# Chapter 8 - Electron Configurations and Periodicity

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```					Chapter 8 - Electron Configurations
and Periodicity
Remember n = principle quantum number (shell)
l = angular momentum quantum number (subshell)
ml = magnetic quantum number (orbital)
ms = spin quantum number

Each electron in an atom is governed by the Pauli Exlculsion
Principle.

• No two electrons in an atom can have the same set of four
quantum numbers.

1
Knowing that each orbital (ml) may accommodate a maximum
of 2 electrons (pair per orbital).

• The Pauli Exclusion Principle tells us that electron pairs
residing in the same orbital (ml or magnetic quantum
number) are required to have opposing spins;

ms = +½ , -½ or

The circle represents one orbital.

With this in mind we can now develop a table that summarizes
where electrons reside in the atom.

2
Principle    Angular      Magnetic Quantum Number Possible          Maximum
Quantum      Momentum Number               of       Electron Spin   Number of
Number or    Quantum      or Orbital       orbitals Quantum         electrons in
Shell        Number                                 Numbers for     subshell
or Subshell                            each orbital
n (1, 2, 3…) l (0 to n-1) ml (–l to +l)             ms (-½, +½)

1          0 (s)             0            1        -½, +½            2

2          0 (s)             0            1        -½, +½            2
1 (p)         -1, 0, +1        3        -½, +½            6

3          0 (s)              0           1        -½, +½             2
1 (p)          -1, 0, +1       3        -½, +½             6
2 (d)      -2, -1, 0, +1, +2   5        -½, +½            10

4          0 (s)              0           1        -½, +½             2
1 (p)          -1, 0, +1       3        -½, +½             6
2 (d)      -2, -1, 0, +1, +2   5        -½, +½            10
3 (f)    -3,-2,-1,0,+1,+2,+3   7        -½, +½            14

3
• That table summerizes the electron configuration of the atom.

• We shall see that this table will tell us how the periodic table
is formed.

An orbital diagram is used to show how the orbitals of a sub-
shell are occupied by electrons.

Let’s consider the orbital diagram for the Boron atom, atomic
number 5

5 electrons, 5 protons

Remember that there is a maximum of 2 electrons per orbital.
4
Starting with the principles you learned in Chapter 7 let’s
consider the orbitals in boron that can accommodate those 5
electrons.

To do this we must consider the relative energies of those
orbitals (circles)
l=0    l=1

3p

Energy
n = 3: 3s             The 5 electrons will fill
starting from the bottom.
2p
n = 2: 2s

n = 1: 1s
5
The filling pattern can be represented by the following:

1s       2s             2p

Where:
Each orbital is represented by a circle.

Each group of orbitals is labeled by its sub shell notation.

Electrons are represented by arrows:

– up for ms = +½ and down for ms = -½
6
The electron configuration for boron can be represented as:

Principle quantum number

2     2    1
Boron:1s 2s 2p                  Letter designation of the
subshell, (l)

Number of electrons in orbital

7
Aufbau Principle
• Every atom has an infinite number of
possible electron configurations.
The configuration associated with the lowest energy
level of the atom is called the “ground state.”

Other configurations correspond to “excited
states.”

Table 8.1 lists the ground state configurations of
atoms up to krypton. (A complete table appears in
Appendix D.)

8
9
The Aufbau principle is a scheme used to reproduce the
ground state electron configurations of atoms by following the
“building up” order.

Listed below is the order in which all the possible sub-shells fill
with electrons.

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

You need not memorize this order. As you will see, it can be
easily obtained.

10
Order for Filling Atomic Subshells

1s
2s     2p
3s     3p     3d
4s     4p     4d 4f
5s     5p     5d 5f
6s     6p     6d 6f

11
Example: Ground state electron configuration of Fe

Fe – 26 electrons

1s2 2s2 2p6 3s2 3p6 4s2 3d6 or 1s2 2s2 2p6 3s2 3p6 3d6 4s2

In class: Write down the ground state configurations for

Br – 35 electrons                Tc – 43 electrons

12
Abbreviations to electron configurations

• Use the atomic symbol of the closest noble gas to represent
the inner shell electrons.

Example: Fe – 26 electrons

Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Note that Ar (18 electrons) has the following configuration:

Ar: 1s2 2s2 2p6 3s2 3p6

The abbreviated form of the electron configuration for Fe is

Fe: [Ar] 4s2 3d6
13
Z=8 Oxygen   1s22s22p4   or [He]2s22p4

Z=12 Magnesium 1s22s22p23s2    or [Ne]3s2

Z=9 Fluorine 1s22s22p5   or [He]2s22p5

14
Configurations and the Periodic
Table
Note that elements within a given family (column) have
similar electron configurations.

For instance, consider the noble gases.

Helium     1s2
Neon       1s22s22p6
Argon      1s22s22p63s23p6
Krypton    1s22s22p63s23p63d104s24p6

15
The Group IIA elements are called the alkaline earth metals.

Beryllium 1s22s2
2 2   6 2
Magnesium 1s 2s 2p 3s
Calcium   1s22s22p63s23p64s2

Electrons that reside in the outermost shell of an atom - or in
other words, those electrons outside the “noble gas core (s2p6)”-
are called valence electrons.

