Download Atomic Structure

Atomic Structure
Unit 2
http://www.unit5.org/chemistry
Guiding Questions
How do we know atoms exist?
How do we know that electrons, protons, and
neutrons exist?
What is radiation and what does it come from?
Is radiation safe?
Where does matter come from?
How are elements formed?
Are all atoms of an element the same?
How do we measure atoms if they are so small?
How do we know what stars are made of?
What is wrong with this picture?
Atomic Theory and Atomic Structure
You should be able to
Discuss the development of the atom from its earliest model to
the modern day atom.
Identify the correct number of subatomic particles for atoms, ions,
and isotopes.
Calculate the average atomic mass of an atom from isotopic data.
Name compounds and write chemical formulas for binary compounds,
ternary compounds (those with polyatomic ions), and acids.

Memorize the chemical formulas and charged of the polyatomic ions and the most common transition metal ions.
Atomic Structure and Periodicity
You should be able to
Identify characteristics of and perform calculations with frequency and
wavelength.
Know the relationship between types of electromagnetic radiation and
energy; for example, gamma rays are the most damaging.
Know what exhibits continuous and line spectra.
Know what each of the four quantum numbers n, l, m, and ms represents.
Identify the four quantum numbers for an electron in an atom.
Write complete and shorthand electron configurations as well as orbital
diagrams for an atom or ion of an element.
Identify the number and location of the valence electrons in an atom.
Apply the trends in atomic properties such as atomic radii, ionization
energy, electronegativity, electron affinity, and ionic size.
The Hellenic Market
Fire
~
Water
Earth
Air
A Brief History of Chemistry
• In fourth century B.C., ancient Greeks proposed
that matter consisted of fundamental particles
called atoms.
• Over the next two millennia, major advances in
chemistry were achieved by alchemists. Their
major goal was to convert certain elements into
others by a process called transmutation.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
The Greeks
History of the Atom
• Not the history of atom, but
the idea of the atom
• In 400 B.C the Greeks tried to
understand matter
(chemicals) and broke them
down into earth, wind, fire,
and air.
~
~
• Democritus and Leucippus
Greek philosophers
Greek Model
“To understand the very large,
we must understand the very small.”
Democritus
• Greek philosopher
• Idea of ‘democracy’
• Idea of ‘atomos’
– Atomos = ‘indivisible’
– ‘Atom’ is derived
• No experiments to support
idea
• Continuous vs. discontinuous
theory of matter
Democritus’s model of atom
No protons, electrons, or neutrons
Solid and INDESTRUCTABLE
Democritus
DEMOCRITUS (400 BC) – First Atomic Hypothesis
Atomos: Greek for “uncuttable”. Chop up a piece of matter until you reach the atomos.
Properties of atoms:
• indestructible.
• changeable, however, into different forms.
• an infinite number of kinds so there are an infinite number of elements.
• hard substances have rough, prickly atoms that stick together.
• liquids have round, smooth atoms that slide over one another.
• smell is caused by atoms interacting with the nose – rough atoms hurt.
• sleep is caused by atoms escaping the brain.
• death – too many escaped or didn’t return.
• the heart is the center of anger.
• the brain is the center of thought.
• the liver is the seat of desire.
“Nothing exists but atoms and space, all else is opinion”.
Four Element Theory
FIRE
• Plato was an atomist
• Thought all matter was
composed of 4 elements:
–
–
–
–
–
Earth (cool, heavy)
Water (wet)
Fire (hot)
Air (light)
Ether (close to heaven)
Hot
Dry
‘MATTER’
AIR
Wet
EARTH
Cold
WATER
Relation of the four elements and the four qualities
Blend these “elements” in different proportions to get all substances
Some Early Ideas on Matter
Anaxagoras
(Greek, born 500 B.C.)
–Suggested every substance had its own kind of “seeds” that clustered together to
make the substance, much as our atoms cluster to make molecules.
Empedocles
(Greek, born in Sicily, 490 B.C.)
–Suggested there were only four basic seeds – earth, air, fire, and water. The
elementary substances (atoms to us) combined in various ways to make
everything.
Democritus
(Thracian, born 470 B.C.)
–Actually proposed the word atom (indivisible) because he believed that all
matter consisted of such tiny units with voids between, an idea quite similar to
our own beliefs. It was rejected by Aristotle and thus lost for 2000 years.
Aristotle
(Greek, born 384 B.C.)
