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THE UNIVERSITY OF BIRMINGHAM. CHM2S1-A Introduction to Quantum Mechanics Dr R. L. Johnston. II: Quantum Mechanics of Atoms and Molecules 5. Electronic Structures of Atoms 5.1 The hydrogen atom and hydrogenic ions. 5.2 Quantum numbers. 5.3 Orbital angular momentum.

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chm2s1 a introduction to quantum mechanics dr r l johnston

THE UNIVERSITY

OF BIRMINGHAM

CHM2S1-AIntroduction to Quantum MechanicsDr R. L. Johnston

II: Quantum Mechanics of Atoms and Molecules

5.Electronic Structures of Atoms

5.1 The hydrogen atom and hydrogenic ions.

5.2 Quantum numbers.

5.3 Orbital angular momentum.

5.4 Atomic orbitals.

5.5 Electron spin.

5.6 Properties of atomic orbitals.

5.7 Many-electron atoms.

slide2
6. Bonding in Molecules

6.1The Born-Oppenheimer approximation.

6.2 Potential energy curves.

6.3 Molecular orbital (MO) theory.

6.4 MO diagrams.

6.5 MOs for 2nd row diatomic molecules.

6.6 Molecular electronic configurations.

6.7 Bond order.

6.8 Paramagnetic molecules.

6.9 Heteronuclear diatomic molecules.

5 electronic structures of atoms
5. Electronic Structures of Atoms

5.1 The Hydrogen Atom and Hydrogenic Ions

  • The series of atoms/ions H, He+, Li2+, Be3+ …

all have: 1 electron (charge –e)

nucleus (charge+Ze)

  • To determine the electronic wavefunction ( = e) and allowed electronic energy levels (E), we must set up and solve the Schrödinger Equation for a single electron in an atom.
slide4

-e

me

+Ze

mN

  • Consider 3-D motion of the electron (-e) relative to the nucleus (+Ze):

Hamiltonian Operator

Kinetic Energy where

Potential Energy where

(electrostatic attraction between electron and nucleus).

r

slide5

angular

wavefunction

(spherical harmonic)

radial

wavefunction

  • We must solve the 2nd order differential equation:
  • Because of the spherical symmetry of the atom, it is convenient to describe the position of the electron in spherical polar coordinates(r,,), rather than Cartesians (x,y,z).
  • Due to this spherical symmetry, the wavefunction can be separated into a product of radial and angular components:

(r,,)=R(r).Y(,)

  • Imposing boundary conditions  3 quantum numbers(n,,m)

 n,,m(r,,)=Rn,(r).Y,m(,)

slide6
5.2 Quantum Numbers

n,,m(r,,)=Rn,(r).Y,m(,)

  •  depends on 3 quantum numbers (n,,m).

1. Principal Quantum Number (n)

  • Positive integer n = 1, 2, 3 …
  • For H and 1-e ions, the electron energy depends only on n

(this is the same result as from the Bohr model).

  • For many-electron atoms E depends on n and .
slide7
2. Orbital Angular Momentum Quantum Number ()
  • For a given value of n,  can take the integer values:

 = 0, 1, 2 … (n-1)

e.g. n = 1 = 0

n = 2  = 0, 1

n = 3  = 0, 1, 2

3. Orbital Magnetic Quantum Number (m)

  • For a given value of , m can take the integer values:

m = 0, 1, 2 …  

e.g.  = 0m = 0

 = 1m = 0, 1

  • There are (2+1) allowed values of m.
slide8
5.3 Orbital Angular Momentum
  • Q.Nos.  andm arise from the angular part of the wavefunction and they relate to the angular momentum of the electron due to its motion around the nucleus (orbital angular momentum).
  • –determines magnitude of angular momentum vector (J)
  • Because  is quantized, so is |J|:
slide9

m

  • The orientation of the angular momentum vector (J) is also quantized – as it depends on m.
  • The component of angular momentum along a reference direction (e.g. the z-axis):

