Transcript Quantum Mechanics 1 - University of Birmingham

CHM2S1-A
Introduction to Quantum Mechanics
Dr R. L. Johnston
I: Foundations of Quantum Mechanics
THE UNIVERSITY
OF BIRMINGHAM
1. Classical Mechanics
1.1 Features of classical mechanics.
1.2 Some relevant equations in classical mechanics.
1.3 Example – The 1-Dimensional Harmonic Oscillator
1.4 Experimental evidence for the breakdown of classical mechanics.
1.5 The Bohr model of the atom.
2. Wave-Particle Duality
2.1 Waves behaving as particles.
2.2 Particles behaving as waves.
2.3 The De Broglie Relationship.
3. Wavefunctions
3.1 Definitions.
3.2 Interpretation of the wavefunction..
3.3 Normalization of the wavefunction.
3.4 Quantization of the wavefunction
3.5 Heisenberg’s Uncertainty Principle.
4. Wave Mechanics
4.1 Operators and observables.
4.2 The Schrödinger equation.
4.3 Particle in a 1-dimensional box.
4.4 Further examples.
Learning Objectives
•
•
•
•
•
•
•
•
•
To appreciate the differences between Classical (CM) and Quantum
Mechanics (QM).
To know the failures in CM that led to the development of QM.
To know how to interpret the wavefunction and how to normalize it.
To appreciate the origins and implications of quantization and the
uncertainty principle.
To understand wave-particle duality and know the relationships
between momentum, frequency, wavelength and energy for “particles”
and “waves”.
To be able to write down the Schrödinger equation for particles: in a 1D box; in 1- and 2-electron atoms; in 1- and 2-electron molecules.
To know the origins and allowed values of atomic quantum numbers
and how the energies and angular momenta of hydrogen atomic
orbitals depend on them.
To be able to sketch the angular and radial nodal properties of atomic
orbitals.
To appreciate the origins of sheilding and its effect on the ordering of
orbital energies in many-electron atoms.
•
•
•
•
•
•
To use the Aufbau Principle, the Pauli Principle and Hund’s Rule to
predict the lowest energy electron configuration for many-electron
atoms.
To appreciate how the Born-Oppenheimer approximation can be used
to separate electronic and nuclear motion in molecules.
To understand how molecular orbitals (MOs) can be generated as
linear combinations of atomic orbitals and the difference between
bonding and antibonding orbitals.
To be able to sketch MOs and their corresponding electron densities.
To construct MO diagrams for homonuclear and heteronuclear diatomic
molecules.
To predict the electron configurations for diatomic molecules, calculate
bond orders and relate these to bond lengths, strengths and vibrational
frequencies.
References
Fundamentals
• P. W. Atkins, J. de Paula, Atkins' Physical Chemistry (7th edn.),
OUP, Oxford, 2001.
• D. O. Hayward, Quantum Mechanics for Chemists (RSC Tutorial
Chemistry Texts 14) Royal Society of Chemistry, 2002.
• W. G. Richards and P. R. Scott, Energy Levels in Atoms and
Molecules (Oxford Chemistry Primers 26) OUP, Oxford, 1994.
More Advanced
• P. W. Atkins and R. S. Friedman, Molecular Quantum Mechanics
(3rd edn.) OUP, Oxford, 1997.
• P. A. Cox, Introduction to Quantum Theory and Atomic Structure
(Oxford Chemistry Primers 37) OUP, Oxford, 1996.
1. Classical Mechanics
• Do the electrons in atoms and molecules obey Newton’s
classical laws of motion?
• We shall see that the answer to this question is “No”.
• This has led to the development of Quantum Mechanics – we
will contrast classical and quantum mechanics.
1.1 Features of Classical Mechanics (CM)
1) CM predicts a precise trajectory for a particle.
velocity v
position r = (x,y,z)
• The exact position (r)and velocity (v) (and hence the momentum
p = mv) of a particle (mass = m) can be known simultaneously
at each point in time.
• Note: position (r),velocity (v) and momentum (p) are vectors,
having magnitude and direction  v = (vx,vy,vz).
2) Any type of motion (translation, vibration, rotation) can have any
value of energy associated with it
– i.e. there is a continuum of energy states.
3) Particles and waves are distinguishable phenomena, with different,
characteristic properties and behaviour.
Particles
Waves
Property
Behaviour
mass
position
velocity
momentum
collisions
wavelength
frequency