–Added the idea of “qualities” – heat, cold, dryness, moisture – as basic elements
which combined as shown in the diagram (previous page).
Hot + dry made fire; hot + wet made air, and so on.
O’Connor Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 26,
Alchemy
• After that chemistry was
ruled by alchemy.
• They believed that that
could take any cheap
metals and turn them into
gold.
• Alchemists were almost
like magicians.
– elixirs, physical immortality
Alchemy
Alchemical symbols for substances…
..
.
......
.
.....
GOLD
SILVER
COPPER
IRON
SAND
transmutation: changing one substance into another
D
In ordinary chemistry, we cannot transmute elements.
Contributions
of alchemists:
Information about elements
- the elements mercury, sulfur, and antimony were discovered
- properties of some elements
Develop lab apparatus / procedures / experimental techniques
- alchemists learned how to prepare acids.
- developed several alloys
- new glassware
Timeline
Greeks
(Democritus ~450 BC)
Discontinuous
theory of matter
ALCHEMY
400 BC
300 AD
Issac Newton
(1642 - 1727)
1000
2000
Greeks
(Aristotle ~350 BC))
Continuous
theory of matter
American
Independence
(1776)
Dalton Model of the Atom
Late 1700’s - John Dalton- England
Teacher- summarized results of his experiments
and those of others
Combined ideas of elements with that of atoms
in Dalton’s Atomic Theory
The Atomic Theory of Matter
• In 1803, Dalton proposed that elements consist of
individual particles called atoms.
• His atomic theory of matter contains four hypotheses:
1. All matter is composed of tiny particles called atoms.
2. All atoms of an element are identical in mass and
fundamental chemical properties.
3. A chemical compound is a substance that always
contains the same atoms in the same ratio.
4. In chemical reactions, atoms from one or more
compounds or elements redistribute or rearrange in
relation to other atoms to form one or more new
compounds. Atoms themselves do not undergo a
change of identity in chemical reactions.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Foundations of Atomic Theory
Law of Conservation of Mass
Mass is neither destroyed nor created during ordinary chemical
reactions.
Law of Definite Proportions
The fact that a chemical compound contains the same elements
in exactly the same proportions by mass regardless of the size
of the sample or source of the compound.
Law of Multiple Proportions
If two or more different compounds are composed of the
same two elements, then the ratio of the masses of the
second element combined with a certain mass of the first
elements is always a ratio of small whole numbers.
Conservation of Atoms
2 H2 + O2
2 H2O
John Dalton
H
H
H2
O
H
O2
+
H2
H
O
H2O
O
H 2O
H
H
O
H
H
4 atoms hydrogen
2 atoms oxygen
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204
4 atoms hydrogen
2 atoms oxygen
Conservation of Mass
High
voltage
electrodes
Before reaction
After reaction
glass
chamber
O2
High
voltage
H2O
O2
5.0 g H2
H2
80
0 g H2
g O2
300 g (mass
of chamber)
+
385 g total
45
? g H2O
40 g O2
300 g (mass
of chamber)
+
385 g total
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204
Law of Definite Proportions
Joseph Louis Proust (1754 – 1826)
• Each compound has a specific ratio of
elements
• It is a ratio by mass
• Water is always 8 grams of oxygen for
every one gram of hydrogen
The Law of Multiple Proportions
• Dalton could not use his theory to determine the
elemental compositions of chemical compounds
because he had no reliable scale of atomic masses.
• Dalton’s data led to a general statement known as the
law of multiple proportions.
• Law states that when two elements form a series of
compounds, the ratios of the masses of the second
element that are present per gram of the first element
can almost always be expressed as the ratios of
integers.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Daltons Atomic Theory
• Dalton stated that
elements consisted of
tiny particles called atoms
• He also called the
elements pure
substances because all
atoms of an element
were identical and that in
particular they had the
same mass.
The Atomic Theory of Matter
• Dalton’s atomic theory is essentially correct, with four
minor modifications:
1. Not all atoms of an element must have precisely the same mass.
2. Atoms of one element can be transformed into another through
nuclear reactions.
3. The composition of many solid compounds are somewhat
variable.
4. Under certain circumstances, some atoms can be divided
(split into smaller particles: i.e. nuclear fission).
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Dalton’s Atomic Theory
1. All matter consists of tiny particles.
Dalton, like the Greeks, called these particles “atoms”.
