Chapter 7
Electronic Configurations and
the Periodic Table
Robert W. Bunsen 1811-1899.
Established the science of
spectroscopy.
Niels H. Bohr 1885-1962.* Related
Periodic Table to atomic atomic
structure and spectral behavior.
1
The Wave Nature of Light
•Modern atomic theory arose out of studies of the
interaction of radiation with matter.
•Electromagnetic radiation moves through a vacuum
with a speed of 2.99792458  108 m/s.
•(We will usually round off to 3 x 108 m/s)
•Electromagnetic waves have characteristic wavelengths
(l) and frequencies (n) .
•Example: visible radiation has wavelengths between
• l = 400 nm (violet) and 750 nm (red).
2
The Wave Nature of Light
Some examples of EM radiation:
visible light
ultraviolet radiation (UV)
infrared radiation (IR)
radio waves
microwaves
x-rays
?sound waves?
3
The Wave Nature of Light
The wavelength, λ, is the distance
between crests in a wave
The frequency, n, of a wave is the
number of cycles which pass a point in
one second.
The speed of a wave, c, is given by its
frequency multiplied by its wavelength:
c = nl
where c = 3 x 108 m.s-1 (the “speed of
light”)
4
Short wavelength
High frequency
Long wavelength
Small frequency
as λ
as λ
ν
ν
5
The Wave Nature of Light
6
7
Quantized Energy and Photons
Some problems affecting science in 1900:
Blackbody radiation
Photoelectric effect
Line spectra
Spiraling of electrons into nucleus
8
Quantized Energy and Photons
Planck: energy can only be absorbed or released from
atoms in certain amounts called quanta.
This energy is proportional to the frequency ,, of the
radiation. The proportionality constant, h, is called
Planck’s constant and is equal to: 6.626  10-34 J.s
E = h
Matter can absorb or emit energy only in
quantum units which are multiples of h :
ΔE = nh
where n=integer
9
Quantized Energy and Photons
The Photoelectric Effect
•The photoelectric effect provides evidence for the
particle nature of light -- “quantization”.
•If light shines on the surface of a metal, there is a
point at which electrons are ejected from the metal.
•The electrons will only be ejected once the threshold
frequency, υ, is reached.
•Below the threshold frequency, no electrons are
ejected.
•This is consistent with the Planck statement that light
energy depends only on frequency.
10
Quantized Energy and Photons
The Photoelectric Effect
11
The Photoelectric Effect
Energy of
ejected
electrons
threshold
frequency
no
frequency, n
•Einstein assumed that light traveled in energy
packets called photons.
•The energy of one photon, E = hn.
12
1. Calculate the wavelength and energy of a photon with
frequency of 2.3 x 1014 Hz
2. Calculate the frequency and energy of a photon with
a wavelength of 420 nm
13
Lothar Meyer 1830-1895.
Codiscoverer of Periodic Table.
Meyer noticed a periodic trend in atomic sizes and
suggested there was an inner structure to the atom,
although it was too early to guess what the nature
of this structure was.
14
Bohr’s Model of the Hydrogen Atom
Bohr’s Model
•Rutherford assumed the electrons orbited the nucleus
analogous to planets around the sun.
•However, a charged particle moving in a circular path
should lose energy.
•This means that the atom should be unstable according
to Rutherford’s theory.
•Bohr noted the line spectra of certain elements and
assumed the electrons were confined to specific energy
states. These were called orbits.
15
Bohr’s Model of the Hydrogen Atom
Line Spectra
Colors from excited gases arise because electrons
move between energy states in the atom. These are
called line spectra.
Na
H
16
Line Spectra
prism
gas discharge
tube
17
Bohr’s Model of the Hydrogen Atom
Absorption
Emission
E4
E3
E2
E1
18
E3
E2
Absorption
E1
N
Emission
19
Bohr’s Model of the Hydrogen Atom
The first orbit in the Bohr model has energy E1 and is
assigned a quantum number, n=1. The next orbit has
energy E2 and quantum number, n=2.
Electrons in the Bohr model can only move between
orbits by absorbing and emitting energy in quanta
(hn).
The amount of energy absorbed or emitted on
movement between states is given by
 E = E f  E i = hn
BUT.......
The Bohr model doesn’t work!
20
The Wave Behavior of Matter
Knowing that light has a particle nature, it seems
reasonable to ask if matter has a wave nature.
Using Einstein’s and Planck’s equations, de Broglie
supposed:
h
l=
mv
The momentum, mv, is a particle property, where as l
is a wave property.
In one equation de Broglie summarized the concepts of
waves and particles, with noticeable effects if the objects
are small.