diffraction
interference
1.2 Revision of Some Relevant Equations in CM
Total energy of particle:
E = Kinetic Energy (KE) + Potential Energy (PE)
T - depends on v
V - depends on r
V depends on the system
e.g. positional, electrostatic PE
E = ½mv2 + V

Note:
E = p2/2m + V
(p = mv)
strictly E, T, V (and r, v, p) are all defined at a particular
time (t) – E(t) etc..
• Consider a 1-dimensional system (straight line translational
motion of a particle under the influence of a potential acting
parallel to the direction of motion):
• Define:
position
velocity
momentum
r=x
v = dx/dt
p = mv = m(dx/dt)
PE
force
V
F = (dV/dx)
• Newton’s 2nd Law of Motion
F = ma = m(dv/dt) = m(d2x/dt2)
acceleration
• Therefore, if we know the forces acting on a particle we can
solve a differential equation to determine it’s trajectory {x(t),p(t)}.
1.3 Example – The 1-Dimensional Harmonic Oscillator
x=0
F
k
m
NB – assuming no friction or
other forces act on the particle
(except F).
x
• The particle experiences a restoring force (F) proportional to its
displacement (x) from its equilibrium position (x=0).
• Hooke’s Law
F = kx
k is the stiffness of the spring (or stretching force constant of the
bond if considering molecular vibrations)
k
• Substituting F into Newton’s 2nd Law we get:
m(d2x/dt2) = kx
a (second order) differential
equation
Solution:
position
of particle
 = /2 =
frequency
(of oscillation)
Note:
ω
x(t) = Asin(t)
1
2π
k
m
k
m
frequency depends only on characteristics of the system
(m,k) – not the amplitude (A)!
x

time period  = 1/ 
+A
t
A
• Assuming that the potential energy V = 0 at x = 0, it can be
shown that the total energy of the harmonic oscillator is given
by:
E = ½kA2
• As the amplitude (A) can take any value, this means that the
energy (E) can also take any value – i.e. energy is continuous.
• At any time (t), the position {x(t)} and velocity {v(t)} can be
determined exactly – i.e. the particle trajectory can be specified
precisely.
• We shall see that these ideas of classical mechanics fail when
we go to the atomic regime (where E and m are very small) –
then we need to consider Quantum Mechanics.
• CM also fails when velocity is very large (as v  c), due to
relativistic effects.
1.4 Experimental Evidence for the Breakdown of Classical Mechanics
• By the early 20th century, there were a number of experimental
results and phenomena that could not be explained by classical
mechanics.
a) Black Body Radiation (Planck 1900)
“UV Catastrophe”
Energy
Radiated
Classical Mechanics
(Rayleigh-Jeans)
2000 K
1750 K
1250 K
0
2000
4000
6000
l/nm
Planck’s Quantum Theory
• Planck (1900) proposed that the light energy emitted by the
black body is quantized in units of h ( = frequency of light).
E = nh
(n = 1, 2, 3, …)
• High frequency light only emitted if thermal energy kT  h.
• h – a quantum of energy.
• Planck’s constant
h ~ 6.6261034 Js
• If h  0 we regain classical mechanics.
• Conclusions:
• Energy is quantized (not continuous).
• Energy can only change by well defined amounts.
b) Heat Capacities (Einstein, Debye 1905-06)
• Heat capacity – relates rise in energy of a material with its rise in
temperature:
CV = (dU/dT)V
• Classical physics

