1910
Oil Drop Experiment
Determined actual charge and mass
of an e-
Rutherford - 1910
 Discovered nucleus and that it was positive
 Gold Foil Experiment
1) Most of atom is empty space (majority of
particles went straight through)
2) nucleus is small, dense and positively
charged (some positive charges were greatly
deflected)
Rutherford’s Gold Foil Experiment
 Alpha particles are helium nuclei
 Particles were fired at a thin sheet of gold foil
 Particle hits on the detecting screen (film) are
recorded
Rutherford’s Findings
 Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected
Conclusions:
 The nucleus is small
 The nucleus is dense
 The nucleus is positively charged
Niels Bohr - 1915
• Proposed early model of atom
• “Planetary Model” electrons orbit nucleus like
planets orbit sun
• Lacks math of modern version
• Has some errors/violates current theory
• Radiation is emitted when electrons move from
one orbit to another
Chadwick
Discovered neutrons
Modern Atomic Theory
(changes from Dalton)
 Atoms cannot be subdivided, created, or
destroyed in ordinary chemical reactions.
However, these changes CAN occur in nuclear
reactions!
 Atoms of an element have a characteristic
average mass which is unique to that element
(isotopes)
Foundations of Atomic Theory
Law of Conservation of Mass: mass is neither created or
destroyed in ordinary chemical reactions
Law of Definite Proportions (composition): compounds
contain same elements in same ratio by mass
Example: NaCl is always 39.9% Na and 60.66% Cl by
mass
Law of Multiple Proportions: 2 or more different compounds
composed of same two elements have ratios of small
whole numbers
Example: CO vs CO2 ratio of oxygen to oxygen is 2 to 1
Book Assignment
• Use sections 4.1 and 4.2 of the textbook to
answer the following questions:
– P.110 #1,4,5
– P.115 #10, 13-15
What is AMU?
Stands for atomic mass unit – used when
describing “relative” atomic masses
This system is used because the actual masses of
atoms are so small
Carbon-12 is the standard to which all other
elements are compared (i.e. hydrogen-1 has a
mass that is 1/12 that of carbon-12 so it’s
mass would be 1 amu)
Atomic Number
Atomic number (Z) of an element is the
number of protons in the nucleus of each atom
of that element.
Element
# of protons
Atomic # (Z)
6
6
Phosphorus
15
15
Gold
79
79
Carbon
Mass Number
Mass number is the number of protons and
neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide
p+
n0
e-
Mass #
Oxygen - 18
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Atomic Masses
Atomic mass is the average of all the naturally
isotopes of that element.
Carbon = 12.011
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.11%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
Isotopes
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
ISOTOPES
• Example:
• Carbon, C, exists in 3 isotopes:
Isotope
Protons
Neutrons
Mass
Carbon-12
6
6
12
Carbon-13
6
7
13
Carbon-14
6
8
14
Isotope symbols
Hyphen notation Nuclear notation
Carbon – 12
12C
6
6p+ and 6no
Carbon – 13
13C
6p+ and 7no
Carbon – 14
14C
6p+ and 8no
6
6
How to Calculate the Average Mass
What is the average atomic mass of sample of
Cesium with 3 isotopes:
75% 133Cs, 20%
0.75 x 133 =
0.20 x 132 =
0.05 x 134 =
Total
=
132Cs,
and
5%
134Cs.
99.75
26.40
6.70
132.85
avg. atomic mass
The Mole
1 dozen = 12
1 gross = 144
1 ream = 500
1 mole = 6.02 x 1023
There are exactly 12 grams of
carbon-12 in one mole of carbon-12.
Avogadro’s Number
6.02 x 1023 is called “Avogadro’s Number” in
honor of the Italian chemist Amadeo Avogadro
(1776-1855).
I didn’t discover it. Its
just named after me!
Amadeo Avogadro
Calculations with Moles:
Converting moles to grams
How many grams of lithium are in 3.50 moles of
lithium?
3.50 mol Li
6.94 g Li
1 mol Li
=
45.1
g Li
Calculations with Moles:
Converting grams to moles
How many moles of lithium are in 18.2 grams of
lithium?
18.2 g Li
1 mol Li
6.94 g Li
=
2.62
mol Li
Calculations with Moles:
Using Avogadro’s Number
How many atoms of lithium are in 3.50 moles of
lithium?
3.50 mol Li 6.022 x 1023 atoms Li
1 mol Li
= 2.11 x 1024 atoms Li
Calculations with Moles:
Using Avogadro’s Number
How many atoms of lithium are in 18.2 g of
lithium?
18.2 g Li 1 mol Li
6.94 g Li
6.022 x 1023 atoms Li
1 mol Li
(18.2)(6.022 x 1023)/6.94
= 1.58 x 1024 atoms Li