Jz = m

  • e.g. for  = 2 (m = 0, 1, 2)

angular momentum vector J

has magnitude |J| = 6and

5 allowed orientations:

Jz = 0, , 2.

slide10
5.4 Atomic Orbitals
  • The wavefunctions (n,,m) describing an electron in an atom are known as atomic orbitals.
  • Each atomic orbital (1-e wavefunction) is uniquely defined by 3 quantum numbers (n,,m).
  • The orbital () gives the spatial distribution of the electron (via the Born interpretation of ||2).
  • Orbitals are often drawn as 3-D surfaces which enclose approx. 90% of the probability of finding the electron.
slide11
Electronic Shells, Sub-shells and Orbitals
  • All orbitals with the same value of n together constitute an electronic shell.
  • For each n, orbitals with the same value of  together constitute an electronic sub-shell.
  • Each sub-shell consists of (2+1) orbitals, each with a different m value.
slide12
Possible combinations of shells and sub-shells.
  • Total number of orbitals in shell n isn2.
electronic shells sub shells and orbitals
Electronic Shells, Sub-shells and Orbitals

3d orbitals

l=2

ml=1

ml=2

ml=2

ml=1

ml=0

n=3

3p orbitals

l=1

ml=1

ml=1

ml=0

3s orbital

l=0

ml=0

shell

sub-shells

orbitals

atomic orbital energy diagram for h
Atomic Orbital Energy Diagram for H
  • Orbital energies depend only on the

principal quantum number n.

  • For a given n value (shell), all sub-shells

() and orbitals (m) have the same energy

– i.e. they are degenerate.

quantum numbers and atomic orbitals
Quantum Numbers and Atomic Orbitals
  • Quantum theory postulates that electrons in atoms are fixed in regions of space corresponding to atomic orbitals. The probability of finding an electron at a particular point in space (r) is proportional to |(r)|2 (Born Interpretation).
  • Orbitals differ from each other in size, shape and orientation:
    • the orbital size is defined by the principle quantum numbern
    • the type of orbital (it’s shape) is defined by the angular momentum quantum number
    • the orientation of the orbital is defined by the orbital magnetic quantum numberm
shapes of atomic orbitals

s orbitals have l = 0 and ml= 0

pz

px

py

  • d orbitals have l = 2 and ml= 2, 1, 0 ,1, 2

d x2 y2

dz2

dxy

dxz

dyz

Shapes of Atomic Orbitals
  • p orbitals have l = 1 and ml= 1 , 0, 1
slide17
5.5 Electron Spin
  • Stern and Gerlach(1921) observed that a beam of Ag atoms is split into 2 beams by an inhomogeneous magnetic field.
  • Dirac(1930) introduced relativistic effects into Quantum Mechanics and showed that to completely describe the state of an electron we must specify:

1. The orbital (n,,m)

2. The spin state of the electron

  • Electron spin is characterised by 2 quantum numbers:

– spin angular momentum q. no.s ( =½ for all electrons)

– spin magnetic q. no.ms ( = ½).

slide18
Spin angular momentum (S) has magnitude:
  • The projection of S on the z-axis:

Sz = ms = ½ 

  • There are two possible electron spin states:

ms = +½ “spin up”

ms = -½ “spin down”

  • Spin is an intrinsic property of the electron and is not connected with orbital motion.
  • Complete specification of an electron in an atom requires 4 quantum numbers (n,,m,ms) – as s is fixed.
slide19

+

5.6 Properties of Atomic Orbitals

1. Shape – determined by the angular wavefunction Y,m(,)

1s orbital (n = 1,  = 0, |J| = 0)

  • Normalized wavefunction:
  • Has no angular dependence

(spherically symmetric, depends only on r).

  • All s orbitals are spherically symmetrical.