16
Valence electrons or Valence shell
electrons:
are primarily involved in chemical reactions.

elements within a given group have the same “valence
shell configuration.”

accounts for the similarity of the chemical properties
among groups of elements. (Chapter 2)

17
The following slide illustrates how the periodic table provides a
sound way to remember the Aufbau sequence.

In many cases you need only the configuration of the outer
elements.

You can determine this from their position on the periodic
table.

The total number of valence electrons for an atom equals
its group number.

18
19
Consider column IV A       Carbon:    2s2 2p2
Silicon:   3s2 3p2
Four outermost electrons   Ge:        4s2 4p2
Hence, group IVA           Sn:        5s2 5p2
Pb:        6s2 6p2

Group IIA                  Be:        2s2
Mg:        3s2
Two valence electrons      Ca:        4s2
2
Sr:        5s
2
Ba:        6s

20
Now let’s consider orbital diagrams again. Recall that the
diagram for ground state configuration of boron is

1s        2s            2p

Now consider carbon (Z = 6) with the ground state
2 2   2
configuration 1s 2s 2p .

21
Three possible arrangements are given in the
following orbital diagrams.
1s    2s        2p
Diagram 1:
Diagram 2:
Diagram 3:

Each state has a different energy characteristics.

Only one of these configurations can be the true ground state
for carbon
22
Hund’s rule states that the lowest energy arrangement (the
“ground state”) of electrons in a sub-shell is obtained by
putting electrons into separate orbitals of the sub shell with
the same spin before pairing electrons.

Looking at carbon again, we see that the ground state
configuration corresponds to diagram 1 when following
Hund’s rule.

1s      2s             2p

23
To apply Hund’s rule to oxygen, whose ground state
configuration is 1s22s22p4, we place the first seven electrons as
follows.

1s    2s        2p

The last electron is paired with one of the 2p electrons to give a
doubly occupied orbital.

1s    2s       2p

Table 8.2 lists more orbital diagrams.

24
25
In Class: Draw the orbital diagrams for the ground state
configurations of vanadium (V) and silicon (Si)

26
Periodic Properties
• The periodic law states that when the elements are arranged
by atomic number, their physical and chemical properties
vary periodically.

• We will look at three periodic properties:

2. Ionization energy
3. Electron affinity

27
1. Periodic properties of atomic radii

When considering main group elements, within each period
(horizontal row), the atomic radius tends to decrease with
increasing atomic number (nuclear charge).

• Think of an increasing number of protons acting upon
electrons with approximately the same energy (same n
shell).

28
Within each group (vertical column), the atomic radius tends
to increase with the period number (or atomic weight).

Remember that size increases with principle quantum number,
n.

Think of electrons piling on another each other with each
successive shell, the onion effect.

29
Figure 8.17:
Representation of
group elements.

30
PERIODIC TABLE OF THE ELEMENTS
1A                                                                                                                                      8A
1                                                                                                                                       2
H                                                                                                                                       He
1.008   2A                                                                                      3A      4A      5A      6A      7A      4.003
3        4                                                                                       5       6       7       8       9      10
Li      Be                                                                                       B       C       N      O        F      Ne
Increasing    6.941   9.012                                                                                   10.81   12.01   14.01   16.00   19.00   20.18

11      12                                                                                      13      14      15      16      17      18
Atomic        Na      Mg                                                                                      Al      Si      P       S       Cl      Ar
3B      4B      5B      6B      7B              8B              1B      2B
22.99   24.31                                                                                   26.98   28.09   30.97   32.07   35.45   39.95

19      20      21      22      23      24      25      26      27      28      29      30      31      32      33      34      35      36
K       Ca      Sc      Ti      V       Cr      Mn      Fe      Co      Ni      Cu      Zn      Ga      Ge      As      Se      Br      Kr
39.10   40.08   44.96   47.88   50.94   52.00   54.94   55.85   58.93   58.69   63.55   65.39   69.72   72.61   74.92   78.96   79.90   83.80
37      38      39      40      41      42      43      44      45      46      47      48      49      50      51      52      53      54
Rb      Sr      Y       Zr      Nb      Mo      Tc      Ru      Rh      Pd      Ag      Cd      In      Sn      Sb      Te       I      Xe
85.47   87.62   88.91   91.22   92.91   95.94   (98)    101.1   102.9   106.4   107.9   112.4   114.8   118.7   121.8   127.6   126.9   131.3
55      56      57      72      73      74      75      76      77      78      79      80      81      82      83      84      85      86
Cs      Ba      La      Hf      Ta      W       Re      Os      Ir      Pt      Au      Hg      Tl      Pb      Bi      Po      At      Rn
132.9   137.3   138.9   178.5   181.0   183.8   186.2   190.2   192.2   195.1   197.0   200.6   204.4   207.2   209.0   (209)   (210)   (222)
87      88      89      104     105     106     107     108     109
Fr      Ra      Ac      Unq     Unp     Unh     Uns     Uno     Une
(223)   226.0   227.0   (261)   (262)   (263)   (262)   (265)   (266)

31
2. Ionization energy

The first ionization energy of an atom is the minimal
energy needed to remove the highest energy (outermost)
electron from the neutral atom.

