Transcript Topic 7. 1 Atomic Structure

 The modern atom has gone through a few stages of
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development
Dalton’s Atomic Therory – idea of an atom
JJ Thompson – 1890 – negative charge (electrons)
Earnest Rutherford – 1911 - positive nucleus (protons)
Niels Bohr – 1913 – orbital shells
Chadwick – 1932 – neutrons
 This is a VERY simplified idea of the atom
 Nucleus
 Protons – positive charge – 1.6 x 10-19C
 Neutrons – no charge
 Diameter order of 10-15m
 Electron “cloud”
 Electrons – negative charge – 1.6 x 10-19C
 Diameter order of 10-10m
 The nucleus is about 100,000 times smaller than the
electron orbits.
 Imagine a pea in the center of a football field with the track
being the orbits.
 Protons and Neutrons have very similar mass.
 Protons and Neutrons are about 1800 times bigger than
electrons.
Dalton’s Atomic Theory
1. All matter is composed of extremely small particles
called atoms.
2. All atoms of a given element are identical.
3. Atoms cannot be created, divided into smaller
particles, or destroyed.
4. Different atoms combine in simple whole number
ratios to form compounds.
5. In a chemical reaction, atoms are separated,
combined or rearranged.
Deomcritus
Atoms
Differences in atoms
1. All matter is composed of extremely small particles
2.
3.
4.
5.
called atoms.
All atoms of a given element are identical.
Atoms cannot be created, divided into smaller
particles, or destroyed. (This part proven wrong)
Different atoms combine in simple whole number
ratios to form compounds.
In a chemical reaction, atoms are separated,
combined or rearranged.
Deomcritus
Atoms
Differences in atoms
Dalton
•Atoms
•Sameness
•Created/destroyed
•Combination
•Rearragement
 J. J. Thomson – 1890-1900
 Used cathode ray tube to
prove existence of
electron.
 Proposed “Plum Pudding
Model”
 Cathode ray tube
 Stream of charged
particles (electrons).
http://www.youtube.com/wa
tch?v=YG-Wz-arcaY
http://www.youtube.com/wa
tch?v=O9Goyscbazk
 Plum Pudding
 J. J. Thompson
 Plum Pudding Model
Deomcritus
Thompson
Atoms
•Atoms composed
of electrons
Differences in atoms
Dalton
•Atoms
•Sameness
•Created/destroyed
•Combination
•Rearragement
 Gold Foil experiment
 Used to prove the existence
of a positively charged core
(Nucleus)
 Fired alpha
particles(2protons and 2
neutrons) into very thin gold
foil.
 The results were “like
firing a large artillery shell
at a sheet of paper and
having the shell come back
and hit you!”
 What should have
happened
• What DID happened
 After performing hundreds of tests and calculations,
Rutherford was able to show that the diameter of the
nucleus is about 105 times smaller than the diameter of the
atom
Deomcritus
Thompson
Atoms
•Atoms composed
of electrons
Differences in atoms
Rutherford
Dalton
•Atoms
•Sameness
•Created/destroyed
•Combination
•Rearragement
•Positively Charged
Nucleus
 Chadwick
 Worked with Rutherford.
 Noted there was energy in the nucleus, but wasn’t the
protons.
 Concluded that neutral particles must also exist in nucleus.
 James Chadwick – 1932
 Bombarded a beryllium target with alpha particles
 Alpha particles are helium nucleus
 Discovered that , carbon was produced with another
particle.
 Concluded this particle had almost identical mass to
proton but no charge.
 Called it a neutron
Deomcritus
Thompson
Atoms
•Atoms composed
of electrons
Differences in atoms
Rutherford
Dalton
•Atoms
•Positively Charged
Nucleus
•Sameness
Chadwick
•Created/destroyed
•Neutrons
exist in
Nucleus
•Combination
•Rearragement
 Three main particles:
 Proton
 Positive
 In nucleus
 Neutrons
 Neutral
 In nucleus
 Electrons
 Negative
 Orbiting the nucleus (not inside)
 If Rutherford’s was correct, electrons orbiting would
undergo centripetal acceleration.
 This would mean they would radiate electromagnetic waves.
 Meaning they would loose energy
 Meaning the atom would collapse on it’s self
 If low-pressure gases are heated or current is passed through
them they glow.
 Different colors correspond to their wavelengths.
 Visible spectrum 400nm(violet) to 750nm(red)
 Gas – slit – slit – prism – viewing screen
 When single element gases such as hydrogen and helium
are excited only specific wave lengths were emitted.
 These are called emission line spectra
 Light – gas vapor – slit – slit – prism – viewing screen
 If white light is pass through the gas the emerging light will
show dark bands called absorption lines.
 They correspond to the emission lines.
 Rutherford’s model didn’t explain why atoms emitted or
absorbed only light at certain wavelengths.
 1885 JJ Balmer showed that hydrogen’s four emission lines
fit a mathematical formula.
 This “Balmer series” also show the pattern continued into
non-visible ultra-violet and infra-red.
 Bohr called these “energy levels”
 Reasoned that the electrons do not lose energy
continuously but instead, lose energy in discrete amounts
called “quanta”.
 He agreed with Rutherford that electrons orbit the nucleus
but only certain orbits were allowed.
 The electric force between protons and electrons holds
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electrons in orbit
Electron never found between these levels. (“jumps”
instantly)
Only radiates energy when it “jumps” down.
Absorbs energy when it “jumps” up.
Total energy stays constant
 Bohr explained the emission and absorption line spectra
with the idea that electrons absorbed only certain quantity
of energy that allowed it to move to a higher orbit or energy
level.
 Each element has its own “finger print”.
 Ground state – lowest energy level –
smallest possible radius
 Excited state – when an electron
absorbs energy and jumps to a higher
energy level.
