Mechanical Universe
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Lesson 49: The Atom
This program explores the history of the atom, from the ancient
Greeks to the early 20th century, when discoveries by J.J. Thomson
and Ernest Rutherford created a new crisis for the world of physics.
Text Assignment: Chapter 49
Instructional Objectives
Be able to summarize the kinetic theory and discuss the size of
atoms.
Be able to compare Thomson's model of an atom with Rutherford's
planetary model of an atom.
Be able to discuss why Rutherford's model of an atom conflicted
with Maxwell's theory of charged particles.
Be able to discuss the significance of Brownian motion in providing
evidence for the existence of atoms.
Bangs, Flashes, and Explosions
Section 4: Atomic Theory and Periodic Trends............................................ 43
Rutherford Experiment Simulation...........................................................................................43
Colored Flames ...............................................................................................................................45
Throwing Flame(test) Flames ....................................................................................................47
Bangs, Flashes, and Explosions – Illustrated Guide to Chemistry Demonstrations
©2005 Chris Schrempp and ExploScience.com. All Rights Reserved
Bengal Flames.................................................................................................................................48
Fiery Tornado...................................................................................................................................50
The Incredible Glowing Pickle (and grapefruit, and onion, and…) ..............................52
Reactions of Lithium, Sodium, and Potassium in Water..................................................54
John Dalton, British Physicist (1766-1844)
– Main contribution: All matter-whether element,
compound, or mixture-is composed of small particles
called atoms
Postulates of Dalton’s Atomic Theory
1.Matter consists of indivisible atoms.
2.All of the atoms of a given chemical element are identical in
mass and in all other properties.
3.Different chemical elements have different kinds of atoms; in
particular their atoms have different masses.
4.Atoms are indestructible and retain their identities in chemical
reactions.
5.A compound forms from its elements through the
combination of atoms of unlike elements in small wholenumber ratios.
Dalton’s Law of Definite Proportions (one compound always
having the same mass relationship)- Suppose a compound is made
of element H and element O, giving the formula H2O. Since the
weight of H is constant and the weight of O is constant, the H:O
weight ratio will always be the same. Hence, the Law of Definite
Proportions.
If H weighs 1 gram and O weighs 16g the mass relationship in
water (H2O)will always be 2g H/ 16 g O = 2:16 =1:8.
Dalton’s Law of Multiple Proportions (for 2 or more compounds
containing the same elements) - Two elements which form
more than one compound between them will have compounds
whose mass ratio between the two elements are related to each
other as a ratio of small whole numbers. H2O2 or H2O but not
H2.7O2.2
Example:
CO2
32.00 g O = 2.66 g O = 2.66
12 g of C
1.0g C
CO
16.00 g O = 1.33 g O = 1.33
12 g of C
1.0g C
The Structure of the Atom
Joseph John Thomson (1856-1940)
– British Physicist
– Main contribution: the discovery of the electron!
How the electron was discovered
Used a Cathode Ray tube
– 2 electrodes from a high voltage source are sealed in a
glass tube in which the air has been evacuated
– When electricity is applied, a lavender beam of light
emits from the electrodes
– Same lavender beam of light produced regardless of
electrode material
– Since same lavender light is emitted from all types of
material, all materials have same type of
particles…ELECTRONS!
– Plum Pudding Model
The Charge on an Electron
Robert Millikan (1868-1953)
–US Physicist
Main contribution: determined the charge on a single electron
Famous Oil Drop Experiment
Charge on a single electron
1.602 x 10-19 Coulombs
Along with JJ Thomson’s work, Millikan’s work helped
determine mass of the electron:
9.109 x 10-31 kg
Ernst Rutherford (1871-1937)
–British Physicist
Main contribution: developed a new model of the atom where
there is a small but massive, positively charged nucleus in the
atom.
The Nuclear Model of the atom
Rutherford’s model based on his famous gold foil experiment
–Used alpha particles as “bullets” (Helium atoms without
electrons, have + charge)
–Fired these particles into very thin gold foil
–Most went through without deflection
–Some particles were deflected at large angles
Rutherford’s Gold Foil Experiment
Explanation of Experiment
Most of the atom is empty space!
