Name ____________________________________
Period ____________
Subatomic Particles & Models of the Atom
1) John Dalton – Combined all research on atoms up to that time and formed the atomic theory of matter
a) Daltons atomic theory of matter had 4 key points:
(I) All elements are made of atoms
(II) All atoms in an element are identical.
(III) Atoms are not created or destroyed in chemical reactions.
(IV) Compounds have definite whole number ratios of elements.
1) JJ Thompson – used a cathode ray to discover electrons and subsequently made the first model of the atom
(I) Cathode Ray Tube - An evacuated glass tube where a stream of electrons flows from the cathode
(negative electrode) to the anode (positive electrode).
(i) Thompson knew that the electrons must have come from the atoms
of the cathode because the air had been pumped out
(ii) Thompson made two big conclusions using his cathode ray
 Thompson concluded that electrons are negatively charged
because they are repelled by the negative side of a magnet and
attracted by the positive side of a magnet
 Thompson concluded that electrons must have mass because
when a paddle wheel is placed in the path of the cathode ray the
paddle wheel is pushed by the electrons
(iii) Thompson named the particles that the cathode ray was made out of
Sir Joseph John Thomson
electrons
(1856-1940)
 Electron – subatomic particle with a very small mass and
negative charge
 Since we are not constantly getting shocked by everything we touch Thompson deduced that an
atom must have these electrons embedded in a positively charged ball
b) Based on these finding Thompson proposed that an atom must looks like “plum pudding” with negative
charged electrons floating in a sphere of positive charge
c) Plum pudding model of an atom – JJ Thompson’s theory that an atom must looks like “plum pudding” with
negative charged electrons floating in a sphere of positive charge
This is J.J. Thompson’s plum pudding model of an atom. It
consists of electrons floating in a sphere of positive charge.
Not unlike plums float in plum pudding
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This is what plum pudding looks like for everyone not
friendly to 1890’s cuisine
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A cathode ray tube. Because the cathode ray is made out of negatively charged electrons it is attracted to the positively
charged portion of the magnet
J.J. Thompson’s Cathode Ray Experiment
When J. J. Thompson passed a cathode ray through 2 oppositely charged plates he found that it bent towards the
positively charged plate. From this he concluded that electrons must have a negative charge. Because the tube had all
the air pumped out of it Thompson also concluded that the electrons must have come from the atoms on the cathode
Before
After
Thompson also noted that the cathode ray has the ability to push a paddle wheel from one side of the tube to the other.
From this Thompson concluded that electrons must have a mass
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2) Ernest Rutherford
a) Wanting to investigate atoms closer Rutherford performed
the gold foil experiment
(I) Rutherford’s Gold foil experiment
(i) Rutherford focused a beam of alpha particles
through a piece of gold foil only 3 or 4 atoms thick
(II) Results of the gold foil experiment
(i) Rutherford found that most of the alpha particles
passed straight through the foil, however a few of
them where deflected
(III) Conclusions of Rutherford gold foil experiment
Ernest Rutherford
(i) Since most of the particles went through the gold
(1871-1937)
foil, Rutherford’s conclusion was that an atom must
be mostly empty space
(ii) Since some of the positively charged Alpha particles were deflected, Rutherford concluded that the
center of an atom must be small solid and positively charged
Rutherford Gold Foil Experiment
When Rutherford focused a ray of positively charges Alfa particles at a piece of gold foil he noticed that most of the
particles went through, while a few of them where deflected. From this data Rutherford concluded that the atom must
be mostly empty space, with a small positively charged nucleus.
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b) Based on these findings Rutherford corrected Thompson’s plum
pudding model of an atom saying that an atom must consist of a
small dense positively charged nucleus with electrons floating in
a cloud on the outside
Rutherford’s correction to Thompson’s
model of the atom. Rutherford
concluded that atoms must be mostly
empty space with electrons on the
outside and protons in the center.
