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Unit 4
THE ATOM
EVOLUTION OF The
Atomic Theory
Main idea
• The development of the atomic theory is a perfect
explain of how the scientific method works.
• This example illustrates the continuous changes many
scientific theories undergo over time.
• This example uses scientific models.
• This example uses numerous tests and experiments.
Democritus – 460 B.C.
• Greek Philosopher
• Only talk and thoughts,
no evidence or experiments.
• At some point you can’t divide
matter anymore.
• Named these smallest pieces an
atom, meaning “indivisible.”
• Atomic model: smallest pieces
Fast forward to 1800’s
• Greek society collapses and harder times force science
to takes a backseat to basic survival.
• Atomic theory lays dormant for about 2000 years.
• The industrial revolution (and technology) during the
1800’s allows science to prosper once again.
• Getting quantitative….
• Law of conservation of mass
• Law of definite proportions
• Law of multiple proportions
Law of conservation of mass - 1782
• Mass is neither destroyed nor created during ordinary
chemical and physical reactions.
• mass of reactants = mass of products
Carbon + Oxygen  Carbon Dioxide
C
12g
+
O2
32g

CO2
44g
Law of DEFINITE PROPORTIONS- 1799
• A given compound always contains the same fixed
ratio of elements.
• The recipes for chemical compounds never change.
• Water is always
• 11.20% Hydrogen
• 88.80% Oxygen
Dalton’s atomic theory - 1803
• English School Teacher
• Developed 1st atomic theory
• Based on experiments
• Atomic Model: hard sphere
Dalton’s atomic theory - 1808
1. All matter is composed of extremely small particles
called atoms.
2. Atoms of a given element are identical in size, mass
and other properties.
3. Atoms cannot be subdivided, created or destroyed.
4. Atoms of different elements combine in simple
whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms can be combined,
separated and rearranged.
Conclusion
1. All matter is composed of extremely small particles
called atoms.
2. Atoms of a given element are identical in size, mass
and other properties.
3. Atoms cannot be subdivided, created or destroyed.
4. Atoms of different elements combine in simple
whole-number ratios to form chemical compounds.
5. In chemical reactions, atoms can be combined,
separated and rearranged.
Discovery of the electrons – 1897
English Physicist
Thomson
Preformed experiments with cathode-ray tubes.
Cathode-Ray Tube
Voltage Source
CATHODE
Cathode Ray
(a stream of negative particles)
ANODE
An object placed between the
cathode ray and the opposite end
of the tube cast a shadow on
the glass.
This showed movement and
direction.
http://www.youtube.com/watch?v=Xt7ZWEDZ_GI&feature=related
A paddle wheel placed on rails between
the electrodes rolled along the rails from
the cathode towards the anode.
This showed cathode ray had a mass or force.
Cathode rays could be manipulated by magnets. The ray was
attracted to the positive side of a magnet and repelled by the
negative side of the magnet.
http://www.youtube.com/watch?v=7YHwMWcxeX8
Conclusions
1. atoms have negatively charged particle called electron
2. atoms are divisible
3. hypothesized that atoms must also contain a positively
charge particle that balance the atoms.
Atomic Model: Plum Pudding
Discovery of the nucleus
• Ernest Rutherford (1911)
• Gold Foil Experiments
• Atomic Model:
Nuclear Model
Gold foil Experiment
Most of the alpha particles
pass straight through.
Some hit something that is
dense and positive.
Conclusion
• Rutherford concluded that the rebounded alpha particles
must have experienced some powerful force within the
atom. And he figured that the source of this force must
occupy a very small amount of space because only a few
of the total number of alpha particles had been affected by
it.
• Atoms have a very densely packed bundle of positive
matter called the nucleus. Later, it will be found to be the
proton.
Planetary model
• Niels Bohr (1912)
• Spectroscopy (Light) Experiments
• Atomic Model: Planetary Model
Planetary model
• Electrons move around the nucleus in fixed
energy levels.
• These energy levels are
like rungs on a ladder where
energy is needed to move up
“the ladder”
• Think of the electrons as a moving fan or bee!
Conclusion
• Suggested that the electrons orbiting the nucleus of
atoms can only have certain discrete energies.
• Electrons lose/gain energy in order to move from one
energy level to another.
• The emission of light (photon) occurs when an
electron moves from a higher to a lower energy orbit
(emission spectra).
• According to Bohr, electrons exist as sub-atomic
particles found in neat orbits around the nucleus.
• Many scientists soon found that this idea of the
electron’s location was an over simplification.
• Electrons can have properties of both particles and waves
which make them much more complex to understand.
Structure of the
atom
The Atom
• The smallest unit of an element that maintains the
properties of that element.
• So what does all of this mean for the shape, structure
and contents of the atom…
Properties OF subatomic particles
Electron
Neutron
Proton
-1
0
+1
Negative
Neutral
Positive
Mass
1/1836 amu
(Small)
1 amu
(Large)
1 amu
(Large)
Location
Outside Nucleus
(Electron Cloud)
Nucleus
Nucleus
Charge
AMU = Atomic Mass Unit
The Atom
• Electrons are located outside the nucleus in the
electron cloud.
• Most of the volume of an atom is occupied by the
electron cloud.
• The protons and neutrons are located inside in the
center of the atom in the nucleus.
• The majority of the mass of an atom is found in the
nucleus (neutrons & protons).
The Atom
• Typically, particles with the same charge repel each
other.
• Nuclear forces - short range forces that help hold the
atom together.
• Proton – proton
• Neutron – neutron
• Neutron – proton
Atomic number &
atomic mass
Atomic Number
Atomic Number = Number Of Proton
Different For Each Element
Atomic NUMBER
• Number of protons = Number of electrons
• Result is a neutral atom
Isotopes
• Atoms of a given element that contain different
numbers of neutrons.
• Remember: Neutrons have a sizable mass.
• Therefore: Atoms of the same element that have
different masses.
• Same number of protons (+) and electrons (-).
Isotope Notation
• Hydrogen – 3
• Hydrogen – 2
• Hydrogen – 1
• Number after element’s name represents isotopes
atomic mass
Atomic mass
Average mass of all the isotopes of an element.
Atomic Mass
Atomic Mass = number of protons + number of neutrons
number of neutrons = atomic mass - atomic number
The Mole
• SI unit for amount of substance.
• the amount of a substance that contains as many particles as
there are atoms in exactly 12g of carbon - 12
• Counting unit like a dozen or π.
Avogadro’s Number
• 6.0221367 X 1023 rounded to 6.02 X 1023
• The number of particles in exactly one mole of a pure
substance.
Molar mass
(Formula Weight)
• The mass of one mole of a pure substance (g/mol)
• Equal to the atomic mass of an element or the
elements in a compound.
• Examples:
Hydrogen – 1.00794g/mol or 1g/mol
Oxygen – 15.994g/mol or 16g/mol
Atom/Mass/Mole Conversions
• Sample Problem
•
•
•
•
Finding mass from moles
Finding moles from mass
Finding atoms from moles
Finding moles from atoms