Transcript Chapter 4
Chapter 4:
Arrangement of
Electrons in Atoms
Chemistry
Development of a New
Atomic Model
There were some problems with the
Rutherford model…It did not answer:
Where the e- were located in the space
outside the nucleus
Why the e- did not crash into the nucleus
Why atoms produce spectra at specific
wavelengths
Properties of Light
Wave-Particle Nature of Light – early
1900’s
A Duel Nature
It was discovered that light and e- both have
wave-like and particle-like properties
Wave Nature of Light
Electromagnetic radiation – form of
energy that exhibits wave-like behavior
as it travels through space
Electromagnetic spectrum
All the forms of electromagnetic radiation
Speed of light in a vacuum
3.0 x 108 m/s
Wave Nature of Light
Wavelength
Distance between two corresponding points on
adjacent waves
λ
nm
Frequency
Number of waves that pass a given point in a
specified time
ν
Hz - Hertz
Wave Nature of Light
Figure 4-1, page 92
Equation
c=λν
Indirectly related!
Spectroscope
Device that separates light into a spectrum that can
be seen
Diffraction Grating – the part of the spectroscope
the separates the light
Particle Nature of Light
Quantum
Minimum quantity of energy that can be lost
or gained by an atom
Equation
E=hν
Direct relationship between quanta and
frequency
Planck’s Constant (h)
h=6.626 x 10-34 Js
Particle Nature of Light
Photon
Individual quantum of light; “packet”
The Hydrogen Atom
Line emission spectrum (Figure 4-5, page 95)
Ground State
Lowest energy state (closest to the nucleus)
Excited State
State of higher energy
**When electron drops from its excited state to its ground
state, a photon is emitted! This produces a bright-line
spectrum. Each element has a characteristic bright-line
spectrum – much like a fingerprint!**
http://jersey.uoregon.edu/vlab/elements/Eleme
nts.html
Particle Nature of Light
Why does an emission spectrum occur?
Atoms get extra energy – voltage
The e- jumps from ground state to excited
state
Atoms return to original energy, e- drops
back down to ground state
Continuous spectrum
Emission of continuous range of frequencies
Particle Nature of Light
Bohr Model of the H atom
1913 – Danish physicist – Niels Bohr
Single e- circled around nucleus in allowed paths or
orbits
e- has fixed E when in this orbit (lowest E closest to
nucleus)
Lot of empty space between nucleus and e- in which
e- cannot be in
E increases as e- moves to farther orbits
http://chemmovies.unl.edu/ChemAnime/BOHRQD/B
OHRQD.html
Particle Nature of Light
Bohr Model (cont)
ONLY explained atoms with one e Therefore – only worked with hydrogen!!
Particle Nature of Light
Spectroscopy
Study of light emitted by excited atoms
Bright line spectrum
The Quantum Model of the
Atom
e- act as both waves and particles!!
De Broglie
1924 – French physicist
e- may have a wave-particle nature
Would explain why e- only had certain orbits
Diffraction
Bending of wave as it passes by edge of object
Interference
Occurs when waves overlap
The Quantum Model of the
Atom
Heisenberg Uncertainty Principle
1927 – German physicist
It is impossible to determine simultaneously
both the position and velocity of an e-
12:28-14:28
The Quantum Model of the
Atom
Schrodinger Wave Equation
1926 – Austrian physicist
Applies to all atoms, treats e- as waves
Nucleus is surrounded by orbitals
Laid foundation for modern quantum theory
Orbital – main energy level; 3D region
around nucleus in which an e- can be found
Cannot pinpoint e- location!!
Quantum Numbers
Quantum Numbers
Solutions to Schrodinger’s wave eqn
Probability of finding an e“address” of eFour Quantum Numbers
Principle
Anglular Momentum
Magnetic
Spin
Principle Quantum
Number
Which main energy level? (“orbital”
“shell”)
Symbol- n
n is normally 1-7 (corresponds to period
on periodic table)
Higher the n, the greater the distance
from the nucleus
Angular Momentum
Quantum Number
What is the shape of the orbital?
F shape
Symbol – l
l = s,p,d,f
When n = 1, l = s
n = 2, l = s,p
n = 3, l = s,p,d
n = 4, l = s,p,d,f
http://www.chemeng.uiuc.edu/~alkgrp/mo/gk12
/quantum/
Magnetic Quantum
Number
Orientation of orbital around nucleus
Symbol – m
s–1
p–3
d–5
f–7
Every orientation can hold 2 e-!!
Figures 4-13, 4-14, 4-15 on page 102-103
Spin Quantum Number
Each e- in one orbital must have opposite
spins
Symbol – s
+½,-½
Two “allowed” values and corresponds to
direction of spin
Electron Configuration
Electron configurations – arrangements
of e- in atoms
Rules:
Aufbau Principle – an e- occupies the lowest
energy first
Hund’s Rule – each orbital is filled with 1efirst and then the 2nd e- will fill
Pauli Exclusion Principle – no 2 e- in the
same atom can have the same set of QN
14:30-18:25
Electron Configuration
Representing electron configurations
Use the periodic table to write!
Know the s,p,d,f block and then let your
fingers do the walking!
Electron Configuration
Representing Electron
Configurations
Three Notations
Orbital Notation
Electron Configuration Notation
Electron Dot Notation
Orbital Notation
Uses a series of lines and arrows to
represent electrons
Examples
Orbital Notation
More examples
Electron Configuration
Notation
Eliminates lines and arrows; adds
superscripts to sublevels to represent
electrons
Long form examples
Electron Configuration
Notation
Short form examples – “noble gas
configuration”
Electron Dot Notation
Outer shell e Inner shell e Highest occupied energy level / highest
principle quantum number
Valence electrons – outermost e Examples
Electron Dot Notation
More examples
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