2s orbital ns orbital

  • As n, orbitals expand – average radius r
slide20

+

-

+

-

-

+

2p orbitals (n = 2,  = 1, |J| = 2)

  • There are 3 degenerate orbitals, with 3 different m values (0,1).
  • e.g. m = 0  Jz = m  = 0  2pz orbital (pointing along z-axis)
  • m = 1  Jz =   2px, 2py orbitals
slide21
2. Nodal Properties

a) Angular Nodes

  • s orbitals – have no angular nodes (spherically symmetric)
  • p orbitals – have one angular node (a nodal plane, where  = 0)
  • No. of angular nodes =  (s = 0, p = 1, d = 2 …)

b) Radial Nodes

  • s orbitals – have a maximum in  at nucleus (r = 0)
  • For other orbitals (p, d, f …),  = 0 at nucleus due to angular nodes.
  • All orbitals decay exponentially at large r values.
  • Radial nodes – nodal surfaces where  = 0, with  changing sign either side of the node.
  • No. of radial nodes =n1 1s  0

2s  1 2p  0

3s  2 3p  1 3d  0

  • Total number of nodes (angular + radial) = n1
angular and radial nodes in

+

+

 = 0

-

-

+

 = 0

2pz orbital

3s orbital

Angular and Radial Nodes in 

Angular Nodes

Radial Nodes

slide23
3. Radial Distribution Functions
  • Electron density at point (r) in space:
  • Probability of finding electron at point r:
  • Radial Distribution Function – probability density of finding electron at distance r from nucleus:
  • P(r)dr – probability of finding electron in a shell of thickness dr at distance r from the nucleus.
  • For non-spherical orbitals (p, d, f …), use:
slide24

4r22

Example: H(1s)

  • Peak in RDF at r = a0 (the Bohr radius for n = 1).
  • The delocalized electron is represented by a wave which has maximum probability at r = a0.

a0

r

slide25
Hydrogenic Ions
  • 1e ions with nuclear charge Z.
  • Max. in 1s RDF at
  • As Z, orbitals contract (rmax).
  • Orbital Energies
  • As Z, orbitals more tightly bound (En – more negative).
slide26

-e(1)

r1

r12

Z = +2e

-e(2)

r2

5.7 Many-Electron Atoms

  • For atoms with 2 or more electrons, the Schrödinger Eqn. must include KE terms for all electrons and PE terms for all e-e and e-n interactions.
  • Example: He (2 electrons)

2 terms in

3 terms in

S.E.

where

slide27
For N electrons:N = (r1,r2,r3 … rN)
  • For > 1 e, S.E. cannot be solved analytically – though good numerical solutions may be obtained from computer calculations.

 Approximations must be made.

The Orbital Approximation

  • The total wavefunction (N) for the N-electron atom is approximated by the product of N 1-e orbitals similar to those of hydrogenic ions:

N = (r1,r2,r3 … rN) =(r1)(r2)(r3)…(rN)

slide28
Orbital Energies for Many-Electron Atoms
  • For H and hydrogenic ions, E depends only on n.
    • e.g. 3s, 3p and 3d orbitals are degenerate.
  • For many-electron atoms, E depends on n and .
  • In shell n, En as  i.e. E(ns) < E(np) < E(nd) …
  • Why ?
slide29
Effective Nuclear Charge
  • An electron at distance r from nucleus experiences an effective nuclear charge:

Zeff < Z

(Z = actual charge on nucleus = atomic number)

  • Electrons inside a sphere of radius r repel the electron and shield (screen) it from the nucleus:

Zeff = Z  

( = shielding constant).

slide30



A

-B

A

B

+

R

Properties of 

  • For large R,  behaves like 2 independent H(1s) AOs.
  • For small R, there is significant overlap between the AOs.
  • Destructive interference between atomic

wavefunctions (AOs) AandB depletion

of electron density between the nuclei

(decrease of  and 2).

  • The depletion of electron density between

the nuclei leads to decreased e-n attraction:

 E() > E(A,B)

  is an antibonding MO.

