CHEMISTRY 161

Chapter 8

(n, l, ml, ms) ATOMIC ORBITALS

n l ml orbitals designation

1 0 0 1 1s

2 0 0 1 2s

1 -1,0,+1 3 2px,2py,2pz

3 0 0 1 3s

1 -1,0,+1 3 3px,3py,3pz

2 -2,-1,0,+1,+2 5 3dxy,3dyz,3dxz,

3dx2-y2,3dz2

4 … … … …

H Atom Orbital Energies

energy level diagram H atom

3s 3p 3d

2s 2p

1s

E

energy depends only on principal quantum number

orbitals with same n but different l are degenerate

1s

E

2s2p

3s3p

3d4s

4p5s

4d

MULTI-ELECTRON ATOM

orbitals with same n and different l are not degenerate

energy depends on n and ml

EXAMPLES [Xe]

EXP I

Periodic Table of the Elements

period

group

chemical reactivity - valence electrons

ns1 ns2

ns2np6

ns2(n-1)dx

PERIODIC TRENDS

3. IONIZATION ENERGIES

4. ELECTRON AFFINITIES

1. ATOMIC RADIUS

2. IONIC RADIUS

ATOMIC RADIUS

MAIN GROUPS

AT

OM

IC R

AD

IUS

AT

OM

IC R

AD

IUS

EXP II

1s, 2s, 3s

3s2s1s

2px, 3px, 4px

4px3px2px

ATOMIC RADIUSMAIN GROUPS

ATOMIC RADIUS

effective nuclear charge

(shielding s vs p orbitals)

IONIC RADII

IONIC RADIUS

ION

IC R

AD

IUS

ION

IC R

AD

IUS

cations are smaller than their atoms

anions are larger than their atoms

Na is 186 pm and Na+ is 95 pm

F is 64 pm and F- is 133 pm

same nuclear charge and repulsion among electrons increases radius

one less electron electrons pulled in by nuclear charge

O < O– < O2–

EXAMPLES

Which is bigger?

Na or Rb Rb higher n, bigger orbitals

K or Ca K poorer screening for Ca

Ca or Ca2+ Ca bigger than cation

Br or Br- Br smaller than anion

QUESTIONThe species F-, Na+,Mg2+ have relative sizes in the order

1 F-< Na+<Mg2+ 2 F-> Na+>Mg2+

3 Na+>Mg2+> F- 4 Na+=Mg2+= F-

5 Mg2+> Na+>F-

QUESTION

1 F-< Na+<Mg2+

2 F-> Na+>Mg2+

3 Na+>Mg2+> F-

4 Na+=Mg2+= F-

5 Mg2+> Na+>F-

Na+ is 95 pm

Mg2+ is 66 pm

F- is 133 pm

ALL 1s22s22p6

ALL are isoelectronic

3. IONIZATION ENERGIES

M(g) M+(g) + e-

energy required to remove an electron from a gas phase atom in its electronic ground state

I1 > 0

first ionization energy(photon)

M+(g) M2+(g) + e-

M2+(g) M3+(g) + e-

second ionization energy

third ionization energy

I2 > 0

I3 > 0

I1 > I2 > I3

Why?

electrons closer to nucleus more tightly held

ION

IZA

TIO

N E

NE

RG

Y

ION

IZA

TIO

N E

NE

RG

Y

first ionization energies decrease

d shell insertion

I E AZ

neff. . 2

2

IONIZATION ENERGY

0

500

1000

1500

2000

2500

0 1 2 3 4 5 6 7 8 9

GROUP NUMBER

ION

IZA

TIO

N E

NE

RG

Y(k

J/m

ol)

1 2 13 14 15 16 17 18

n=1

n=2

n=3

n=4

1. closed shells are energetically most stable

2. half-filled shells are energetically very stable

DERIVATION OF IONIZATION ENERGIES

noble gases have the highest ionization energy

4. ELECTRON AFFINITIES

the energy change associated with the addition

of an electron to a gaseous atom

X(g) + e– X–(g)

electron affinity can be positive or negative

Why?

EL

EC

TR

ON

AF

FIN

ITY

EL

EC

TR

ON

AF

FIN

ITY

general trend

-200

-100

0

100

200

300

400

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19

ATOMIC NUMBER

-ELE

CT

RO

N A

FF

INIT

Y

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

Al

Si

P

S

Cl

Ar

1. closed shells are energetically most stable

2. half-filled shells are energetically very stable

DERIVATION OF

ELECTRON AFFINITIES