CV,m = 3R (for all T).
• Experiment

CV,m < 3R (CV as T).
• At low T, heat capacity of solids determined by
vibrations of solid.
• Einstein and Debye adopted Planck’s hypothesis.
• Conclusion: vibrational energy in solids is quantized:
– vibrational frequencies of solids can
only have certain values ()
– vibrational energy can only change
by integer multiples of h.
c) Photoelectric Effect (Einstein 1905)
h
e
Photoelectrons ejected with
kinetic
- energy:
eP
h
o
te
le
c
tro
n
s
Ek = h - F
Metal surface
work function = F
•
•
•
•
•
Ideas of Planck applied to electromagnetic radiation.
No electrons are ejected (regardless of light intensity) unless 
exceeds a threshold value characteristic of the metal.
Ek independent of light intensity but linearly dependent on .
Even if light intensity is low, electrons are ejected if  is above the
threshold. (Number of electrons ejected increases with light
intensity).
Conclusion:
Light consists of discrete packets (quanta) of
energy = photons (Lewis, 1922).
d) Atomic and Molecular Spectroscopy
•
It was found that atoms and molecules absorb and emit light only at
specific discrete frequencies spectral lines (not continuously!).
•
e.g. Hydrogen atom emission spectrum (Balmer 1885)
n1 = 1  Lyman
n1 = 2  Balmer
n1 = 3  Paschen
n1 = 4  Brackett
n1 = 5  Pfund
 1

ν 1
1
ν    RH  2  2 
n

c λ
n
2 
 1
•
Empirical fit to spectral lines (Rydberg-Ritz): n1, n2 (> n1) = integers.
•
Rydberg constant RH = 109,737.3 cm-1 (but can also be expressed
in energy or frequency units).
Revision: Electromagnetic Radiation
A – Amplitude
l – wavelength
 - frequency
c =  x lor  = c / l
wavenumber  =  /c = 1 / l
c (velocity of light in vacuum) = 2.9979 x 108 m s-1.
1.5 The Bohr Model of the Atom
• The H-atom emission spectrum was rationalized by Bohr (1913):
– Energies of H atom are restricted to certain discrete values
(i.e. electron is restricted to well-defined circular orbits,
labelled by quantum number n).
– Energy (light) absorbed in discrete amounts (quanta =
photons), corresponding to differences between these
restricted values:
e
E = E2  E1 = h
n2
n1
p+
E2
E2
h
E1
E1
h
Absorption
Emission
• Conclusion: Spectroscopy provides direct evidence for quantization of
energies (electronic, vibrational, rotational etc.) of atoms and molecules.
Limitations of Bohr Model & Rydberg-Ritz Equation
•
The model only works for hydrogen (and other one electron
ions) – ignores e-e repulsion.
•
Does not explain fine structure of spectral lines.
•
Note: The Bohr model (assuming circular electron orbits) is
fundamentally incorrect.
2. Wave-Particle Duality
•
Remember: Classically, particles and waves are distinct:
– Particles – characterised by position, mass, velocity.
– Waves – characterised by wavelength, frequency.
•
By the 1920s, however, it was becoming apparent that
sometimes matter (classically particles) can behave like waves
and radiation (classically waves) can behave like particles.
2.1 Waves Behaving as Particles
a) The Photoelectric Effect
Electromagnetic radiation of frequency  can be thought of as
being made up of particles (photons), each with energy E = h .
This is the basis of Photoelectron Spectroscopy (PES).
b) Spectroscopy
Discrete spectral lines of atoms and molecules correspond to
the absorption or emission of a photon of energy h , causing
the atom/molecule to change between energy levels: E = h .
Many different types of spectroscopy are possible.
c) The Compton Effect (1923)
•
Experiment: A monochromatic beam of X-rays (li) = incident on
a graphite block.
•
Observation: Some of the X-rays passing through the block are
found to have longer wavelengths (ls).
Intensity
ls
li

