Transcript A is for atom

Developing the Modern
Atomic Model
SCH4U
Learning Goals/Success Criteria
• I will be able to discuss the development of the atom from earliest
atomic theory to the modern theory of the atom used today.
• I will be able to explain how experimental observations and
inferences made by Ernest Rutherford and Niels Bohr contributed to
the development of the planetary model of the hydrogen atom.
– Be able to use appropriate terminology related to atomic theory such as orbital,
emission spectrum, wavelength, frequency, energy level, photon etc.
– Explain Rutherford’s gold foil experiment, What Rutherford expected to happen
but what he observe instead.
– Describe the limitations to Rutherford’s model
– Bohr’s theory of the hydrogen ion spectrum
28/03/2016
A is
for
ATOM
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Evidence for nuclear atom using
alpha particle scattering.
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The Nuclear Atom
History of the Atom - Democritus
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From the ancient Greeks through to the
19th century, there has been the question
‘what is matter made from?’
The idea of atoms was first proposed by
Greek Philosopher Democritus in 530 B.C.
The concept was that matter could only be split in half
and half again until indivisible units were reached.
ATOMS
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The Nuclear Atom
History of the Atom – John Dalton
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In 1808, John Dalton (a teacher!)
proposed the modern ATOMIC
THEORY.
It simply states that all elements are
made up of atoms and an element is
only made up from one type of atom.
Dalton's view of tiny indivisible spheres
remained unchallenged until the end of
the 19th Century.
Dalton’s
Atom
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Dalton
• Lavoisier – Law of Conservation of Mass
• Proust – Law of Definite Proportions
• Dalton – Law of Multiple Proportions
Dalton’s Atomic Theory – 1805
-all matter consists of tiny particles called atoms
-atoms cannot be subdivided, created or destroyed
-all atoms of an element are identical
-atoms of different elements are different from each other.
-atoms of different types combine in specific ratios to form compounds.
The Nuclear Atom
Discovering the Electron
28/03/2016
In 1897, Joseph (JJ) Thomson (British
Physicist) was experimenting with
electrical currents through gases.
The cathode rays he produced could be
deflected or moved when in
electromagnetic fields.
Cathode rays were made up of tiny
negatively charged particles –
ELECTRONS.
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The Nuclear Atom
Thomson’s Model of the Atom
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From his evidence, Thomson
proposed that atoms were made up of
just tiny electrons.
He accounted for the neutrality of
atoms by the stating the electrons
existed in a ‘soup of positive charge’.
Sometimes referred to the plumpudding model.
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The Nuclear Atom
Radioactivity
28/03/2016
Around the same time, Henri Becquerel
discovered that some unstable elements
gave off smaller particles –
RADIOACTIVITY.
Therefore atoms must be divisible and
made up of smaller parts –
SUBATOMIC PARTICLES.
Marie and Pierre Curie and Ernest
Rutherford confirmed this.
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DIG – The Dating Game
Radiation
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
Alpha Particle: Positively charge. Large in comparison.
Essentially a Helium Nucleus (as proved by Rutherford)

Beta Particle: Negatively charged. Light. (later to be
shown as electrons)