2. Atoms of one element can neither be subdivided nor changed into
atoms of any other element.
3. Atoms can neither be created nor destroyed.
4. All atoms of the same element are identical in mass, size, and
other properties.
5. Atoms of one element differ in mass and other properties from
atoms of other elements.
6. In compounds, atoms of different elements combine in simple, whole
number ratios.
Dalton’s Atomic Theory
1. All matter is made of tiny indivisible particles
called atoms.
2. Atoms of the same element are identical,
those of different atoms are different.
3. Atoms of different elements combine in whole
number ratios to form compounds
4. Chemical reactions involve the rearrangement
of atoms. No new atoms are created or
destroyed.
California WEB
Structure of Atoms
• Scientist began to wonder what an atom
was like.
• Was it solid throughout with no internal
structure or was it made up of smaller,
subatomic particles?
• It was not until the late 1800’s that
evidence became available that atoms
were composed of smaller parts.
History: On The Human Side
1834
1895
1896
1896
1897
1898
1899
1900
Michael Faraday - electrolysis experiments
suggested electrical nature of matter
Wilhelm Roentgen - discovered X-rays when
cathode rays strike anode
Henri Becquerel - discovered "uranic rays" and
radioactivity
Marie (Marya Sklodowska) and Pierre Curie discovered that radiation is a property of the
atom, and not due to chemical reaction.
(Marie named this property radioactivity.)
Joseph J. Thomson - discovered the electron
through Crookes tube experiments
Marie and Piere Curie - discovered the
radioactive elements polonium and radium
Ernest Rutherford - discovered alpha and beta
particles
Paul Villard - discovered gamma rays
1919
1932
1934
1938
1940
1941
1942
1903
1910
1911
Ernest Rutherford and Frederick Soddy established laws of radioactive decay and
transformation
Frederick Soddy - proposed the isotope concept
to explain the existence of more than one atomic
weight of radioelements
Ernest Rutherford - used alpha particles to
explore gold foil; discovered the nucleus and the
proton; proposed the nuclear theory of the atom
1944
1964
Ernest Rutherford - announced the first artificial
transmutation of atoms
James Chadwick - discovered the neutron by
alpha particle bombardment of Beryllium
Frederick Joliet and Irene Joliet Curie - produced
the first artificial radioisotope
Otto Hahn, Fritz Strassmann, Lise Meitner, and
Otto Frisch - discovered nuclear fission of
uranium-235 by neutron bombardment
Edwin M McMillan and Philip Abelson discovered the first transuranium element,
neptunium, by neutron irradiation of uranium in a
cyclotron
Glenn T. Seaborg, Edwin M. McMillan, Joseph
W. Kennedy and Arthur C. Wahl - announced
discovery of plutonium from beta particle
emission of neptunium
Enrico Fermi - produced the first nuclear fission
chain-reaction
Glenn T. Seaborg - proposed a new format for
the periodic table to show that a new actinide series of 14
elements would fall below and be analogous to the 14
lanthanide-series elements.
Murray Gell-Mann hypothesized that quarks are the
fundamental particles that make up all known subatomic
particles except leptons.
Radioactivity (1896)
1. rays or particles produced by
unstable nuclei
a. Alpha Rays – helium nucleus
b. Beta Part. – high speed electron
c. Gamma ray – high energy x-ray
2. Discovered by Becquerel –
exposed photographic film
3. Further work by Curies
Antoine-Henri Becquerel
(1852 - 1908)
Radioactivity
• One of the pieces of evidence for the
fact that atoms are made of smaller
particles came from the work of
Marie Curie (1876 - 1934).
• She discovered radioactivity, the
spontaneous disintegration of some
elements into smaller pieces.
The Effect of an Obstruction on
Cathode Rays
High
voltage
source of
high voltage
shadow
cathode
yellow-green
fluorescence
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 117
Crooke’s Tube
-
voltage
source
William Crookes
+
vacuum tube
metal disks
magnet
A Cathode Ray Tube
Source of
Electrical
Potential
Stream of negative
particles (electrons)
Metal Plate
Gas-filled
glass tube
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58
Metal plate
Background Information
•
•
•
•
Cathode Rays
Form when high voltage is applied across
electrodes in a partially evacuated tube.
Originate at the cathode (negative electrode)
and move to the anode (positive electrode)
Carry energy and can do work
Travel in straight lines in the absence of an
external field
Cathode Ray Experiment
1897 Experimentation
• Using a cathode ray tube, Thomson was
able to deflect cathode rays with an
electrical field.