3) Current model of an atom
a) Today, a simplified model of the atom is
called the solar system or planetary model of
an atom Since this model has electrons
rotating around a nucleolus like planets
around the sun
b) Solar system or planetary model of an atom
– electrons rotate around the nucleolus of an
atom at particular distances like planets
around the sun
Bohr’s solar system model said that electrons rotate at set
distances away from the nucleus
Numerically Describing atoms
A) What is an atom made out of?
1) A atom is made out of three smaller (subatomic)particles, protons, neutrons, and electrons
Mass (AMU)
Location
Charge
Symbol
Proton
1
nucleus
+
p+
Neutron
1
nucleus
0
e-
Electron
Effectively 0
(very small)
outside
-
n0
Subatomic Particle
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B) What is the difference between one atom and another?
1) In 1914 John Moseley, an English chemists discovered that the fundamental
difference between every different type if atom is the number of protons in its
nucleolus
a) Atomic number – the number of protons in an atoms nucleolus
2) Because we aren’t getting shocked by everything we touch Moseley further
concluded that most atoms are neutral, meaning that they have the same number
of protons and electrons
C) How do chemists describe atoms?
1) Chemists have come up with a number of ways to describe particular atoms
John Moseley
a) Atomic number – the number of protons in an atoms nucleolus
(1856 –1968)
b) Mass Number (atomic mass)- total number of protons and neutrons in the
nucleus of a particular atom.
(I) This number is expressed in AMU’s or atomic mass units
(II) 1 AMU is an extremely small unit for mass, it is only useful when describing 1 atom because it is so small
(III) an AMU is equal to the weight of 1/12 the weight of a carbon atom with a mass number of 12
(i) (in other words 1 proton or neutron is about 1 AMU)
D) Are there different types of atoms?
1) Atoms can have varying numbers of electrons and neutrons (the proton number is always the same because that
is what determines the identity of an atom)
a) Ions - Formed when an atom gains or looses electrons.
(I) Since the electron and proton number isn’t equal, ions have a charge
(i) The Charge of an ion is equal to the atomic number minus the number of electrons
Ion charge = (p+) – (e-)
Mg +2 = 2 less electrons than protons (2 more positives than negatives)
Atomic number
12 p+
(number of protons)
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10 e-
Number of electrons
Charge
Ion charge = (p+) – (e-)
12 p+ – 10 e- = +2
charge
N-3 = 3 more electrons than protons (3 more negatives than positives)
Atomic number
7 p+
(number of protons)
Number of electrons
10 eCharge
7 p+ – 10 e- = -3
Ion charge = (p+) – (e-)
charge
2) Isotopes - any of the different types of atoms of the same element, each having a different atomic mass
a) All isotopes of an element have the same number of protons (atomic number is what gives the atom its
identity) but the number of neutrons may vary
b) Chemists represent isotopes on paper using Isotopic Notation
(I) Isotopic Notation – A Shorthand way of representing an isotope of an element.
(i) There is two different forms of isotopic notation
Two isotopes of Helium. They both have the same number of protons but a different number of neutrons
Wright the isotope for chlorine with 17 p+, 17 e-, and 16n0
Isotopic notion 1)
Full name – mass number of isotope
I.E. = CHLORINE - 33
Isotopic notation 2)
Atomic symbol
MassNumber
AtomicNumber
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I.E. =
Atomic
number
6
6
Neutrons
number
6
7
6
8
Mass Number
33
17
Cl
Hyphen
Notation
Full Istopic
Notation
1. Why are Moles useful for atoms?
a. Atomic mass - mass of one atom of an element measured in AMU (atomic mass units)
i. H = 1 AMU
ii. O = 16AMU
iii. C = 12AMU
b. The number 6.02x1023 is useful in chemistry because 1 mole of AMU or 6.02x1023 AMU is 1
gram:
i. 1 H atom = 1 AMU
1mole of H atoms = 1g
ii. 1 O atom=16AMU
1mole of O atoms= 16g
iii. 1 C atom=12AMU
1mole of O atoms = 12g
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c. Therefore because a mole or 6.02x10 of AMUs 1 gram . . . a mole or 6.02x1023 atoms of any
element is equal to its atomic mass in grams:
i. 1mole of H atoms
is
6.02x1023 H atoms
weighs 1 gram
ii. 1mole of O atoms
is
6.02x1023 O atoms weighs 16 grams
iii. 1mole of C atoms
is
6.02x1023 C atoms
weighs 12 gram
2. Why are moles useful for molecules and formula units?
a. Formula mass - mass of all the atoms in a single molecule or formula unit of a compound.