Gamma Rays: Neutrally charged. No mass – Energy.
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The Nuclear Atom
Rutherford’s Experiment
28/03/2016
Rutherford and his colleagues bombarded a thin foil of gold
with a beam of alpha particles and then onto a fluorescent
screen.
Small amounts were
deflected.
Fluorescent
Screen
99.9% passed straight
through unaffected.
Thin Gold Foil
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The Nuclear Atom
28/03/2016
Alpha Particle Scattering
Why were alpha particles scattered?
To explain back scattering Rutherford proposed the
Nuclear Model of the Atom.
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The Nuclear Atom
Alpha Particle Scattering
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Alpha particles are positive.
High speed alpha particle bullet
travels through atom.
The electrons have little effect since
they are very light and the electrons
in the pudding model are very spread
out.
Very little deflection.
Does not support observations.
Plum-Pudding Model
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The Nuclear Atom
Alpha Particle Scattering
28/03/2016
What is observed is that alpha
particles in some instances are
strongly deflected.
Alpha Source
With electrons practically
dismissed, the only electrostatic
force available could be a positive
charge somewhere within the
atom.
An atom
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The Nuclear Atom
28/03/2016
Rutherford’s Model
He suggested that all of the atom’s positive
charge, together with most of its mass, is
concentrated
in the centre.
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Alpha particles which travel close to the
nucleus are strongly deflected. The degree
of deflection depends on how close it
approaches.
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The Nuclear Atom
Rutherford’s Model
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The nucleus must be very small in comparison
to the atom.
This will account for the vast majority making it
through unaffected.
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The Nuclear Atom
28/03/2016
Rutherford’s Nuclear Model of an Atom
In summary,
He envisioned an atom that had a
positively charged nucleus in the
centre.
The atom was mostly empty
space.
An he deemed it reasonable that
electrons orbit this nucleus like
planets orbit the Sun.
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Nuclear Model of an
Atom
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The Nuclear Atom
28/03/2016
Rutherford’s Nuclear Model of an Atom
The model appeared flawless and
convinced most of the scientific
community.
Rutherford and his colleagues (Hans
Geiger and Ernest Marsden) were able
to precisely predict the effects of:
Alpha particle energy
Thickness of sample
Different metals
However there was a problem…
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Nuclear Model of an
Atom
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The Nuclear Atom
The Problem…
28/03/2016
As the electrons move in circles, they would lose energy.
Losing energy would slow them down.
Therefore they would be pulled into the positively charged
nucleus.
It has been calculated that a Rutherford atom would only
exist for about 1 billionth of a second!
The answer lies within QUANTUM MECHANICS –
when things get really small!!
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Rutherford’s model could not
explain:
• Why the electrons did not
lose energy as they
orbited.
• What held the protons
together in the nucleus.
• The origins of emission
spectra of gases could not
be explained.
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The Nuclear Atom
28/03/2016
James Chadwick’s Protons
The number of protons in a nucleus did not match the atomic
weight of the atom.
Therefore a third neutrally charged particle must exist!
Alpha Radiation
Neutron Released
Beryllium Foil
These he named NEUTRONS.
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Atomic Spectra
A glass prisim can be
used to generate a colour
spectrum.
If this the light generated
by a hot (glowing) gas is
viewed through a prism
specific colour lines are
seen as AN EMISSION
LINE SPECTRUM.
If light is shone through a
cold sample of the same
gas, the same specific
colour lines are absent
and appear as an
ABSORPTION LINE
SPECTRUM.
Bohr’s Atom
Bohr - brought the concept of
quantization into atomic
theory.
Electrons could only move in
certain specific orbits
corresponding to specific
amounts of energy.
These ENERGY LEVELS radiated
out from the nucleus with
higher energies being further
away.
Electrons do not radiate energy
in these orbits.
Energy is gained or lost when
they move between orbits.
This model enabled Bohr to
explain the hydrogen
spectrum.
On what basis could Bohr make these claims?
In 1900, a physicist named Max Planck proposed an idea:
Planck suggested that matter,
 at the atomic level, can absorb or emit only discrete
quantities of energy.
 Each of these specific quantities.is called a quantum of
energy.
 In other words, Planck said that the energy of an atom is
quantized. (can exist only in certain discrete amounts).
 a quantum is an extremely small "packet" of energy.
Line
Spectra
Gases absorb certain
frequencies of light.
Each gas absorbs a
unique combination of
frequencies – each
frequency corresponding
to a unique colour.
So each gas has a
unique set colours which
is known at its “line
spectra” – because they
are unique they can be
used to identify a gas –
similar to fingerprints.
Absorption & Emission spectrum
ABSORBED LIGHT
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EMITTED LIGHT
• In absorption spectrum radiation is again absorbed
by electrons being excited to higher energy levels.
• The same frequencies (colours) are again emitted
when the excited electrons drop to the ground state
in an emission spectrum.
The Bohr-Rutherford model
http://www.upscale.utoronto.ca/GeneralInterest/Harrison/BohrModel/Fla
sh/BohrModel.html
Emission Spectrum
Excited electrons dropping
from unstable energy levels
radiate energy in the form of
light.
The frequency (colour) of the
radiation is directly related to
the energy gap between the
energy levels.
Since each element has its own
unique series of energy levels,
each element also has its own
unique series of
emission/absorption lines.
The line spectrum can therefore
be used to identify each
element much like a fingerprint.
References
• Simon Ball
• Nelson 12
• Mcgraw-Hill Ryerson