There must be very small, dense areas where the alpha particles
hit, are deflected, and bounce back
These areas must be positively charged!
Other Discoveries
These scattering experiments led to the determination of the +
charge on the nucleus called the Atomic Number, Z.
1911 – Rutherford discovered protons by striking lighter
elements, like nitrogen, with alpha particles
Led to new definition of element – a substance whose atoms
contain all the same atomic number (proton #)

The Discovery of the Neutron
James Chadwick (1891-1977)
–British Physicist
Main Contribution: discovered atom was made of neutral
particles that were as massive as protons - these he called
neutrons!
Isotopes
Some elements have 2 or more different versions of the same
atom!
These different versions all have the same number of protons but
differing numbers of neutrons; these are called isotopes!
NIELS BOHR



Depicted the atom as a small, positively charged nucleus
surrounded by waves of electrons in orbit — similar in
structure to the solar system, but with electrostatic forces
providing attraction, rather than gravity, and with waves
spread over entire orbits instead of localized planets.
The theory that electrons travel in discrete orbits around the
atom's nucleus, with the chemical properties of the element
being largely determined by the number of electrons in each
of the outer orbits.
The idea that an electron could drop from a higher-energy
orbit to a lower one, emitting a photon (light quantum) of
discrete energy (this became the basis for quantum theory).
DRAW THIS
# of Protons determines what element it will be
- a nuclide is a particular nucleus characterized by
a definite atomic number and mass number
Mass Number
Atomic Number, Z
One nuclide of sulfur One nuclide of phosphorus
Atomic Number = Number of Protons
Mass Number = Number of Protons + Neutrons
Example:
If you look up potassium (K) in the periodic table, it has an
atomic number of 19, meaning that all potassium atoms and ions
contain 19 protons.
Example:
A particle with 6 protons and an atomic mass number of 14 has 8
neutrons.
REVIEW Determine the number of electrons, protons and neutrons for:
13
5
58
B
58
28
108
Ni
22
Ni
19
9
28
11
7
54
Xe
64
Na
14
F
REVIEW
29
Cu
14
N
6
C
ISOTOPES
Isotopes are atoms of the same element that have a different
number of neutrons. Therefore, isotopes have the following
characteristics:
ISOTOPES
1.Isotopes have the same atomic number (same number of
protons), but a different atomic mass number (a different
number of neutrons).
2.Isotopes behave the same chemically, because they are
the same element. The only difference is that one is
heavier than the other, because of the additional neutrons
For example, carbon-12 and carbon-14 are both isotopes
of carbon. Carbon-12 has 6 neutrons; carbon-14 has 8
neutrons.
Isotope
32
S
33
S
34
S
36
S
Relative atomic mass Abundance (%)
31.9720707
95.02
32.9714585
0.75
33.9678668
4.21
35.9670809
0.02
The average weight of Sulfur is 32.066 a.m.u. and is the
accepted mass for Sulfur, (but not mass of all isotopes.)
LAB PARTNER EXERCISE ISOTOPES
1) Rubidium has two common isotopes, 85Rb and 87Rb. If
the abundance of 85Rb is 72.2% and the abundance of
87Rb is 27.8%, what is the average atomic mass of
rubidium?
85.56 amu
2) Uranium has three common isotopes. If the abundance
of 234U is 0.01%, the abundance of 235U is 0.71%, and the
abundance of 238U is 99.28%, what is the average atomic
mass of uranium?
237.98 amu
3) Titanium has five common isotopes: 46Ti (8.0%), 47Ti
(7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is the
average atomic mass of titanium?
47.92 amu
4) Explain why atoms have different isotopes. In other
words, how is it that helium can exist in three different
forms?
Neutrons exist to stabilize the nucleus – without them,
the nucleus would consist of nothing but positivelycharged protons in close proximity to one another.
