Download The Modern Atomic Model

11/17/2015
Chapter 5: Electrons in Atoms
Bohr’s Model
• Niels Bohr
• Changed Rutherford’s model because it lacked
the ability to describe chemical properties.
• Rutherford’s also did not account for atomic
collapse.
• Changed the model based upon how atoms
absorb/emit light.
• Hydrogen atom
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Bohr’s Contributions
•Bohr discovered that electrons are not
free, random-motion particles.
•Major Contribution:
• Electrons are found in circular paths, or
orbits, around the nucleus.
• Fixed Energy
Bohr (cont.)
•Energy levels are like
rungs of a ladder.
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Energy Levels
• Higher energy level = further from
nucleus
• Quantum  Move between energy
levels.
Energy Levels
Bohr’s Ladder
(Increasing Energy)
6th Level
5th Level
4th Level
3rd Level
•Energy levels are not
equidistant.
2nd Level
1st Level
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Electrons in Energy Levels
• Degenerate vs. Excited Energy Levels
• Electrons are lazy.
• Emission of light.
Energy  Lazy = Happy
n=4
n=3
n=2
n=1
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Flaws with the Bohr Model
•Bohr’s Model:
• Showed electrons in controlled orbits.
(like planets) – often called “Planetary
Model”.
• Experiments greater than 1 electron
systems failed to reproduce this motion.
Oversights of Bohr
• Electrons are not only moving around the
nucleus but then are spinning as they move.
• 2 types of spin (more later)
• Uncertainty of location of e-
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Heisenberg’s Uncertainty
• Werner Heisenberg – you cannot know both
the momentum and the location of an e-, you
either know one or the other.
• “Heisenberg’s Uncertainty Principle”
The Modern Atomic Model
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Erwin Schrödinger
• 1926
• Created a math formula to explain how emoved in systems with more than 1 e-.
(Schrödinger Equation)
• E- exist in orbitals of probability.
Quantum Mechanical Model (e- cloud model)
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Hierarchy of the Atom
Quantum Numbers
• Quantum numbers describe the probable location of ein atoms.
• 3,1,0,+½
• ‘principal quantum number’ (n)  The energy levels in
the QM model are labeled by the (n). [n= 1, 2, 3, 4, ….]
• n ≠ 0  electrons do not exist in the nucleus
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QM Model – Sublevels
• The secondary quantum # ‘L’ (l)describes the shape of
the sublevel.
• l=0  s sublevel (spherical)
• l=1  p sublevel (dumbbell shaped)
• l=2  d sublevel (clover shaped)
• l=3  f sublevel (complex shape)
3,1,0, +½
• Range: l = 0…(n-1)
Magnetic Quantum Number (ml)
• The 3rd quantum number is represented (ml).
• This number divides the sublevels into their
individual orbitals which hold electrons.
3,1,0, +½
• Range: ml=-l … +l
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Shapes of Atomic Orbitals
Shapes of Atomic Orbitals
2s
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Spin Quantum Number (ms)
• Electrons are not only in motion based on their
velocities, they are also rotating upon on axis.
• Electrons can either have a +1/2 or -1/2 for their
‘ms’ values.
Orbital Occupancy
• Each orbital, in each energy level can only
hold a maximum of 2 e-.
• s sublevel  1 orbital, 2 e• p sublevel  3 orbitals, 6 e• d sublevel  5 orbitals, 10 e• f sublevel  7 orbitals, 14 e-
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The Atom
Identifying Quantum Numbers
Identify the following quantum numbers and the information the
principal and secondary quantum numbers tell us.
• EX:
a) 2,1,+1,+½  n =
l=
ml=
ms =
b) 3,2,0,-½  n =
l=
ml=
ms =
c) 1,0,0,+½  n =
l=
ml=
ms =
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