l
li
ls
• Explanation: The scattered X-rays undergo elastic collisions with
electrons in the graphite.
– Momentum (and energy) transferred from X-rays to electrons.
• Conclusion: Light (electromagnetic radiation) possesses momentum.
• Momentum of photon
p = h/l
• Energy of photon
E = h = hc/ l
• Applying the laws of conservation
of energy and momentum we get:
li
p=h/ls
ls

e
p=mev
 h 
1  cos 
Δλ  λ s  λ i   

m
c
 e 
2.2 Particles Behaving as Waves
Electron Diffraction (Davisson and Germer, 1925)
Davisson and Germer showed that
a beam of electrons could be diffracted
from the surface of a nickel crystal.
Diffraction is a wave property – arises
due to interference between scattered
waves.
This forms the basis of electron
diffraction – an analytical technique for
determining the structures of molecules,
solids and surfaces (e.g. LEED).
NB: Other “particles” (e.g. neutrons,
protons, He atoms) can also be
diffracted by crystals.
2.3 The De Broglie Relationship (1924)
• In 1924 (i.e. one year before Davisson and Germer’s
experiment), De Broglie predicted that all matter has wave-like
properties.
• A particle, of mass m, travelling at velocity v, has linear
momentum p = mv.
• By analogy with photons, the associated wavelength of the
particle (l) is given by:
h
h
λ 
p mv
3. Wavefunctions
•
•
A particle trajectory is a classical concept.
In Quantum Mechanics, a “particle” (e.g. an electron) does not
follow a definite trajectory {r(t),p(t)}, but rather it is best described
as being distributed through space like a wave.
3.1 Definitions
•
•
Wavefunction () – a wave representing the spatial distribution of a
“particle”.
e.g. electrons in an atom are described by a wavefunction centred
on the nucleus.
•  is a function of the coordinates defining the position of the
classical particle:
(x)
– 3-D
(x,y,z) = (r) = (r,,) (e.g. atoms)
•  may be time dependent – e.g. (x,y,z,t)
– 1-D
The Importance of 
•  completely defines the system (e.g. electron in an atom or
molecule).
• If  is known, we can determine any observable property (e.g.
energy, vibrational frequencies, …) of the system.
• QM provides the tools to determine  computationally, to
interpret  and to use  to determine properties of the system.
3.2 Interpretation of the Wavefunction
• In QM, a “particle” is distributed in space like a wave.
• We cannot define a position for the particle.
• Instead we define a probability of finding the particle at any point
in space.
The Born Interpretation (1926)
“The square of the wavefunction at any point in space is
proportional to the probability of finding the particle
at that point.”
•
Note: the wavefunction () itself has no physical meaning.
1-D System
• If the wavefunction at point x is (x), the probability of finding
the particle in the infinitesimally small region (dx) between x and
x+dx is:
P(x)  (x)2 dx
probability density
• (x) – the magnitude of  at point x.
Why write 2 instead of 2 ?
• Because  may be imaginary or complex  2 would be
negative or complex.
• BUT: probability must be real and positive (0  P  1).
• For the general case, where  is complex ( = a + ib) then:
2 = *
where * is the complex conjugate of .
(* = a – ib)
(NB i   1 )
3-D System
• If the wavefunction at r = (x,y,z) is (r), the probability of finding
the particle in the infinitesimal volume element d (= dxdydz) is:
P(r)  (r)2 d
• If (r) is the wavefunction describing
the spatial distribution of an electron
in an atom or molecule, then:
(r)2 = (r) – the electron density at point r
3.3 Normalization of the Wavefunction
• As mentioned above, probability:
P(r)  (r)2 d
• What is the proportionality constant?
• If  is such that the sum of (r)2 at all points in space = 1, then:
P(x) = (x)2 dx
1-D
P(r) = (r)2 d
3-D
• As summing over an infinite number of infinitesimal steps = integration,
this means:
2