  •  is cylindrically symmetrical about the internuclear (A-B) axis

 labelled as a  MO (or sometimes *, as it is antibonding).

slide31
Electrons in orbitals with low n (core orbitals) have greater electron density (higher probability of finding the electron) close to the nucleus  experience higher Zeff  lower energy.
  • Electrons in orbitals with same n, but

lower , have peaks in RDF which are

closer to the nucleus  better at shielding

other electrons and better at penetrating

shielding:

Zeff ns > np > nd …

E ns < np < nd …

  • Orbitals with same n and  are still

degenerate

e.g. E(2px) = E(2py) = E(2pz)

slide32
Electronic Configurations
  • To determine the electronic configuration of an atom (how the electrons are distributed among the atomic orbitals), we need to know:
    • The available orbitals and their relative energies.
    • The number of electrons.
    • The rules for filling the orbitals.
slide33
Rules for Filling Orbitals

1. The Pauli Exclusion Principle

No two electrons in a particular atom or ion can have the same values of all 4 quantum numbers (n,,m,ms).

  • No more than 2e can occupy a given orbital
  • If 2e are in the same orbital (same n,,m), they must have opposite (= paired) spins.

H1s1

He1s2

Li 1s22s1

X

slide34
2. The Aufbau (Building-up) Principle
  • Add available electrons into orbitals – starting from the lowest in energy, put maximum of 2e in each orbital.
  • General order of orbital occupation:

1s < 2s < 2p < 3s < 3p < 4s ~ 3d < 4p < 5s < 4d …

  • Usually completely fill each sub-shell () before starting to fill another.
    • Exceptions – e.g. in the 1st row transition metals:

Cr[Ar]4s13d5

Cu[Ar]4s13d10

slide36

Variation of Orbital Energy with Z

The 3d and 4s orbitals are close in

energy for the first transition metal

series.

For heavier atoms (higher atomic

number Z), E(3d) < E(4s).

As Z, the  sub-shells of the inner

(lower n) orbitals become approx.

degenerate

slide37
3. Spatial Separation
  • Electrons occupy different orbitals (m) of a given sub-shell () before starting to pair electrons.

e.g. N[He]2s22p3 = [He]2s22px12py12pz1

  • Electrons in different orbitals on average are further apart lower e-e repulsion  lower E.
slide38
4. Hund’s Rule (of Maximum Multiplicity)
  • Provided rules 1 and 2 are satisfied, an atom in its ground state adopts the configuration with the maximum number ofunpaired (= parallel) spins.

e.g. C[He]2s22p2 = [He]2s22px12py1

(equivalent to … 2px12pz1 and 2py12pz1).

Parallel Spins Antiparallel Spins

  • Electrons with parallel spins () tend to avoid each other better than those with antiparallel spins ()  reduced e-e repulsion  lower E.
  • This is known as spin correlation.
6 bonding in molecules

e

rA

rB

RAB

HA

HB

PE

KE

nuclear K.E.

electron K.E.

e-n attraction

n-n repulsion

6. Bonding in Molecules

Exact solution of the Schrödinger Equation is not possible for any molecule – even the simplest molecule H2+.

Full Hamiltonian operator for H2+:

slide40
6.1 The Born-Oppenheimer Approximation
  • Nuclei are much heavier (thousands of times) than electrons.

 they move much more slowly.

  • In the Born-Oppenheimer approximation, nuclei are treated as being stationary.
  • Consider motion of electrons relative to fixed nuclei.
  • Total wavefunction () is split into the product of electronic and nuclear wavefunctions:

 =e.n

slide41

e

rA

rB

R

HA

HB

electron K.E.

e-n attraction

n-n repulsion

(constant)

Example 1: H2+

  • Fix RAB (  R).
  • Schrödinger Equation for electronic

motion:

  • Note: although the electronic SE is solved for fixed RAB(= R), the solutions (e,Ee) depend on the value of R.
slide42

e1

r1A

r1B

r12

R

HA

HB

r2A

r2B

e2

Example 2: H2

  • Electronic Hamiltonian:
  • SE for molecules containing other elements

(atomic no. Z):