• The rays bent towards the positive pole,
indicating that they are negatively
charged.
The Effect of an Electric Field on
Cathode Rays
source of
high voltage
High
voltage
cathode
negative
plate
_
+
anode
positive
plate
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 117
Cathode Ray Experiment
Displacement
Volts
Anodes / collimators
Cathode
+
Deflection
region
Drift region
Thomson’s Calculations
Cathode Ray Experiment
• Thomson used magnetic and electric fields to measure and
calculate the ratio of the cathode ray’s mass to its charge.
Electric
=
deflection
Magnetic
=
deflection
charge of
ray particle
x
charge of
ray particle
x
length of
electric
x
deflection region x
field
mass of ray
velocity of 2
x
particle
ray particle
length of
drift region
length of
magnetic
x
deflection region x
field
mass of ray
velocity of
x
particle
ray particle
length of
drift region
magnetic deflection
electric deflection
magnetic field
=
electric field
x
velocity
Thomson
PAPER
Conclusions
• He compared the value with the mass/ charge ratio for the
lightest charged particle.
• By comparison, Thomson estimated that the cathode ray
particle weighed 1/1000 as much as hydrogen, the lightest
atom.
• He concluded that atoms do contain subatomic particles - atoms
are divisible into smaller particles.
• This conclusion contradicted Dalton’s postulate and was not
widely accepted by fellow physicists and chemists of his day.
• Since any electrode material produces an identical ray, cathode
ray particles are present in all types of matter - a universal
negatively charged subatomic particle later named the electron
J.J. Thomson
• He proved that atoms of
any element can be
made to emit tiny
negative particles.
• From this he concluded
that ALL atoms must
contain these negative
particles.
• He knew that atoms did
not have a net negative
charge and so there must
be balancing the negative
charge.
J.J. Thomson
William Thomson
(Lord Kelvin)
• In 1910 proposed
the Plum Pudding
model
– Negative electrons
were embedded into
a positively charged
spherical cloud.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
Spherical cloud of
Positive charge
Electrons
Thomson Model of the Atom
• J.J. Thomson discovered the electron and knew that
electrons could be emitted from matter (1897).
• William Thomson proposed that atoms consist of small,
negative electrons embedded in a massive, positive
sphere.
• The electrons were like currants in a plum pudding.
• This is called the ‘plum pudding’ model of the atom.
-
-
electrons
-
-
-
Rutherford
Ernest Rutherford (1871-1937)
PAPER
• Learned physics in
J.J. Thomson’ lab.
• Noticed that ‘alpha’
particles were
sometime deflected
by something in the
air.
• Gold-foil experiment
Animation by Raymond Chang – All rights reserved.
Rutherford ‘Scattering’
• In 1909 Rutherford undertook a series of experiments
• He fired a (alpha) particles at a very thin sample of gold foil
• According to the Thomson model the a particles would only
be slightly deflected
• Rutherford discovered that they were deflected through large
angles and could even be reflected straight back to the source
Lead collimator
Gold foil
a particle
source
q
Rutherford’s Apparatus
Rutherford received the 1908 Nobel Prize in Chemistry for his pioneering work in nuclear chemistry.
beam of alpha particles
radioactive
substance
circular ZnS - coated
fluorescent screen
gold foil
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120
What He Expected
• The alpha particles to pass through
without changing direction (very much)
• Because
• The positive charges were spread out
evenly. Alone they were not enough to
stop the alpha particles
California WEB
What he expected…
Because he thought the mass was
evenly distributed in the atom.
-
-
-
-
What he got…
richocheting
alpha particles
The Predicted Result:
expected
path
expected
marks on screen
Observed Result:
mark on
screen
likely alpha
particle path
Interpreting the
Observed Deflections
.
.
.
.
.
.
beam of
alpha
particles
.
.
.
.
.
undeflected
particles
.
.
.
.
.
gold foil
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120
.
deflected particle
Density and the Atom
• Since most of the particles went through,
the atom was mostly empty.
• Because the alpha rays were deflected so
much, the positive pieces it was striking
were heavy.
• Small volume and big mass = big density
• This small dense positive area is the
nucleus
California WEB
Rutherford Scattering (cont.)