Ptotal 1D    ψx  dx  1


2
  
2
Ptotal 3D    ψr  dτ     ψx, y, z  dxdydz  1

  
• i.e. the probability that the particle is somewhere in space = 1.
• In this case,  is said to be a normalized wavefunction.
How to Normalize the Wavefunction
• If  is not normalized, then:
2
 ψr  dτ  A  1
• A corresponding normalized wavefunction (Norm) can be
defined:
1
ψ Norm r  
ψr 
A
2
such that:
 ψ Norm r  dτ  1
• The factor (1/A) is known as the normalization constant
(sometimes represented by N).
3.4 Quantization of the Wavefunction
The Born interpretation of  places restrictions
on the form of the wavefunction:
(a)  must be continuous (no breaks);
(b) The gradient of  (d/dx) must be
continuous (no kinks);
(c)  must have a single value at any point in
space;
(d)  must be finite everywhere;
(e)  cannot be zero everywhere.
• Other restrictions (boundary conditions) depend on the exact system.
• These restrictions on  mean that only certain wavefunctions and  only
certain energies of the system are allowed.
 Quantization of   Quantization of E
3.5 Heisenberg’s Uncertainty Principle
“It is impossible to specify simultaneously, with precision, both the momentum
and the position of a particle*”
(*if it is described by Quantum Mechanics)
Heisenberg (1927)
px.x  h / 4
x
px
(or /2, where  = h/2).
– uncertainty in position
– uncertainty in momentum (in the x-direction)
•
If we know the position (x) exactly, we know nothing about momentum (px).
•
If we know the momentum (px) exactly, we know nothing about position (x).
•
i.e. there is no concept of a particle trajectory {x(t),px(t)} in QM (which applies to
small particles).
•
NB – for macroscopic objects, x and px can be very small when compared
with x and px  so one can define a trajectory.
•
Much of classical mechanics can be understood in the limit h  0.
How to Understand the Uncertainty Principle
• To localize a wavefunction () in space (i.e. to specify the
particle’s position accurately, small x) many waves of
different wavelengths (l) must be superimposed  large px
(p = h/l).
2 ~ 1
• The Uncertainty Principle imposes a fundamental (not
experimental) limitation on how precisely we can know (or
determine) various observables.
• Note – to determine a particle’s position accurately requires use
of short radiation (high momentum) radiation. Photons colliding
with the particle causes a change of momentum (Compton
effect)  uncertainty in p.

The observer perturbs the system.
• Zero-Point Energy (vibrational energy Evib  0, even at T = 0 K)
is also a consequence of the Uncertainty Principle:
– If vibration ceases at T = 0, then position and momentum
both = 0 (violating the UP).
• Note: There is no restriction on the precision in simultaneously
knowing/measuring the position along a given direction (x) and
the momentum along another, perpendicular direction (z):
 pz.x = 0
• Similar uncertainty relationships apply to other pairs of
observables.
• e.g. the energy (E) and lifetime () of a state:
E.  
(a) This leads to “lifetime broadening” of spectral lines:
– Short-lived excited states ( well defined, small ) possess
large uncertainty in the energy (large E) of the state.
 Broad peaks in the spectrum.
(b) Shorter laser pulses (e.g. femtosecond, attosecond lasers) have
broader energy (therefore wavelength) band widths.
(1 fs = 1015 s, 1 as = 1018 s)
4. Wave Mechanics
•
Recall – the wavefunction () contains all the information we need to
know about any particular system.
•
How do we determine  and use it to deduce properties of the
system?
4.1 Operators and Observables
•
If  is the wavefunction representing a system, we can write:
ˆ ψ  Qψ
Q
where
Q – “observable” property of system (e.g. energy,
momentum, dipole moment …)
ˆ – operator corresponding to observable Q.
Q
• This is an eigenvalue equation and can be rewritten as:
ˆ ψ  Q  ψ
Q
operator Q acting on
function 
(eigenfunction)
function  multiplied
by a number Q
(eigenvalue)
(Note:  can’t be cancelled).
Examples:
d/dx (eax) = a eax
d2/dx2 (sin ax) = a2 sin ax
To find  and calculate the properties (observables) of a system:
ˆ
1. Construct relevant operator Q
ˆ ψ  Qψ
2. Set up equation Q
3. Solve equation  allowed values of  and Q.
Quantum Mechanical Position and Momentum Operators
1. Operator for position in the x-direction is just multiplication by x
ˆxψ  xψ
2. Operator for linear momentum in the x-direction:
ˆ xψ  pxψ
p