Because there are different ways of
stabilizing the protons, there are different isotopes
ATOM
Subatomic Particles
Mass
Charge
Proton
1 a.m.u
+1
Neutron
1 a.m.u
0
Electron
0 a.m.u
-1
a.m.u = atomic mass unit – 1/12 of a carbon atom
How Many: (Mass # in a.m.u) Protons
1- Hydrogen: (1.008)
1
2 - Helium: (4.002)
2
3 – Lithium: (6.941)
3
4 – Beryllium: (9.012)
4
5 – Boron: (10.811)
5
6 – Carbon(12.017)
6
Neutrons
0
2
3
5
5
6
I MISS
CONVERSIONS

2H + O -> H2O
Two atoms of hydrogen and one atoms of oxygen react
together to make two molecules of water.
If the atom is so small, so how are we going to pick up and
mix 2 atoms of each reactant?
How do we slap together 2 trillion atoms of hydrogen (H)
and 1 trillion atoms of oxygen (O)?
Is the relative weight of these two huge masses any
different than the simple atoms? No
nting.'
1 Million atoms H = 1 atom H
1 Million atoms O = 1 atom O
THE MOLE
• The mole is the weight scale at the chemical fruit stand:
it allows different substances to be measured in a
comparable way
• Moles make it easier to interpret chemical equations.
Thus the equation: 2H + O = H2O
• is understood as "two moles of hydrogen plus one mole of
oxygen yields one moles of water."
RUN BACK TO
BIOLOGY
WHILE YOU
STILL HAVE
THE CHANCE!
And at this point God looked
down from heaven and said to
all confused chemists,
“LET THERE BE ZOOLOGY”
SO HOW MUCH IS THIS
MOLE?
IT EQUATES TO SAND IN THE SAHARA.
602,200,000,000,000,000,000 grains of sand
SO LETS BREAK IT DOWN A LITTLE
A MOLE LETS US KEEP OUR NUMBERS SIMPLE
WITHOUT LOSING TRACK OF ALL THOSE ATOMS.
The term dozen represents the number - 12 eggs
A deck of cards represents the number - 52 cards
The term ton represents the number – 2000 pounds
The term mole represents the number:
23
6.022 x 10
particles/mol
A MOLE OF ANYTHING EQUALS A MOLE OF
ANTHING ELSE (in # not weight)
Pass around vials of one mole of each substance.
• A dozen eggs will make a nice sized omelet, but a mole of
eggs will fill all of the oceans on earth more than
30 million times over.
• It would take 10 billion chickens laying 10 eggs per day
more than 10 billion years to lay that mole of eggs.
If I asked you how much does a dozen weigh ,
You would say that depends on what the dozen is.
The same goes for Moles: it simply depends on what it is.
All moles weigh differently (depends on the type of atom)
• So why would need to use such a
big number?
• Using absolute numbers of atoms would take a
MOLAR MASS – Mass of One Mole (g/mol)
ON WHAT IS 6.022 × 1023 BASED?
A mole is the amount of substance that is the same number of
atoms as atoms in 12 grams of Carbon, the number of carbon
defined as Avogadro's number.
• Moles are useful in chemical calculations, because they
enable the calculation of yields (and other values)when
confronted with different masses (weights) of elements.
• one oxygen atom weighs almost 16 times as much as a
hydrogen atom.
For example, helium has an atomic weight of 4.00 AMU.
Therefore, 4.00 grams of helium will contain one mole of
helium atoms.
You can also work with fractions (or multiples) of moles:
Not on lecture notes.
Mole/Weight Relationship Examples Using
Helium
Moles Helium
# Helium
Grams Helium
Atoms
1/4
1.505 x 1023
1g
1/2
3.01 x 1023
2g
1
6.02 x 1023
4g
2
1.204 x 1024
8g
10
6.02 x 1024
40 g
Molecular Weight (not atomic wt.)
• When atoms form molecules, the atoms bond together,
and the molecule's weight is the combined weight of all of
its parts.