21
Quantum Mechanics and Atomic Orbitals
•Schrödinger proposed an equation that contains both
wave and particle terms.
(impress
2
2
2
2
h      

your
 2  2  2   V ( x, y, z ) = ih

2m  x
y
z 
t
friends)
•Solving the equation leads to the wave function Ψ These
are later referred to as orbitals.
•Wave functions (orbitals, Ψ ) are mathematical
quantities.
•The square of the wave function, Ψ 2 gives the probability
of finding the electron in some region of space.
•Orbitals have shapes and spatial orientations.
• The old Newtonian, classical theory is deterministic.
22
•The new quantum wave theory is probabilistic.
Quantum Mechanics and Atomic
Orbitals
Ψ2 = probability of finding electron
in box
23
Quantum Mechanics and Atomic
Orbitals
Orbitals and Quantum Numbers
If we solve the Schrödinger equation, we get wave
functions (orbitals), energies and quantum numbers.
Schrödinger’s equation requires 3 quantum numbers:
Principal Quantum Number, n. This is the same as
Bohr’s n. As n becomes larger, the atom becomes
larger and the electron is further from the nucleus.
24
Quantum Mechanics and Atomic
Orbitals
Orbitals and Quantum Numbers
Azimuthal Quantum Number, l. This quantum number
depends on the value of n.
l=0,1,2,....n-1
We usually use letters for l (s, p, d and f for l = 0, 1, 2, and 3).
l has to do with orbital shape.
l=
0
s
1
p
2
d
3
f
4
g
Magnetic Quantum Number, ml. This quantum number
depends on l. The magnetic quantum number has integral
values between -l and +l. Magnetic quantum numbers give the
3D orientation of each orbital.
25
n
1
2
3
4
l
0 (s)
0 (s)
1 (p)
0 (s)
1 (p)
2 (d)
0 (s)
1 (p)
2 (d)
3 (f)
Sub
Shell Shell
ml
K
0
1s
0
2s
L
1,0,-1
2p
0
3s
M
1,0,-1
3p
2,1,0,-1,-2
3d
0
4s
1,0,-1
4p
N
2,1,0,-1,-2
4d
3,2,1,0,-1,-2,-3 4f
26
Representation of Orbitals
The s Orbitals
All s-orbitals are spherical in shape.
As n increases, the s-orbitals get larger.
27
Representation of Orbitals
The p Orbitals
There are three p-orbitals px, py, and pz in each l=1 subshell.
These correspond to allowed values of ml
of -1, 0, and +1.)
The orbitals are dumbbell shaped.
As n increases, the p-orbitals get larger.
All p-orbitals have a node at the nucleus.
28
Representation of Orbitals
The p Orbitals
29
Representation of Orbitals
The d and f Orbitals
There are 5 d-orbitals in l=2 subshells.
There are 7 f-orbitals in l=3 subshells
Four of the d-orbitals have four lobes each.
One d-orbital has two lobes and a collar.
(We will not be concerned with the f-orbitals)
30
Representation of Orbitals
The d Orbitals
31
Quantum Mechanics and Atomic
Orbitals
Orbitals and Quantum Numbers
Orbitals can be ranked in terms of increasing energy to
yield an Aufbau diagram.
Note that the Aufbau diagram on next slide is for a
single electron system.
As n increases, note that the spacing between energy
levels becomes smaller.
32
Quantum Mechanics and Atomic Orbitals
Orbitals and Quantum
Numbers
Each square
is
an orbital.
There are n2 orbitals
in each shell.
Note that orbitals with same
value of n have the same energy
33
Orbitals in Many Electron Atoms (everything above
Energies of Orbitals
hydrogen)
Note: 2p above 2s
3p above 3s
3d above 3p
BUT: 4s below 3d
All due to screening
(next chapter).
Rule:for given value of n
energy of orbital increases
with increasing value of l
Example: for n=3
l= 0
1
2
s < p < d
Increasing energy
34
Orbitals in Many Electron Atoms
Mnemonic Device to remember Orbital Sequence
1s 2s 2p3s 3p4s 3d4p5s 4d5p6s
1s
2s
3s
4s
5s
6s
When filled with electrons, it’s
2p
3p 3d
4p 4d 4f
5p 5d 5f
6p
1s22s22p63s23p64s23d104p65s2...
5g
(superscripts are numbers of
electrons in orbital series)
35
Orbitals in Many Electron Atoms
Electron Spin and the Pauli Exclusion Principle
Studies in the 1920s demonstrated that the electron has
a property called spin.