Rutherford interpreted this result by suggesting that
the a particles interacted with very small and heavy
particles
Particle bounces off
of atom?
Case A
Case B
Particle goes through
atom?
Particle attracts
to atom?
Case C
Case D
.
Particle path is altered
as it passes through atom?
Table: hypothetical description of alpha particles
(based on properties of alpha radiation)
observation
hypothesis
alpha rays don’t diffract
... alpha radiation is a stream of particles
alpha rays deflect towards a negatively
charged plate and away from a positively
charged plate
... alpha particles have a positive charge
alpha rays are deflected only slightly by
an electric field; a cathode ray passing
through the same field is deflected
strongly
... alpha particles either have much
lower charge or much greater mass
than electrons
Copyright © 1997-2005 by Fred Senese
Explanation of Alpha-Scattering Results
Alpha particles
Nucleus
+
+
-
-
+
+
-
+
+
-
+
-
+
-
-
Plum-pudding atom
Nuclear atom
Thomson’s model
Rutherford’s model
Results of foil experiment if plumpudding had been correct.
Electrons scattered
throughout
-
+
-
positive
charges
+
+
+
+
-
-
+
+
+
-
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 57
-
Rutherford’s
Gold-Leaf
Experiment
Conclusions:
Atom is mostly empty space
Nucleus has (+) charge
Electrons float around nucleus
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120
The Rutherford Atom
-
-
n+
-
-
-
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 57
This is the modern atom model.
Electrons are in constant motion around the nucleus, protons and neutrons
jiggle within the nucleus, and quarks jiggle within the protons and neutrons.
This picture is quite distorted. If we drew the atom to scale and made protons
and neutrons a centimeter in diameter, then the electrons and quarks would
be less than the diameter of a hair and the entire atom's diameter would be
greater than the length of thirty football fields! 99.999999999999% of an
atom's volume is just empty space!
Website “The Particle Adventure”
Scale of the atom.
While an atom is tiny, the nucleus is ten thousand
times smaller than the atom and the quarks and
electrons are at least ten thousand times smaller
than that. We don't know exactly how small quarks
and electrons are; they are definitely smaller than
10-18 meters, and they might literally be points, but
we do not know.
It is also possible that quarks and electrons are
not fundamental after all, and will turn out to be
made up of other, more fundamental particles.
(Oh, will this madness ever end?)
Website “The Particle Adventure”
Physicists have developed a theory called The Standard Model that explains what
the world is and what holds it together. It is a simple and comprehensive theory that
explains all the hundreds of particles and complex interactions with only:
6 quarks.
6 leptons. The best-known lepton is the electron.
Force carrier particles, like the photon. We will talk about these particles later.
All the known matter particles are composites of quarks and leptons, and they
interact by exchanging force carrier particles.
The Standard Model is a good theory. Experiments have verified its predictions to
incredible precision, and all the particles predicted by this theory have been found.
But it does not explain everything. For example, gravity is not included in the
Standard Model.
Website “The Particle Adventure”
Discovery of the electron
1807 Davy suggested that electrical forces held
compound together.
1833 Faraday related atomic mass and the
electricity needed to free an element during
electrolysis experiments.
1891 Stoney proposed that electricity exists in units
he called electrons.
1897 Thomson first quantitatively measured the
properties of electrons.
Oil Drop Experiment
oil droplets
+
Small hole
Charged plate
. . ... ................................... .
........................................
...... . ........ ........... ......................
..... ..... . . ........ ..... ............................................. .. oil atomizer
............... . ................................................ .
. . . . . . ...... .. . .... .
.. . . . .. . . . .
.. ..........
..
.
.
.
.
......
.
.......... .. .
................... ... .
.. ...................... .
..................... .. .
...... ..........................................
............................... ... .. . .
.. ....... ..... ..... .. .....
....... ... . . ...... ... .
.
.
. .... .. . ..
..... ... ..
.......... .. .. . .
.. .
..
.. . .... .
Charged plate
Robert Millikan
Telescope
oil droplet
under observation
Bohr Atom
The Planetary Model of the Atom
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr Model of Atom
Increasing energy
of orbits
n=3
e-
n=2
e-
n=1
ee-
e-
e-
ee-
e-
e-
eA photon is emitted
with energy E = hf
The Bohr model of the atom, like many ideas in
the history of science, was at first prompted by
and later partially disproved by experimentation.
http://en.wikipedia.org/wiki/Category:Chemistry
Quantum Mechanical Model
Niels Bohr &
Albert Einstein
Modern atomic theory describes the
electronic structure of the atom as the
probability of finding electrons within certain
regions of space (orbitals).
