Download Atomic Structure

Unit 3
History of the Atom
400 B.C. – Democritus & Leucippus
 Beach  sand  smaller piece of
sand  atomos (indivisible)
 Everything is composed of
imperishable, indivisible elements
called atomos.
Aristotle’s views
All substances are made of 4 elements
 Fire, Air, Earth, Water
Blend these elements in different
proportions to get all substances
So who was right?
Greeks settled argument by…
Debating!
Aristotle was a better speaker so he
won.
His views carried on through the
Middle Ages.
1808 - Dalton
English schoolteacher
Recognized that elements were
made of atoms
Combined LOTS of research
Dalton’s Atomic Theory
Dalton’s Atomic Theory
1. All matter is made of tiny, indivisible
particles called atoms.
2. Atoms of a given element are
identical in their physical and
chemical properties.
3. Atoms of different elements differ in
their physical & chemical
properties
Dalton’s Atomic Theory
4. Atoms of different elements
combine in simple whole # ratios to
form compounds
5. In a chemical reaction, atoms are
combined, separated, & rearranged,
but never created, destroyed, or
changed.
From Dalton’s Theory…
Law of Definite Proportions (#4)
 2 samples of a compound have the
same proportions by mass
 500 kg NaCl = 60.66% Cl & 39.34% Na
 2 mg NaCl = 60.66% Cl & 39.34% Na
Law of Conservation of Mass (#5)
 Mass of reactants = mass of products
From Dalton’s Theory…
Law of Multiple Proportions
 If 2 or more compounds are composed
of the same elements, the ratio of the
mass of the elements is always a small,
whole #
 NO 1.14 g O: 1 g N
 NO2 2.28 g O: 1 g N
 O [NO]: O [NO2] = 1.14 : 2.28 = 1 : 2
Another way to look at it…
 Water (H2O) has 8 g of oxygen per g of
hydrogen.
 Hydrogen peroxide (H2O2) is 16 g of
oxygen per g of hydrogen.
 16 to 8 is a 2 to 1 ratio of oxygen.
 Always whole #s because you have to add
a whole atom --- you can’t add a piece of
an atom.
Dalton’s Atomic Model
The Solid Sphere Model
The research continues…
1897 – JJ Thomson – cathode ray tube
Movie
Pumped air out of glass tube and
applied voltage to metal electrodes at
either end of the tube
 Anode = positive charge
 Cathode = negative charge
Thomson’s Experiment
Voltage source
-
+
Metal Disks
Thomson’s Experiment
Voltage source
 Passing
+
an electric current makes a
beam appear to move from the
negative to the positive end
Thomson’s Experiment
Voltage source
+
 By adding an electric field he found that
the moving pieces were negative
 Also placed a paddle wheel in the center –
and it turned! So they must have mass!
Thomson’s Conclusions
The cathode ray consists of particles
that have mass & a negative charge
 Particles called ELECTRONS
Plum pudding model
 Negative electrons in a
ball of positive charge
The research continues…
We know that an electron has a
negative charge and atoms are
neutral
Mass of an electron MUCH less
than an atom
Something must be missing…
Rutherford - 1909
Beam of small, positively charged
particles called ALPHA particles
Aimed at thin gold foil (a few
atoms thick)
Measured the angles of deflection
Movie
Lead
block
Uranium
Flourescent
Screen
Gold Foil
What Rutherford Expected
 The Plum Pudding Model
 Alpha particles would pass through
without changing direction very much.
 Because…
 The positive charges were spread out
evenly & would not stop the positive
alpha particles.
What he expected
Because
Because, he thought the positive
charges were evenly distributed in
the atom
Because, he thought
the positive charges
were evenly
distributed in the
atom
What he got
Rutherford’s Explanation
 Atom is mostly empty.
 Small dense,
positive piece
at center.
 Alpha particles
are deflected by
it if they get hit the
dense positive center.