    • e-n attraction terms  e2  Ze2
    • n-n repulsion terms +e2  +Z2e2.

electron K.E.

e-n attraction

n-n repulsion

e-e repulsion

slide43

V(R)

Re

R

0

-De

R1

R2

6.2 Potential Energy Curves

  • For a diatomic molecule, solving the electronic SE for different fixed positions of the nuclei (i.e. fixed inter-nuclear distances, R) gives the molecular potential energy curve V(R).
  • Re = equilibrium bond length
  • De = PE well depth = dissociation energy.
  • As R  , V(R)  0 (dissociation limit).
  • This can be extended to larger molecules:
    • e.g. for a general triatomic molecule we get a potential energy surfaceV(R1,R2,).
slide44
How can we determine  and E for an electron in a molecule?
  • SE can be solved exactly for (within the B-O approximation) but it is complicated and SE cannot be solved exactly for > 1e.
  • We need to make more approximations  Molecular Orbital theory.

6.3 Molecular Orbital Theory

  • Electrons in molecules have spatial distributions which are described by 1e wavefunctions called molecular orbitals (MOs) – analogous to atomic orbitals (AOs).
  • Let  represent a MO and  an AO.
  • MOs are spatially delocalised over the molecule.
  • Probability of finding electron at point r in space in MO :

P(r) = | (r)|2d

slide45
Linear Combination of Atomic Orbitals (LCAO) Approximation
  • Construct MOs () as linear combination of AOs ():
  • ci= coefficients (numbers) = contribution of ith AO to the MO.
  • N AOs  N MOs.
  • Justification
    • When electron is close to one nucleus (A) it experiences an electrostatic (Coulomb) attraction that is greater than that to B.
    •  MO wavefunction () close to A, resembles an atomic orbital centred on A (A).
slide46

R

A

B

A

B

Example: H2+ and H2

  • MOs formed as linear combinations of H(1s) AOs.
  • 2 AOs (A,B)  2 MOs (+,)

+ = N+(A+B) in-phase (bonding)

 = N(AB) out-of-phase (antibonding)

slide47

R

+

B

A

A

B

R

Properties of +

  • For large R, + behaves like 2 independent H(1s) AOs.
  • For small R, there is significant overlap between the AOs.
  • Constructive interference between atomic

wavefunctions (AOs) AandB build-up

of electron density between the nuclei

(increase of  and 2).

slide48

+

+

+

side view

cross section

  • The accumulation of electron density between the nuclei leads to increased e-n attraction (electrons interact strongly with both nuclei):

 E(+) < E(A,B)

 + is a bonding MO.

 Covalent bonding – due to sharing of electrons.

  • + is cylindrically symmetrical about the internuclear (A-B) axis

 labelled as a  MO:

slide49



node

-B

A

A

B

R

Properties of 

  • For large R,  behaves like 2 independent H(1s) AOs.
  • For small R, there is significant overlap between the AOs.
  • Destructive interference between atomic

wavefunctions (AOs) AandB deplete

of electron density between the nuclei

(decrease of  and 2).

slide50

nodal plane

+

  • The depletion of electron density between the nuclei leads to decreased e-n attraction.

 E() > E(A,B)

  is an antibonding MO.

  •  is cylindrically symmetrical about the internuclear (A-B) axis

 labelled as a * MO (* denotes antibonding character):

slide51
Normalization of the MO Wavefunction
  • What are the normalization constants (N+ and N)?

+ = N+(A+B)

 = N(AB)

  • Normalization condition:

where  = + or 

  • e.g. for +

atomic orbitals are normalized:

define overlap integral between orbitals AandB:

normalization constant:

slide52
Similarly, for  we get:
  • Since 0 < SAB < 1, this means that N > N+.
  • Similar arguments can be used to show that the antibonding orbital () is raised in energy by more than the bonding orbital (+) is lowered in energy when a bond is formed.
  • Note: if we ignore overlap, then N+ = N = 1/2
slide53

 = 2*

Energy

A=H(1sA)

B=H(1sB)

+ = 1

V(R)



R

Re

0

+

-De

6.4 Molecular Orbital Diagrams

Example 1. H2+

  • Ground state configuration: (+)1

1e in bonding orbital  bound state.