Development of Atomic Models
Thomson model
Rutherford model
In the nineteenth century, Thomson described
the atom as a ball of positive charge containing
a number of electrons.
In the early twentieth century, Rutherford
showed that most of an atom's mass is
concentrated in a small, positively charged
region called the nucleus.
Bohr model
After Rutherford's discovery, Bohr proposed
that electrons travel in definite orbits around
the nucleus.
Quantum mechanical model
Modern atomic theory described the
electronic structure of the atom as the
probability of finding electrons within
certain regions of space.
Modern View
• The atom is mostly empty space
• Two regions
– Nucleus
• protons and neutrons
– Electron cloud
• region where you might find an electron
Dalton proposes the
indivisible unit of an
element is the atom.
Review
Models
of the
Atom
Thomson discovers
electrons, believed to
reside within a sphere of
uniform positive charge
(the “plum-pudding model).
Rutherford demonstrates
the existence of a positively
charged nucleus that
contains nearly all the
mass of an atom.
Bohr proposes fixed
circular orbits around
the nucleus for electrons.
In the current model of the atom,
electrons occupy regions of space
(orbitals) around the nucleus
determined by their energies.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Models of the Atom
"In science, a wrong theory can be valuable and better than no theory at all."
- Sir William L. Bragg
e
e
+
e
+
e
+
+
e
+e
+
e
e
+ e + e
Dalton’s
Greek model
model
(400
(1803)
B.C.)
Thomson’s plum-pudding
model (1897)
Bohr’s model
(1913)
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
-
- +
Rutherford’s model
(1909)
Charge-cloud model
(present)
Models of the Atom
e
e
+
e
-
+
e
+
+
e
+e
+e
e
+ e + e
Dalton’s
model
Greek model
(1803)
(400 B.C.)
1803 John Dalton
pictures atoms as
tiny, indestructible
particles, with no
internal structure.
1800
-
Thomson’s plum-pudding
model (1897)
- +
Rutherford’s model
(1909)
1897 J.J. Thomson, a British
1911 New Zealander
scientist, discovers the electron,
leading to his "plum-pudding"
model. He pictures electrons
embedded in a sphere of
positive electric charge.
Ernest Rutherford states
that an atom has a dense,
positively charged nucleus.
Electrons move randomly in
the space around the nucleus.
1805 ..................... 1895
1900
1905
1910
1904 Hantaro Nagaoka, a
Japanese physicist, suggests
that an atom has a central
nucleus. Electrons move in
orbits like the rings around Saturn.
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
1915
Bohr’s model
(1913)
1926 Erwin Schrödinger
1913 In Niels Bohr's
model, the electrons move
in spherical orbits at fixed
distances from the nucleus.
1920
1925
Charge-cloud model
(present)
1930
develops mathematical
equations to describe the
motion of electrons in
atoms. His work leads to
the electron cloud model.
1935
1940
1945
1924 Frenchman Louis
1932 James
de Broglie proposes that
moving particles like electrons
have some properties of waves.
Within a few years evidence is
collected to support his idea.
Chadwick, a British
physicist, confirms the
existence of neutrons,
which have no charge.
Atomic nuclei contain
neutrons and positively
charged protons.
Bohr Model
Neils Bohr
Planetary
model
After Rutherford’s discovery, Bohr
proposed that electrons travel in definite
orbits around the nucleus.
Bohr’s contributions
• Bohr’s contributions to the understanding of
atomic structure:
1. Electrons can occupy only certain regions of space,
called orbits.
2. Orbits closer to the nucleus are more stable —
they are at lower energy levels.
3. Electrons can move from one orbit to another by
absorbing or emitting energy, giving rise to
characteristic spectra.
• Bohr’s model could not explain
the spectra of atoms heavier
than hydrogen.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Particles in the Atom
Electrons
(-) charge
no mass
located outside the nucleus
1 amu
located inside the nucleus
1 amu
located inside the nucleus
Protons
(+) charge
Neutrons
no charge
Subatomic particles
Name
Symbol
Relative
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
no
0
1
1.67 x 10-24
Discovery of the Neutron
9
4
Be
+
4
2
He
12
6
C
+
1
0
n
James Chadwick bombarded beryllium-9 with alpha particles,
carbon-12 atoms were formed, and neutrons were emitted.