 dψ

 pxψ
i dx
(solve first order differential equation   , px).
 d
ˆp x   
 i  dx
Constructing Kinetic and Potential Energy QM Operators
1. Write down classical expression in terms of position and momentum.
2. Introduce QM operators for position and momentum.
Examples
1. Kinetic Energy Operator in 1-D
px2
CM Tx 
2m
2. KE Operator in 3-D
CM
2
p
T

2m

ˆ
T
px2  py2  pz2
2m
ˆ
T
x
QM
2
2 
2
ˆ
p

d
x
ˆ 

T

x
2m
2m  dx 2
QM
2
2  2
2
2
ˆ
p




ˆ 

T

 2  2
2

2m
2m  x
y
z
partial derivatives
operate on (x,y,z)
ˆ (a function of position)
3. Potential Energy Operator V
 PE operator corresponds to multiplication by V(x), V(x,y,z) etc.




“del-squared”
2



2

2m

4.2 The Schrödinger Equation (1926)
• The central equation in Quantum Mechanics.
• Observable = total energy of system.
Schrödinger Equation
where
ˆ T
ˆ V
ˆ
H
ˆ ψ  Eψ
H
ˆ
H
Hamiltonian Operator
E Total Energy
and
E = T + V.
• SE can be set up for any physical system.
ˆ depends on the system.
• The form of H
• Solve SE   and corresponding E.
Examples
1. Particle Moving in 1-D
ˆ ψT
ˆ ψV
ˆ ψ  Eψ
H
(x)
 2   2 ψ 

 Vx ψ  Eψ
2


2m  x 
• The form of V(x) depends on the physical situation:
– Free particle
V(x) = 0 for all x.
– Harmonic oscillator
V(x) = ½kx2
2. Particle Moving in 3-D
• SE 
or
(x,y,z)
 2   2 ψ  2 ψ  2 ψ 

 2  2  Vx, y, z ψ  Eψ
2

2m  x
y
z 
2 2

 ψ  Vx, y, z ψ  Eψ
2m
Note: The SE is a second order
differential equation
4.3 Particle in a I-D Box
System
– Particle of mass m in 1-D box of length L.
– Position x = 0L.
– Particle cannot escape from box as PE V(x)=  for x = 0, L (walls).
– PE inside box: V(x)= 0 for 0< x < L.


1-D Schrödinger Eqn.
 2   2 ψ 

 Eψ
2
2m  x 
PE (V)
(V = 0 inside box).
0
0
x
L
 2   2 ψ 

 Eψ
2


2m  x 
• This is a second order differential equation – with general
solutions of the form:
 = A sin kx + B cos kx
 2ψ 

  k 2 A sin kx  B cos kx   k 2 ψ
 x 2 


• SE 



 2   2 ψ    2 
2





k
ψ  Eψ
2




2m  x   2m 
k 2 2
E
2m
(i.e. E depends on k).
Restrictions on 
• In principle Schrödinger Eqn. has an infinite number of solutions.
• So far we have general solutions:
– any value of {A, B, k}  any value of {,E}.
• BUT – due to the Born interpretation of , only certain values of 
are physically acceptable:
– outside box (x<0, x>L) V =  
  2 = 0

impossible for particle
to be outside the box
 = 0 outside box.
–  must be a continuous function

Boundary Conditions
 = 0 at x = 0
 = 0 at x = L.
Effect of Boundary Conditions
 = A sin kx + B cos kx = B
1. x = 0
0