  • Excited state configuration: ()1

1e in antibonding orbital  unbound state.

  • Energies defined relative to dissociation

H2+ H + H+

slide54

 = 2*

Energy

A=H(1sA)

B=H(1sB)

+ = 1

Example 2. H2

  • Ground state configuration: (+)2

2e in bonding orbital  bound state.

  • H2 has a shorter stronger bond than H2+ (more bonding electrons).
  • Note: De(H2) < 2De(H2+) due to e-e repulsion.
slide55

 = 2*

Energy

A=He(1sA)

B=He(1sB)

+ = 1

Example 3. He2

  • Ground state configuration: (+)2()2

No net covalent bonding (bonding and a-b contributions cancel out).

Only weak dispersion forces hold He atoms together (see Intermolecular Forces lectures).

  • He2+ has the configuration (+)2()1and does have net covalent bonding.
slide56

+

+

+

+

A(2s)+B (2px)

A(2px)+B (2pz)

6.5 MOs for 2nd Row Diatomic Molecules

  • Valence AOs = 2s, 2px, 2py, 2pz
  • Core AOs = 1s (not involved in bonding)
  • Linear combinations of 2s orbitals:

 = N(A(2s)B (2s))  1(2s) and 2 *(2s) (as for H2)

  • Some combinations are not allowed – zero net overlap = “orthogonal orbitals”:
slide57

2*

4*

4*

2*

2p

2p

1

3

3

1

  • The 2p orbitals interact to give -type (0 angular nodes with respect to the molecular axis) and -type (1 angular node) MOs, which can be bonding (,) or antibonding (*,*).

2pz-2pz overlap >  overlap

 3-4* splitting > 1-2*.

slide58
6.6 Molecular Electronic Configurations
  • Follow same rules as for atomic electronic configurations (Aufbau principle, Hund’s rule etc.).
  • Note: the ordering of MOs can vary – e.g. the 3(2p) and 1(2p) MOs are sometimes reversed:

Due to 2s-2p mixing

(hybridization) which

raises 3 and lowers

2*.

As Z, the 2s-2p separation

increases, so s-p mixing

is weaker.

3 > 1(B2, C2, N2)

3 < 1(O2, F2)

6 7 bond order
6.7 Bond Order
  • The strength of a covalent bond is the net outcome of occupying bonding and antibonding orbitals.
  • Bond Order b = ½(NB N*)

NB = number of electrons in bonding MOs

N* = number of electrons in antibonding MOs

  • Examples
6 8 paramagnetic molecules

4*

2*

2p

2p

1

3

6.8 Paramagnetic Molecules
  • Even with even numbers of electrons, certain molecules are paramagnetic (i.e. they have unpaired electron spins).
  • e.g. O2
    • Ground state electronic configuration: (1)2 (2*)2 (3)2 (1)4 (2*)2
    • There are 2 electrons in the antibonding pair of 2* orbitals.
    • From Hund’s rule – the lowest energy configuration has the most unpaired spins = 2.
    • The magnetic effects of these 2 electrons do not cancel out.
6 9 heteronuclear diatomic molecules

HF

*(1s-2p)

(1s+2p)

non-bonding

F(2px,2py)

AOs.

6.9 Heteronuclear Diatomic Molecules
  • Generally AOs of different atoms have different energies – depending on relative electronegativities of the atoms.
  • The MO closest in energy to an

AO has more character

(greater LCAO coefficient) of

that AO.

  •  bonding and antibonding orbitals

usually have opposite characters.

  • e.g. HF

 = 0.19H(1s)+0.98F(2pz)

* = 0.98H(1s)0.19F(2pz)