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 764
*Walter Boethe
Humor
Two atoms are walking down the street.
One atom says to the other, “Hey! I think I lost an electron!”
The other says, “Are you sure??”
“Yes, I’m positive!”
A neutron walks into a restaurant and orders a couple
of drinks. As she is about to leave, she asks the waiter
how much she owes. The waiter replies, “For you,
No Charge!!!”
Subatomic Particles
ATOM
ATOM
NUCLEUS
NUCLEUS
ELECTRONS
ELECTRONS
PROTONS
PROTONS
NEUTRONS
NEUTRONS
POSITIVE
Positive
CHARGE
Charge
NEUTRAL
Neutral
CHARGE
Charge
NEGATIVE
CHARGE
Negative Charge
equal in a
Atomic
Most Number
of the atom’s mass.
neutral atom
equals the # of...
QUARKS
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Subatomic Particles
• Quarks
– component of
protons &
neutrons
– 6 types
– 3 quarks =
1 proton or
1 neutron
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
He
Bohr - Rutherford diagrams
• Putting all this together, we get B-R diagrams
• To draw them you must know the # of protons, neutrons,
and electrons (2,8,8,2 filling order)
• Draw protons (p+), (n0) in circle (i.e. “nucleus”)
• Draw electrons around in shells
He
2 p+
2 n0
Li
Li shorthand
3 p+
4 n0
3 p+
4 n0
2e– 1e–
Draw Be, B, Al and shorthand diagrams for O, Na
Be
B
Al
4 p+
5 n°
O
5 p+
6 n°
13 p+
14 n°
Na
8 p+ 2e– 6e–
8 n°
11 p+ 2e– 8e– 1e–
12 n°
Mass Number
• mass # = protons + neutrons
• always a whole number
Neutron
+
• NOT on the
Periodic Table!
Electrons
Nucleus
e-
+
e-
e-
+
+
+
+
Nucleus
e-
ee-
Carbon-12
Neutrons 6
Protons
6
Electrons 6
Proton
Isotopes
• Atoms of the same element with different
mass numbers.
• Nuclear symbol:
Mass #
12
Atomic #
6
• Hyphen notation: carbon-12
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
C
Isotopes
Neutron
+
Electrons
Nucleus
+
+
+
+
+
Nucleus
Proton
Proton
Nucleus
Carbon-12
Neutrons 6
Protons
6
Electrons 6
+
+
+
+
Neutron
Electrons
+
+
Carbon-14
Neutrons 8
Protons
6
Electrons 6
Nucleus
6Li
7Li
3 p+
3 n0
3 p+
4 n0
2e– 1e–
2e– 1e–
Neutron
Neutron
Electrons
Electrons
+
Nucleus
+
+
Nucleus
+
Nucleus
Lithium-6
Neutrons 3
Protons
3
Electrons 3
Proton
+
+
Nucleus
Lithium-7
Neutrons 4
Protons
3
Electrons 3
Proton
17
Cl
Isotopes
37
• Chlorine-37
– atomic #:
17
– mass #:
37
– # of protons:
17
– # of electrons: 17
– # of neutrons: 20
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
37
17
Cl
Relative Atomic Mass
•
12C
atom = 1.992 × 10-23 g
• atomic mass unit (amu)
• 1 amu = 1/12 the mass of a 12C atom
• 1 p = 1.007276 amu
1 n = 1.008665 amu
1 e- = 0.0005486 amu
+
Electrons
Nucleus
+
Neutron
+
+
+
+
Nucleus
Carbon-12
Neutrons 6
Protons
6
Electrons 6
Proton
Average Atomic Mass
• weighted average of all isotopes
• on the Periodic Table
• round to 2 decimal places
Avg.
(mass)(%) + (mass)(%)
Atomic =
100
Mass
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Average Atomic Mass
• EX: Calculate the avg. atomic mass of oxygen if its
abundance in nature is 99.76% 16O, 0.04% 17O, and
0.20% 18O.
Avg.
(16)(99.76) + (17)(0.04) + (18)(0.20)
16.00
Atomic =
=
amu
100
Mass
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Average Atomic Mass
• EX: Find chlorine’s average atomic mass
if approximately 8 of every 10 atoms are
chlorine-35 and 2 are chlorine-37.
Avg.