=0 B=0
 = A sin kx
for all x
 = A sin kL = 0
2. x = L
A=0 ? 
(or  = 0 for all x)
sin kL = 0
1

kL = n
sin kL = 0 ?

n = 1, 2, 3, …
(n  0, or  = 0 for all x)
Allowed Wavefunctions and Energies
•
k is restricted to a discrete set of values:
•
Allowed wavefunctions:
•
•
Normalization:
Allowed energies:

n = A sin(nx/L)
A = (2/L) 
ψn 
k 2 2 n 2 π 2 2
En 

2m
2mL 2
En 
k = n/L
n 2h2
8mL 2
nx 

 L 

2
sin
L
Quantum Numbers
•
There is a discrete energy state (En),
corresponding to a discrete wavefunction
(n), for each integer value of n.
•
Quantization – occurs due to boundary
conditions and requirement for  to be
physically reasonable (Born interpretation).
•
n is a Quantum Number – labels each
allowed state (n) of the system and
determines its energy (En).
•
Knowing n, we can calculate n and En.
Properties of the Wavefunction
ψn 
nx 

 L 

2
sin
L
•
Wavefunctions are standing waves:
•
The first 5 normalized wavefunctions for the particle in the 1-D
box:
•
Successive functions possess one more half-wave ( they have a
shorter wavelength).
•
Nodes in the wavefunction – points at which n = 0 (excluding the
ends which are constrained to be zero).
•
Number of nodes = (n-1)
1  0; 2  1; 3  2 …
Curvature of the Wavefunction
• If y = f(x)
• In QM
dy/dx = gradient of y (with respect to x).
d2y/dx2 = curvature of y.
Kinetic Energy  curvature of 
• Higher curvature  (shorter l)  higher KE
• For the particle in the 1-D box (V=0):
KE
  2ψn
E n  Tn  
 x 2

 n2

 L2

 2ψ

T   2  ..... 
 x



Energies
•
En 
n 2h2
8mL 2
En  n2/L2

En as n (more nodes in n)
En as L (shorter box)
n (or L)


curvature of n
KE  En
E 
2
E
node
L
1
L1
L2
•
En  n2

n
3
2
1
•
energy levels get further apart as n
E
E3 
E2 
9h 2
8mL 2
4h 2
8mL 2
h2
E1 
8mL 2
0
•
8mL 2
Zero-Point Energy (ZPE) – lowest energy of particle in box:
ZPE  E min  E 1 
•
ZPE 
h2
CM Emin = 0
h2
8mL 2
QM E = 0 corresponds to  = 0 everywhere (forbidden).
•
If V(x) = V  0, everywhere in box, all energies are shifted by V.
V0
V=0
En 
2 2
n h
8mL
2
V
E2=E2+V
E2
E1
V
E1=E1+V
Density Distribution of the Particle in the 1-D Box
•
The probability of finding the particle
between x and x+dx (in the state
represented by n) is:
Pn(x) = n(x)2 dx = (n(x))2 dx (n is real)

•
Pn x  
2
L
2  nx

sin 
dx
 L 
Note: probability is not uniform
– n 0 at walls (x = 0, L) for all n.
– n2 = 0 at nodes (where n = 0).
2=

2
2
4.4 Further Examples
(a) Particle in a 2-D Square or 3-D Cubic Box
• Similar to 1-D case, but   (x,y) or (x,y,z).
• Solutions are now defined by 2 or 3 quantum numbers
e.g. [n,m, En,m]; [n,m,l, En,m,l].
• Wavefunctions can be represented as contour plots in 2-D
(b) Harmonic Oscillator
• Similar to particle in 1-D box, but PE V(x) = ½kx2
(c) Electron in an Atom or Molecule
ˆ
T
ˆ
V
3-D KE operator
PE due to electrostatic interactions between electron and all
other electrons and nuclei.