+
+
Conclusions from Rutherford
 The small dense, positive place at the
center is called the NUCLEUS
 Radius of nucleus is less than 1/10000 the
radius of the atom
 PROTON – positive charged particle in
nucleus
 Charge is exactly equal but opposite to an
electron
 BUT still not enough mass
Still More Research…
NEUTRONS
 Found in nucleus with protons
 Do not have a charge
 Same mass as protons
The Nuclear Model - Rutherford
Electrons revolve around nucleus
in elliptical orbits
Also called planetary
model
The Bohr Model
Electrons are found in certain
levels or shells around the nucleus
More research…
1924 – de Broglie - Electrons behave
like waves around the nucleus
Heisenberg’s Uncertainty Principle
 A ceiling fan
Planck & Einstein – quantum theory
Electrons found in clouds instead of
strict orbitals
Today’s model – Quantum model
Historical models of the atom (solid sphere, plum pudding,
nuclear model, & planetary model all predicted exact locations
of particles. The modern quantum theories combine all of this
research with more recent findings that suggest a certain level
of unpredictability.
Subatomic particles
Name
Relative
Symbol Charge mass
Actual
mass (g)
Electron
e-
-1
1/1840 9.11 x 10-28
Proton
p+
+1
1
1.67 x 10-24
Neutron
n0
0
1
1.67 x 10-24
Most of the mass is in the nucleus!
The Atom
All atoms have protons and electrons
Most atoms have neutrons
Elements differ from each other in
the number of protons in an atom
Protons & neutrons made of quarks
Atomic Number
The number of protons in an atom
Same in all atoms of an element
Example – Hydrogen=1
Atomic number also reveals the
number of electrons in a neutral atom
Mass Number
Number of particles in the nucleus
 = # of Protons + # of neutrons
Example – Neon has a mass # of 20
Can vary among atoms of an element
Different elements can have the same
mass number
Not specifically on periodic table
Using Atomic Symbols
Each element has a name & a symbol
Examples –
 Sulfur = S
 Sodium = Na
The subscript to the right tells you
how many atoms are present
 S8 = 8 sulfur atoms
Using Atomic Symbols
 Contain the symbol of the element, the
mass number, and the atomic number.
Mass
number
Atomic
number
X
Symbols/Notation
 Find the …
 Atomic number
 Mass Number
 number of protons
 number of neutrons
 number of electrons
 Name
24
11
Na
Symbols/Notation
 Find
the …
–Atomic number
–Mass Number
–number of protons
–number of neutrons
–number of electrons
–Name
80
35
Br
Isotopes
All atoms of an element have the
same number of protons but not
necessarily the same # of neutrons
ISOTOPE – Atoms of the same
element with different numbers of
neutrons
Naming Isotopes
Name – mass number
Examples
 Helium-3, Helium-4
Symbols/Notation
3
2
4
He
2
He
Name That Element
 if
an element has an atomic
number of 34 and a mass number
of 78 what is the
–number of protons
–number of neutrons
–number of electrons
–Complete symbol
– Name
Name That Element
 if
an element has 91 protons and
140 neutrons what is the
–Atomic number
–Mass number
–number of electrons
–Complete symbol
– Name
Name That Element
 if
an element has 78 electrons and
117 neutrons what is the
–Atomic number
–Mass number
–number of protons
–Complete symbol
–Name
Atomic Mass
There are different isotopes of each
element that all have different masses
Therefore we look at AVERAGE
atomic mass for an element
Based on abundance of each isotope
in nature
Atomic Mass Example
You have five rocks -- four with a
mass of 50 g, and one with a mass
of 60 g. What is the average mass of
the rocks?
(50 + 50 + 50 + 50 + 60) / 5
260/5 = 52 g
Atomic Mass
Atoms are so small that even
picograms aren’t useful
1 g = 10-12 pg
So… we use a different unit
Atomic Mass Units (amu)
AKA a Dalton (Da)
Defined as 1/12 the mass of a carbon-
12 atom.
Very close to mass number
Example – Oxygen (O)
 Mass number = 16
 (Average) atomic mass = 15.999 amu
Calculating averages
Average atomic mass = (% as decimal
x mass) + (% as decimal x mass)
Calculate the atomic mass of copper if
copper has two isotopes. 69.1% has a
mass of 62.93 amu and the rest has a
mass of 64.93 amu.
Atomic Mass of Compounds
Add the masses of the parts!
EXAMPLE
 H2O = 2 hydrogen atoms + 1 oxygen
 H2O = 2 (1.0079) + 15.999
 H2O = 18.015 amu in 1 molecule
Another number problem…
Most samples of elements have
LOTS of atoms
We want things to be simple!