measured in AMU (atomic mass units)
i. 1 H2O molecule = 18AMU
ii. 1 H2CO3 molecule = 62 AMU
iii. 1 Fe(NO3)2 formula unit = 180 AMU
b. But remember that 1 mole of AMU = 1gram so…
i. 1 H2O molecule = 18AMU
1 mole H2O molecules =18g
ii. 1 H2CO3 molecule = 62 AMU
1 mole H2CO3 molecules = 62g
iii. 1 Fe(NO3)2 formula unit = 180 AMU
1 mole Fe(NO3)2 formula unit = 180g
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c. Therefore because a mole or 6.02x10 of AMUs 1 gram . . . a mole or 6.02x1023 atoms of any
molecule, formula unit or compound is equal to its formula mass in grams:
i. 1mole of H2O molecules
is
6.02x1023 H2O molecules
weighs 18 grams
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ii. 1mole of H2CO3 molecules is
6.02x10 H2CO3 molecules weighs 62 grams
iii. 1mole of Fe(NO3)2
is
6.02x1023 Fe(NO3)2
weighs 180 grams
iv. Molar Mass – The mass of one mole of an element in g/mol.
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1. Round all molar masses to 1 decimal place
The Mole Practice:
1 mole
= 6.02x1023 particles
=
molar mass
1 mole Ne
= 6.02x1023 atoms Neon
=
20.2g
1 mole CO2
= 6.02x1023 molicules CO2
=
44g
1 mole CaF2
= 6.02x1023 formula units (Ionic)
=
78g
NaCl =54.43 g/mol
1 mole NaCl =
6.02x1023 molicules =
54.43 g/mol
MgO = 40.29 g/mol
1 mole MgO =
6.02x1023 molecules =
40.29 g/mol
N2 = 28.00 g/mol
2N2 = 56.00 g/mol
=
=
6.02x1023 molecules
1.204x1024 (mol x 2) molecules
Atomic mass of H2O = 18 AMU = Molar mass of H2O = 18g/mol = Mass of 1 mole H2O = 18g
E) Why are there decimals for masses on the periodic table of elements?
1) The “mass numbers” on the periodic table of elements are not decimals because those atoms have portions of
protons neutrons and electrons (that’s impossible)
2) That number is actually the Average Atomic Mass of that element
a) Average atomic mass - The combined average mass of all an element’s isotopes and how often those
isotopes occur.
(I) This average is a weighted average, which means that the isotopes that occur more often are count more
than those that hardly occur at all
Difference between straight average and weighted average:
There are thee known isotopes of hydrogen:
Hydrogen – 1 (accounts for 99.9844% of all hydrogen in existence)
Hydrogen – 2 (accounts for 0.0155% of all hydrogen in existence)
Hydrogen – 3 (accounts for 0.0001% of all hydrogen in existence)
This means that if a beaker contains 10,000 hydrogen atoms, 9,998 of them would be Hydrogen-1, 1 of them would be
Hydrogen -2 and one of them would be Hydrogen – 3.
What would be a better guess for the average mass of hydrogen atoms in the beaker:
Straight average:
Add up the masses and divide by the number of samples
(1 AMU + 2 AMU + 3 AMU)  3 = 2
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This means that the typical weight of 1 hydrogen atom in the jar is 2AMU’s
Weighted average:
The masses of each isotope multiplied by the amount it occurs all added together:
1 AMU (0.999844) = 0.999844 AMU
2 AMU (0.000155) = 0.00031 AMU
3 AMU (0.000001)= 0.000003 AMU
(0.999844AMU) + (0.00031AMU) + (0.000003AMU) = 1.0002AMU
This means that the typical weight of 1 hydrogen atom in the jar is 1AMU and very seldom there is a heavier isotope
Which number is more accurate for the average weight of hydrogen in the jar?
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