This introduced a new (and fourth) quantum number,
the electron spin (or just spin) quantum number, ms
There are two values for ms: +1/2 and -1/2
36
Orbitals in Many Electron Atoms
Electron Spin and the Pauli Exclusion Principle
ms=+1/2
ms =-1/2
37
Orbitals in Many Electron Atoms
Electron Spin and the Pauli Exclusion Principle
Pauli’s Exclusion Principle: no two electrons can have
the same set of 4 quantum numbers.
Therefore, two electrons in the same orbital must have
opposite (+1/2 and -1/2) spins.
Can use arrows ( and ) to represent +1/2 and -1/2
spins, respectively.
For example, in the 1s orbital, n=1, l=0 and ml=0.
One electron can be placed (+1/2) and the other
at
(-1/2). But, that’s it! No more electrons in the
1s orbital. Why?
38
Electron Configurations
Periods 1, 2, and 3
Electron configurations tell us in which orbitals the
electrons for an element are located.
Three rules:
•electrons fill orbitals starting with lowest n and
moving upwards;
•no two electrons can fill one orbital with the same
spin (Pauli);
•for degenerate orbitals (orbitals with same
energy), electrons fill each orbital singly before any
orbital gets a second electron (Hund’s rule).
39
3d
(Hund’s Rule)
4s
3p
Hund’s Rule
3s
2p
Hund’s Rule
2s
1s
1s2 2s2 2p6 3s2 3p6 4s2 3d10
40
Electron Configurations
41
Electron Configurations and the Periodic
Table
42
Electron Configurations and the Periodic
Table
There is a shorthand way of writing electron
configurations
Write the core electrons corresponding to the preceding
filled Noble gas in square brackets.
Write the valence electrons explicitly.
Example, P: 1s22s22p63s23p3
but Ne is : 1s22s22p6
Therefore, P: [Ne]3s23p3.
43
Development of the Periodic Table
•How do we organize elements in a meaningful way that
will allow us to make predictions about undiscovered
elements?
•Arrange elements to reflect the trends in chemical and
physical properties.
•First attempt (Mendeleev and Meyer) arranged the
elements in order of increasing atomic weight.
•Modern periodic table: arrange elements in order of
increasing atomic number.
•Elements in same column (group) have same or similar
properties.
•Properties are periodic.
44
Effective Nuclear Charge and Screening
Screening occurs when you have more than one
electron to consider.
this electron
e-
e
e
+
e
is partially shielded
from this nucleus
by all the other
surrounding electrons
in atom
This electron experiences an effective nuclear charge, Zeff
45
Effective nuclear charge, Zeff, and screening
Outer (valence) electrons experience an effective
(not full) nuclear charge because of screening (or
shielding).
Effective nuclear charge
Zeff = Z – no. of core electrons
(Z = actual nuclear charge)
e.g., for Na (1s22s22p63s1)
Mg (1s22s22p63s2)
Al (1s22s22p63s23p1)
Zeff = 11-10 = +1
Zeff = 12-10 = +2
Zeff = 13-10 = +3
Note that Zeff increases across a row
46
The Periodic Table
The Periodic Table is used to organize the 114 elements
in a meaningful way
As a consequence of this organization, there are
periodic properties associated with the periodic table.
47
Development of the Periodic Table
48
Don’t forget to review historical questions
regarding the Periodic Table that are found
in the Problem Bank:
Dalton, Mendeleev, Rutherford, Lavoisier, Bunsen,
Berzelius, Seaborg, Curie, Davy, Bohr, Moseley,
Soddy, the Ancients, all discussed in Chapters 2 and 7.
Review groups (families) of elements in Chapter 2:
alkali metals, alkaline earths, rare earths (lanthanides),
actinides, transuraniums, halogens, noble (inert)
gases.
49
Sizes of Atoms
Covalent (atomic) radii
Array of Au atoms:
(1 Å = 10-10 m)
In Au, internuclear distance = 2.88 Å
Therefore, atomic radius = 2.88 = 1.44 Å
2
50
Sizes of Atoms
•As a consequence of the ordering in the periodic
table, properties of elements vary periodically.
•Atomic size varies consistently through the periodic
table.
•As we move down a group, the atoms become larger.
Each row (period)
is a new shell.
51
Sizes of Atoms
•As we move across a period, atoms become smaller.
There are two factors at work:
•principal quantum number, n, and
•the effective nuclear charge, Zeff.
•As we move across the periodic table, the number of
core electrons remains constant. But, the nuclear
charge increases.
•Therefore, there is an increased attraction between the
nucleus and the outermost electrons. This attraction
causes the atomic radius to decrease.