(35)(8) + (37)(2)
Atomic =
= 35.40 amu
10
Mass
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Mass Spectrophotometer
magnetic field
heaviest
ions
stream
of ions of
different
masses
electron
beam
gas
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 138
lightest
ions
Weighing atoms
gas sample
enters here
.
ions accelerate
towards charged
slit
magnetic field
deflects lightest ions
most
filament current
ionizes the gas
The first mass spectrograph was
built in 1919 by F. W. Aston, who
received the 1922 Nobel Prize for
this accomplishment
ions separated by mass
expose film
• mass spectrometry is used to experimentally determine isotopic masses
and abundances
• interpreting mass spectra
• average atomic weights
- computed from isotopic masses and abundances
- significant figures of tabulated atomic weights gives some idea
of natural variation in isotopic abundances
Copyright © 1997-2005 by Fred Senese
Mass Spectrometry
198
200
202
Photographic plate
196
-
+
Stream of positive ions
Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320
199
201
204
Mass spectrum of mercury vapor
Mass Spectrum for Mercury
(The photographic record has been converted to a scale of relative number of atoms)
The percent natural abundances
for mercury isotopes are:
Hg-196
Hg-198
Hg-199
Hg-200
Hg-201
Hg-202
Hg-204
Relative number of atoms
30
25
198
0.146%
10.02%
16.84%
23.13%
13.22%
29.80%
6.85%
196
200
199
15
10
5
197
198
201
204
Mass spectrum of mercury vapor
20
196
202
199
200
Mass number
201
202
203
204
80
Mercury Isotopes
Hg
200.59
The percent natural abundances
for mercury isotopes are:
A
B
C
D
E
F
G
Hg-196
Hg-198
Hg-199
Hg-200
Hg-201
Hg-202
Hg-204
0.146%
10.02%
16.84%
23.13%
13.22%
29.80%
6.85%
(% "A")(mass "A") + (% "B")(mass "B") + (% "C")(mass "C") + (% "D")(mass "D") + (% "E")(mass "E") + (% F)(mass F) + (% G)(mass G) = AAM
(0.00146)(196) + (0.1002)(198) + (0.1684)(199) + (0.2313)(200) + (0.1322)(201) + (0.2980)(202) + (0.0685)(204) = x
0.28616 + 19.8396 + 33.5116 + 46.2600 + 26.5722 + 60.1960 + 13.974 = x
x = 200.63956 amu
92
Separation of Isotopes
U
238
Natural uranium, atomic weight = 238.029 g/mol
Density is 19 g/cm3. Melting point 1000oC.
Two main isotopes:
238
92
235
92
U
U
99.3%
0.7%
(238 amu) x (0.993) + (235 amu) x (0.007)
236.334 amu + 1.645 amu
Because isotopes are chemically identical
(same electronic structure), they cannot be
separated by chemistry.
So Physics separates them by diffusion or
centrifuge (mass spectrograph is too slow)…
237.979 amu
17
Average Atomic Mass
Cl
35.453
• Assume you have only two atoms of chlorine.
• One atom has a mass of 35 amu (Cl-35)
• The other atom has a mass of 36 amu (Cl-36)
• What is the average mass of these two isotopes?
35.5 amu
• Looking at the average atomic mass printed on the
periodic table...approximately what percentage is Cl-35
and Cl-36?
55% Cl-35 and 45% Cl-36 is a good approximation
Naming Isotopes
• Put the mass number after the name of
the element
• carbon- 12
• carbon -14
• uranium-235
California WEB
Calculating averages
• 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?
• Total mass = (4 x 50) + (1 x 60) = 260 g
• Average mass = (4 x 50) + (1 x 60) = 260 g
5
5
• Average mass = 4 x 50 + 1 x 60 = 260 g
5
5
5
California WEB
Calculating averages
• Average mass = 4 x 50 + 1 x 60 = 260 g
5 5
5
• Average mass = .8 x 50 + .2 x 60
• 80% of the rocks were 50 grams
• 20% of the rocks were 60 grams
• Average = % as decimal x mass +
% as decimal x mass +
% as decimal x mass +
California WEB
Isotopes
• Because of the existence of isotopes, the mass of a
collection of atoms has an average value.
• Average mass = ATOMIC WEIGHT
• Boron is 20% B-10 and 80% B-11.
That is, B-11 is 80 percent abundant on earth.
• For boron atomic weight
= 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu