The Structure of the Atom
Chemistry
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Learning Objectives for this Chapter:
1. Describe changes in the atomic model over time
and why those changes were necessitated by
experimental evidence.
2. Describe structure of atoms (protons, neutrons,
electrons) and differentiate among these
particles in terms of mass, charges and location
in the atom.
3. Compare the magnitude and range of the 4
fundamental forces.
4. Apply the mole concept and the Law of
Conservation of Mass.
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It started a long time ago…
Today
460 – 370 BC
Democritus
Beginning of Atomism
You cannot divide something in half forever.
The smallest piece of matter is called an atom.
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Dalton’s Postulates
460 – 370 BC
1808
Democritus
Atomism
Today
1. All elements are made of tiny indivisible
particles called atoms.
2. All atoms of the same element are the
Over 2,000 years later
John Dalton comes up with
the first “modern” atomic
theory.
same, but different from atoms of every
other element.
3. Chemical reactions rearrange atoms but do
not create, destroy, or convert atoms from one
element to another.
4.
Compounds are made from combining
atoms in simple whole number ratios.
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Another kind of light?
460 – 370 BC
1808
Democritus
Atomism
Dalton
“Modern”
atomic theory
1870
Today
William Crookes invents a tube in which
virtually all the gas has been removed.
Under high voltage, a ray was emitted
from the cathode end of the tube.
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It started a long time ago…
460 – 370 BC
1808
Democritus
Atomism
Dalton
“Modern”
atomic theory
1870
Crookes
Cathode rays
Cathode rays must
be negative.
1897
Today
J.J. Thomson
discovers the
electron
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It started a long time ago…
460 – 370 BC
1808
Democritus
Atomism
Dalton
“Modern”
atomic theory
1870
Crookes
Cathode rays
1897 1910
Thomson
Discovery
of the electron
Today
Ernest Rutherford
discovers the
nucleus
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Size and mass
The size of the atom
comes mostly from
the space occupied
by the electrons
The mass of the
atom comes mostly
from the nucleus
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neutrons
electrons
protons
What happens when you change the number of protons?
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6 protons in
carbon
7 protons in
nitrogen
8 protons in
oxygen
You obtain a different element!
The number of protons is also called the
atomic number for that element.
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neutrons
electrons
protons
What happens when you change the number of electrons?
You get an ion – a charged particle.
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A neutral sodium atom
11 protons
Na
The protons and
electrons cancel each
other out
Balanced
charges
A positive sodium ion
1+
Na
One proton is not neutralized by
an electron, making this a +1
charged atom
One electron short
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8 protons
A negative oxygen ion
-2
O
Two electrons are not
neutralized by protons,
making this a –2 charged
atom
Two extra electrons
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The electron cloud
Electrons are very light and fast.
They are not organized along orbits around the nucleus.
Except for mass, virtually every property of atoms is determined
by electrons, including size and chemical bonding
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neutrons
electrons
protons
What happens when you change the number of neutrons?
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Atomic number
Neutrons act as “glue.”
They hold protons together in the nucleus.
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Mass number = 6 p + 6 n = 12
12C
“carbon-12”
mass number: total number of protons and neutrons in a
nucleus.
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Name
Mass
number
# protons
# neutrons
Carbon-12
Carbon-13
Carbon-14
12
13
14
6
6
6
6
7
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isotopes: atoms or elements that have the same number of
protons in the nucleus but different number of neutrons
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Atomic particle determination
a. Neutrons = mass number of nucleotide –
atomic number of element
b. Protons = atomic number (Z) of element
c. Electrons = atomic number of element
d. Atomic Mass – average mass of isotopes of
an element
e. Mass number – the total protons and
neutrons in the nucleus of an atom
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V Atomic Mass Formulae
1. Mole – the amount of a substance that contains the
same number of units as the number of atoms in
exactly 12 g of Carbon – 16. This is Avogadro’s number
of 6.022 x 1023 particles.
2. Mass to number of atoms =
given mass in g of element
6.022 x 1023 atoms
Molar mass in g of element
3. Number of atoms to mole:
# atoms
mole
6.022 x 1023 atoms
4. Mass to moles:
given mass in g
moles
Molar mass in g
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Number of neutrons for each of 100 lithium atoms
randomly sampled from nature
It’s an AVERAGE mass!
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Isotope periodic table (first 4 rows)
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Forces found in Nature:
1. Gravitational – relatively very weak but
active over long distances
2. Electromagnetic – intermediate strength
3. Strong Nuclear – strongest force that
binds protons and neutrons together in the
nucleus
4. Weak Nuclear – involved in the decay of
many elementary particles
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Examples of the Weak Nuclear force.
Some atoms are unstable and, to become stable,
decay by emitting particles or energy or both.
Alpha particles, a
Beta, b
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Nuclear Particles
Proton
Mass: 1 amu
Charge: +1
Neutron
Mass: 1 amu
Charge: 0
(1 amu = 1/12 mass of Carbon 12 atom)
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Electrons
Negative
0
e
-1
Positive
0
e
positron
+1
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Uranium Isotopes
238
235
U
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U
92
Low Radioactive Fission
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Nuclear Reactions
• Nuclear Decay - a spontaneous process in which an
unstable nucleus ejects a particle and changes to
another nucleus.
– Alpha decay
– Beta decay
• Beta Minus
• Positron
• Fission - a nucleus splits into two fragments of
roughly equal size
• Fusion - Two nuclei combine to form a heavier
nucleus.
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Alpha Decay
• This occurs when a helium nucleus is released.
• This occurs only with very heavy elements.
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Beta (b-) Decay
• A beta particle (negative electron) is released
when a nucleus has too many neutrons for the
protons present. A neutron converts to a
proton and electron leaving a greater number
of protons.
• Neutron decay:
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Positron (b+) Decay
• Positron decay occurs when a nucleus has too
many protons for the neutrons present. A
proton converts to a neutron. A neutrino is
also released.
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Fission
Fusion
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Learning Objectives for this Chapter: Atomic
Structure
1. Describe changes in the atomic model over time
and why those changes were necessitated by
experimental evidence.
2. Describe structure of atoms (protons, neutrons,
electrons) and differentiate among these
particles in terms of mass, charges and location
in the atom.
3. Compare the magnitude and range of the 4
fundamental forces.
4. Apply the mole concept and the Law of
Conservation of Mass to make calculations.
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