52
Sizes of Atoms
In general, atomic size increases this way
in the table.
inc
inc
inc
Atomic Size
53
Atomic
Radii
54
Problem: Estimate the As-I bond length
Radii:
As = 1.19 A
I = 1.33 A
1.33
I
1.19
Bond Length=
1.33 + 1.19=
2.52 A
As
55
Ionization Energy
• The first ionization energy, I1, is the amount of energy
required to remove an electron from a gaseous atom:
Na(g)  Na+(g) + e-.
•The second ionization energy, I2, is the energy required
to remove an electron from a gaseous ion:
Na+(g)  Na2+(g) + e-.
The larger the ionization energy, the more difficult it is
to remove the electron.
There is a sharp increase in ionization energy
when a core electron is removed.
56
Ionization Energy
element at end of row
element at beginning of row
57
Ionization Energy
Periodic Trends in Ionization Energy
•Ionization energy decreases down a group.
Group IA
Li (520 kJ/mol)
Na (496)
K (419)
Rb (403)
Cs (376)
This means that the outermost electron is more readily
removed as we go down a group. Why?
As the atom gets bigger, it becomes easier to remove an
electron from the most spatially extended orbital.
58
Ionization Energy
•Ionization energy generally increases across a
• period.
•As we move across a period, Zeff increases.
Therefore, it becomes more difficult to remove an
electron.
Period 2:
Li Be B C N O F Ne
IE: 520 899 801 1086 1402 1314 1681 2081
(all in kJ/mol)
59
Ionization
Energy
60
Ionization Energy
In general, ionization energy increases this way
in the table.
inc
inc
inc
Ionization Energy
61
Electron Affinities
•Electron affinity is the opposite of ionization energy.
•Electron affinity is the energy change when a gaseous
atom gains an electron to form a gaseous ion:
Cl(g) + e-  Cl-(g)
•Electron affinities are usually exothermic (as in the
above example)
inc
inc
inc
Electron Affinity
62
Electron Affinities
The added electron in Cl is placed in the 3p orbital to form
the stable 3p6 electron configuration.
63
Ionic Radii
Just as atom size is periodic, ion size is also periodic.
Cations (+ ions)
•To form cations, outermost or valence electrons are
removed.
•The effective nuclear charge has increased.
•Therefore, the cation is smaller than the parent.
Anions (- ions)
To form anions, electrons are added to the
outermost orbital.
•The nuclear charge has remained the same, but
the number of screening electrons is increased.
•Therefore, anions are larger than their parents.
64
Ionic Radii
For ions of the same charge, ion size increases down a
group.
All the members of an isoelectronic series have the same
number of electrons.
As nuclear charge increases in an isoelectronic series the
ions become smaller:
O2-
F-
Na+
Mg2+ Al3+
(all have 10 e’s)
65
Ionic Radii
Isoelectronic Species
O2Z=8
#e-=10
F-
Ne
Na+
Z=9
#e-=10
Z=10
#e-=10
Z=11
#e-=10
Mg2+
Z=12
#e-=10
Attraction to nucleus increases
Size Decreases
66
Ionic and Atomic Radii
67
Ionic Radii
Which of the following species is the largest?
B
B+ Al
Al+
Which of the following species is the smallest?
P
P-
S
S-
Which of the following species is the largest?
N- P+
P-
P
Which of the following species is the smallest?
Rb+ K+
Cl-
Ar
68
Metals, Nonmetals, and Metalloids
Metals
•Metallic character refers to the properties of metals
(shiny or lustrous, malleable and ductile)
• Metals react with acids (H+)
•Metallic character increases down a group.
inc
•Metallic character decreases across a period.
dec
inc
inc
•Metals have low ionization energies.
69
Metals
•When metals are oxidized they tend to form
characteristics cations.
•All group 1A metals form M+ ions.
•All group 2A metals form M2+ ions.
•Most transition metals have variable charges.
Metal oxides are basic:
CaO(s) + H2O
Ca(OH)2
Na2O(s) + H2O
2NaOH
70
Nonmetals
When nonmetals react with metals,
nonmetals tend to gain electrons
metals tend to lose electrons:
metal + nonmetal  salt
e.g.
2Al(s) + 3Br2(l)  2AlBr3(s)
Mg(s) + P  ?
K(s) + O2 (g)  ?
Ca(s) + Cl2  ?
Nonmetal oxides tend to be acidic:
CO2 + H2O  H2CO3
NO2 + H2O  HNO3
SO3 + H2O  H2SO4
71
Semimetals
Semimetal oxides tend to be amphoteric (they can
react either as weak acids or weak bases):
B2O3 + 3H2O  2B(OH)3
As an acid:
B2O3 + 6NaOH  2Na3BO3 + 3H2O
strong base
As a base:
B2O3 + 4HCl  2Cl2B(OH) + H2O
strong acid
72