• For example, every water molecule (H2O) has two atoms
of hydrogen and one atom of oxygen.
• One mole of water molecules will contain two moles of
hydrogen and one mole of oxygen.
• A bottle filled with exactly 18.02 g water will contain 6.02
x 1023 water molecules (not atoms).
• The concept of fractions and multiples described above
also applies to molecules: 9.01 g of water would contain
1/2 mole, or 3.01 x 1023 molecules.
• You can calculate the molecular weight of any compound
simply by summing the weights of atoms that make up
that compound.
Mass to Mole Conversions:
Grams to moles:
( ___ grams) x (1 mole / ___ grams )
Moles to grams:
(Number of moles) x (____grams / mole)
Moles to atoms: (maintain scientific notation and sig figs)
(Number of moles) x (6.022 x 1023 atoms / 1 mole)
1. How many moles in 125 g of Hydrogen?
_124_mol_
2. How many moles in 63.54 g of Copper?
_1.000 mol
3. How many moles in 34.7 g of Lithium?
5.00_ mol
4. How many moles in 84.04 g of Nitrogen?
_6.000 mol
5. How many moles in 95.7 g of Carbon?
_7.97_ mol
6. How many moles in 126.54 g of Hydrogen?
125.54_ mol
7. How many moles in 98 g of Chlorine?
__2.8__ mol
8. How many moles in 298 g of Potassium?
7.62_ mol
9. How many moles in 410 g of Sulfur?
__13 mol
10.
How many grams in 0.5 moles of Hydrogen?
_.5g
11.
How many grams in 1.5 moles of Magnesium?
36g
12.
How many grams in 5.70 moles of Silver?
615g_
13.
How many grams in 4.220 moles of Uranium?
1004g
14.
How many grams in .15 moles of Bromine?
__12g
15.
How many grams in 8.8 moles of Oxygen?
__140g
16.
How many grams in 9.2 moles of Argon?
__370g
17.
How many grams in 10.10 moles of Aluminum?272.5g
18.
How many grams in 101 moles of Helium?
19.
How many atoms in 0.12 mol of cadmium
7.2 x 1022
20.
How many atoms in 2.2mol of xenon
1.3 x 1024
21.
How many atoms in 8.0 mol of carbon
4.8 x 1024
22.
How many atoms in 5.5 mol of cadmium
3.3 x 1024
23.
How many atoms in .33 mol of magnesium 1.9 x 1023
24.
How many atoms in 1.59 mol of Americium 9.57 x 1023
25.
How many atoms in 0.120 mol of cadmium 7.23 x 1022
404g
A.
How many grams of carbon are needed to give 3 moles
of carbon?
B.
What is the molar weight of H2O2.?
C.
How many moles of sulfur dioxide, SO2 (g), are in
2000 grams of the gas.
D.
Determine the number of grams in 4 moles of H2O.
E.
Determine the number of molecules of H2O in 3 moles
H2O.
F.
Determine the number of moles of CO2 in 454 grams.
G.
Determine the mass in grams of 3.60 mol of H2SO4.
A.
How many grams of carbon are needed to give 3 moles of
carbon?
1. Calculate the molar mass for carbon. Look up the atomic
weight/mass in the periodic table. The molar mass for
carbon is 1 mole C = 12.0 grams
2. Determine the mass needed to provide 3 moles of hydrogen.
1 mole C = 12.0 grams
2 mole C = 24.0 grams
3 mole C = 36.0 grams
The practical way is to multiply the molar mass by the
number of moles.
B. What is the molar weight of H2O2?
The formula weight for H2O2 =
2(weight from hydrogen) + 2(weight from oxygen)
= formula weight for H2O2
2 H atoms x 1.008 amu + 2 O atom x 16.00 amu =
= 34.02 amu
The molar mass for H2O2 = 34.016 grams
IF I TELL YOU, “I NEED A MOLE OF HYDRODGEN
PEROIXIDE”
YOU WOULD BRING ME 6.022 × 1023 HYDROGEN
PEROXIDE ATOMS IN THE FORM OF 34.02g.