Scientists created a new unit
The MOLE
 # of atoms in exactly 12 g of Carbon-12
Avogadro’s number = the number
of particles in one mole of an
element
Avogadro’s number = 6.022 x 1023
particles (atoms or molecules)
Avogadro as a conversion factor
0.30 mol F = _________ atoms F
0.30 mol F
6.022 x 1023 atoms F
1
mol F
Molar Mass
Mass in grams of 1 mole of an element
Units are g/mol
Molar mass of 1 mole = atomic mass
Copper (Cu) atomic mass = 63.55 amu
Therefore molar mass = 63.55 g/mol
Molar mass as a conversion factor
3.50 mol Cu = _________ g Cu
3.50 mol Cu
63.55 g Cu
1
mol Cu
Molar mass as a conversion factor
 2 step conversions – 3g Cu = _________
atoms
 Calculating compounds – molar mass
of H2O?
 2 step conversions with compounds –
 4g H2O = _____ atoms
Atomic Nuclei Review
Nucleus made of protons & neutrons
Isotope – nucleus with same # of
protons & different # neutrons
Why don’t nuclei repel each other?
1935 Yukawa – attractive force
between neutrons & protons
stronger than proton repulsion
Only works when particles very
close together
Not as strong in large nucleus
Spontaneous Nuclear Change
AKA transmutation
All to increase stability
1. Fission – large nucleus splits
2. Fusion – small nuclei combine
3. Radioactive decay – release of
particles
Radioactive Decay
1. Convert neutrons to protons
 Emission of high energy negative
particles called BETA PARTICLES
 If a neutron loses its negative
charge, it becomes a proton
 Increases atomic number, no
change to mass number
Radioactive Decay
1. Convert neutrons to protons
Example
234
90
Th
234
Pa
91
+
0
-1
β
Radioactive Decay
2. Convert protons to neutrons
Nucleus captures an electron
If a proton combines with an
electron, it forms a neutron
Atomic number decreases by 1, no
change in mass number
238
92
U
238
91
Pa
Radioactive Decay
3. Losing ALPHA PARTICLES
An alpha particle is a positive
particle identical to a helium-4
4
nucleus ( 2 He)
Example
238
92
U
234
90
4
Th + He
2
Radioactive Decay
Also capable of producing a high
energy photon called GAMMA RAY
EMISSION
Does not change nucleus
Types of Radiation
ALPHA
BETA
GAMMA
Symbol
How change
the nucleus
Penetration
Low
Medium
High
Shielding
provided by
Skin
Paper,
clothing
Lead
Danger
Low
Medium
High
Nuclear reactors
 Uranium atoms split (fission) and
releases neutrons
 Neutrons trigger more uranium to split
 nuclear chain reaction
 Boron absorbs neutrons so is used to
control the speed of the chain reaction
 Fission products like iodine & cesium
also produce alpha, beta, & gamma
particles to stabilize
Nuclear reactors
 Heat must be removed to allow decay
to continue (decay = stabilization) so
cooling systems are vital
 Fukushima, Japan
 Power failed from earthquake
 Diesel generators failed from tsunami
 Release of radioactive isotopes of cesium &
iodide
Half Life
Rate of decay of a radioactive sample
Constant – not influenced by
temperature, pressure, etc.
Often used to determine age –
RADIOACTIVE DATING
Half Life
Most Carbon on Earth is C-12
All animals have C-14 from plants
All living plants & animals have fixed
ratio of C-14 to C-12
Once dead, C-14 levels decrease
Measure ratio & compare to known
Half-Life Example
C-14 half life = 5715 years
You find an artifact with a C-14:C-12
ratio that is 1/8 the modern ratio.
How old is the artifact?
Half-Life Example
1. Determine the number of half-
lives that the artifact has
undergone.
2. Multiply the # of half-lives by the
length of the half-life
Half-Life Example
Ratio is one eighth the modern ratio
1  ½
½  ¼ ¼  1/8
 3 half lives
5715 years x 3 half lives
 =17,145 years old
More Half-Life Examples
The half-life of Ra-226 is 1599 years.
How many years are needed for the
decay of 15/16 a given amount?
Assume the half-life of a substance
is 3.824 days. How much time will
it take for ¼ of a sample to remain?
More Half-Life Examples
Assume the half-life of a substance
is 3.0 minutes. How long will it
take for 16 mg to become 1.0 mg?