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
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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
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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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