C. How many moles of sulfur dioxide, SO2 (g), are in 2000
grams of the gas?
1. Look up the atomic weights in the periodic table for
S and O.
The atomic weight for sulfur is 32.07 amu
The atomic weight for oxygen is 16.00 amu
2. Calculate the formula weight for SO2 . Add up the masses
from all the atoms in the formula
The formula weight for sulfur dioxide is:
32.07 amu S + 2 x (16.00 amu O) = 64.07 amu SO2
3. Determine the molar mass. Molar mass is a mass in grams
that is numerically the same as the formula weight.
1 mole SO2 = 64.07 grams SO2 = 64. grams SO2 (to the nearest
gram )
4. Convert 2000 grams of SO2 to moles. The "conversion
factor" is the molar mass.
(2000 grams SO2 )(1 mole SO2 /64. grams SO2 ) = 31.25 moles
SO2 = 31 moles SO2 (rounded to 2 sf)
D. Determine the number of grams in 4 moles of H2O?
Formula mass H2O = (2 x 1.0) + (1 x 16.0) = 18 .0
1 mole H2O = formula mass H2O = 18.0 grams H2O
4 moles H2O x .0 grams / 1 mole = 72.0 grams H2O
E) Determine the number of moles in 88 grams of CO2
Formula Mass CO2 = (1 x 12) + (2 x 16) = 44
1 mole of CO2 = formula mass CO2 = 44 grams CO2
88 grams CO2 x 1 mole CO2 / 44 grams CO2 = 2 moles CO2
E. Determine the number of molecules of H2O in 3 moles H2O
1 mole H2O = 6.023 X 1023 molecules H2O
3 moles H2O x 6.023 X 1023 molecules H2O / 1 mole H2O =
18.069 X 1023 molecules H2O = 1.8069 X 1024 molecules H2O
F. Determine the number of moles of CO2 in 454 grams?
The atomic mass of C is 12.01 and the atomic mass of O is
16.00.
The formula mass of CO2 is: 12.01 + 2(16.00) = 44.01 AMU
Thus, one mole of CO2 weights 44.01 grams. This relation
provides a conversion factor to go from grams to moles. Using
the factor 1 mol/44.01 g:
moles CO2 = 454 g x 1 mol/44.01 g = 10.3 moles CO2
G. Determine the mass in grams of 3.60 mol of H2SO4.
The atomic mass is 1.008 for H; 32.06 for S; 16.00 for O.
Formula mass H2SO4: 2(1.008) + 32.06 + 4(16.00)=98.08AMU
Thus, one mole of H2SO4 weights 98.08 grams.
Using the factor 98.08 g / 1 mol:
3.60 mol x 98.08 g / 1 mol = 353 g H2SO4
Problems Calculate the molar mass (g/mol)
1) HNO3
2) HC2H3O2
3) C12H22O11
4)
1) 63.0 g/mol HNO3
2) 60.0 g/mol HC2H3O2
3) 342 g/mol
Atoms
Monoatomic (He, Li, Ag, Hg)
7 Diatomic Molecules (O, N, H, F, Cl, Br, I)
BrIFN HONCl
Multiatomic Molecules(S8, P4)
Compounds CO, CH4, H2SO4
Oxygen 16.0 amu
Sulfur 32.1 amu
Sulfur atom is twice as heavy
Helium (4 amu) is one fourth as heavy as Oxygen (16 amu)
Phosphorous (31.0) Nitrogen (14.0)
P=2.21
Relative Weights
1 Atomic Mass Unit : 1/12 of an atom of Carbon
(12 atoms C = 12 amu)
Water Molecule 18 amu
(1.5 times heavier than C of 12 amu)
Sulfuric Acid Molecule 98.1 amu HSO4
(8.175 times heavier than C of 12 amu)
2H2
O2
2H2 O
+
+
2 *2.02g
32.00g
+
= 2*18.02g