 Once an electron jumps back to a
lower state it gives off energy in the
form of a photon.
 These photons are the emission
spectrum.
 The amount jumped
correlates to the energy of
the photon.
 Greater the jump means
the greater the energy is
emitted.
 Each jump corresponds to
a different amount of
energy being released. This
means we can calculate the
frequency and wavelength
of light that will be
produced.
 E = hf
 E = energy of a quantum
 h = Planck’s constant (6.63 x 10-34Js)
 f = frequency
 An electron in a hydrogen atom drops from energy level E4
to energy level E2. What frequency of the emitted photon,
and which line line in the emission spectrum corresponds
to this event?
 First find the amount of energy lost
 Elost = E4 – E2
 Elost = (-0.85eV) – (-3.40eV)
 Elost = 2.55 eV
 Second, convert eV into J.
 1eV = 1.6 x 10-19J
 Answer: 4.08 x 10-19J
 Third use Planck’s equations to find the frequency.
 E = hf
 f = 6.15 x 1014 Hz
 Fourth decide which line corresponds to this even.
 Answer: Green light
 v=fλ
 Practice C, pg 769 in book, #2 – 5
 Definitions
 Nucleon – any of the constituents of a nucleus. Protons
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and neutrons.
Atomic Number – The number of protons in the nucleus.
Nucleon Number – The number of nucleons in the
nucleus. AKA the mass number. (protons + neutrons)
Isotope – Nuclei which contain the same number of
protons but different numbers of neutrons.
Nuclide – the nucleus of an atom. The nuclides of
isotopes are different, even though they are the same
element.
 Atomic Number (proton number), Z
 How many protons there are.
 This is what defines the element.
 Ex. Hydrogen Z =1, Oxygen Z = 8 Carbon Z = 6
 Nucleon Number (mass number), A
 How many nucleons there are.
 Protons + neutrons
 Number of neutrons, N
 Mass number = atomic number + number of neutrons
 A=Z+N
 Standard notation is: A over Z in front of element(X)
*****Draw on board*****
Isotopes
 More evidence for neutrons is the existence of isotopes.
 When nuclei of the same element have different numbers
of neutrons.
 Carbon has 6 isotopes: Carbon-11, Carbon-12, Carbon-13,
Carbon-14, Carbon-15, Carbon-16.
 All have 6 protons but each has different number of
neutrons.
 The different isotopes don’t exist in nature in equal
amounts.
 Carbon:
 C – 12 is most abundant (98.9%)
 C – 13 is next (1.1%)
 This is where atomic mass comes from. It’s the weighted
average mass of all the different isotopes.
 Nuclei of different atoms are known as nuclides.
 Ex. C – 12, C – 14
 Both are carbon but different isotopes
 Their nuclei have different numbers of neutrons.
 These are different nuclides.
 How do like charge (protons), stay stuck together?
 We already know that like charges repel each other.
 We have also seen that they are stronger than gravitational
forces.
 Strong Force – The force that binds the nucleus together.
 It is an attractive force that acts between all nucleons.
 Short – range interactions only (up to 10-15m)
 7.3.3 - Define the term unified atomic mass unit.
 7.3.4 - Apply the Einstein mass-energy equivalence
relationship.
 7.3.5 - Define the concepts of mass defect, binding energy
and binding energy per nucleon.
 7.3.6 - Draw and annotate a graph showing the variation
with nucleon number of the binding energy per nucleon.
 7.3.7 - Solve problems involving mass defect and binding
energy.
 Because the mass of an atom is so small a new unit was
created.
 Some times called “Atomic mass unit”
 1 u = 1.66053886 x 10-27 kg
 12u = one atom of carbon-12
 Mass of a nucleus is sometimes expressed in terms of rest
energy.
 A particle has a certain amount of energy associated with
its mass.
Relationship between rest energy and mass:
ER = mc2
 It doesn’t always happen with nuclear processes.
 Some times mass is converted or lost in the form of energy.
 1u = 931.49 MeV
 So that means that one proton IS 938.3MeV of energy.
 Mass is energy, energy is mass THEY ARE THE SAME
THING!!! AHHHHHH!!!!!!
Check out the table
 What happens when you place two negative charged
particles next to each other?
 What happens when you place two positively charged
particles next to each other?
 So why doesn’t a nucleus explode?
 It shouldn’t stay together.
 Strong Force
 Attractive force
 Independent of electric charge
 Very short range
 Neutrons!!!
 Spread the protons apart to help balance electrical repulsion
and strong attraction
 Particles in a stable nuclease need an input of energy to
break the strong nuclear force.
 When to unbound particles come together energy is
released. (think nuclear reactions)
 Turns out these quantities of energy are the same.
 Called the binding energy
 Binding energy is the energy it takes to hold the atom
together.
 Equal to the
 Recall that mass is energy.
 Carbon – 12
 Atom of carbon – lighter, less rest energy
 Constituent parts of – heavier, more rest energy
 What happen to that little bit of matter?
 It is used as the energy to bind together the atom.
The difference in the two masses is known as mass defect (∆m)
 Binding energy = mass defect x (speed of light)2
 Ebind = ∆m c2
 E = mc2
 The nucleus of the deuterium atom, called deuteron,
consists of a proton and a neutron. Given that the atomic
mass of deuterium is 2.014 102u, calculate the deuteron’s
binding energy in MeV.
 Answer: 2.224MeV
 If the phosphorus has a mass of 30.973 762u, then what is
the binding energy that holds the nucleus together in MeV?
 Answer:
 Practice A, pg 795 in book, #1,3-4
 Answers:
 1) 160.65MeV, 342.05MeV
 2) 0.764MeV
 3) 7.933MeV
 4) 7.5701 MeV/nucleon