For a lithium atom, the first ionization energy is illustrated
by:

2   1          +     2      −
Li (1s 2s ) → Li (1s ) + e
Ionization energy = 520 kJ/mol

32
Ionization energy trends

There is a general trend that ionization energies increase with
atomic number within a given period (row).

This follows the trend in size, as it is more difficult to
remove an electron that is closer to the nucleus.

For the same reason, we find that ionization energies,
again following the trend in size, decrease as we descend
a column of elements.

33
Figure 8.18: Ionization energy versus atomic number.

34
General Trend for Ionization Energy
PERIODIC TABLE OF THE ELEMENTS
1A                                                                                                                                      8A
1                                                                                                                                       2
H                                                                                                                                       He
1.008   2A                                                                                      3A      4A      5A      6A      7A      4.003
3       4                                                                                       5       6       7       8       9      10
Li     Be                                                                                       B       C       N      O        F      Ne
6.941   9.012                                                                                   10.81   12.01   14.01   16.00   19.00   20.18
11      12                                                                                      13      14      15      16      17      18
Na      Mg                                                                                      Al      Si      P       S       Cl      Ar
22.99   24.31   3B      4B      5B      6B      7B              8B              1B      2B      26.98   28.09   30.97   32.07   35.45   39.95

Increasing     19      20      21      22      23      24      25      26      27      28      29      30      31      32      33      34      35      36
K       Ca      Sc      Ti      V       Cr      Mn      Fe      Co      Ni      Cu      Zn      Ga      Ge      As      Se      Br      Kr
Ionization     39.10   40.08   44.96   47.88   50.94   52.00   54.94   55.85   58.93   58.69   63.55   65.39   69.72   72.61   74.92   78.96   79.90   83.80
37      38      39      40      41      42      43      44      45      46      47      48      49      50      51      52      53      54
Energy         Rb      Sr      Y       Zr      Nb      Mo      Tc      Ru      Rh      Pd      Ag      Cd      In      Sn      Sb      Te       I      Xe
85.47   87.62   88.91   91.22   92.91   95.94   (98)    101.1   102.9   106.4   107.9   112.4   114.8   118.7   121.8   127.6   126.9   131.3
55      56      57      72      73      74      75      76      77      78      79      80      81      82      83      84      85      86
Cs      Ba      La      Hf      Ta      W       Re      Os      Ir      Pt      Au      Hg      Tl      Pb      Bi      Po      At      Rn
132.9   137.3   138.9   178.5   181.0   183.8   186.2   190.2   192.2   195.1   197.0   200.6   204.4   207.2   209.0   (209)   (210)   (222)
87      88      89      104     105     106     107     108     109
Fr      Ra      Ac      Unq     Unp     Unh     Uns     Uno     Une
(223)   226.0   227.0   (261)   (262)   (263)   (262)   (265)   (266)

Increasing Ionization Energy (not perfect)

35
More on ionization trends.
Within rows (periods) we should be able to note that noble gases are the
most difficult to ionize. (see Figure 8.18 again).

- Filled s + p orbitals (s2p6) are very stable (meaning resistance to
chemical changes, reactions)

- Filled s (s2) orbitals are relatively stable when compared to s2p1, e.g.
Be vs. B, and Mg vs. Al (Group IIA vs. IIIA)

s             p               s          p
- Half filled p (p3) orbitals are relatively stable when compared to
elements with p4 valence shell configuration. (Group VA vs. VIA)

p                             p

36
nd
2 Ionization Energies
– The energies required at each step are known as the
first ionization energy, the second ionization energy,
and so forth.

1st Ionization Energy, I1: A(g) => A+(g) + e-

2nd Ionization Energy, I2: A+(g) => A2+(g) + e-

A general trend follows,

I 2 > I 1,   In >……..> I3 > I2 > I1

• Why?
37
3. Electron Affinity

• The electron affinity is the energy change for the process
of adding an electron to a neutral atom in the gaseous
state to form a negative ion.

For a chlorine atom, the first electron affinity is illustrated by:

– ([ Ne ]3s 2 3p 5 ) + e −         −           2    6
Cl                    +       → Cl ([ Ne ]3s 3p )
Electron Affinity = -349 kJ/mol

38
The more negative the electron affinity, the more stable the
negative ion that is formed.

Broadly speaking, the general trend goes from lower left
to upper right as electron affinities become more negative.

The noble gases (ns2 np6) and alkali earths (ns2) do not
have measurable electron affinities.

Table 8.4 gives the electron affinities of the main-group
elements.

39
Electron Affinity Trends
PERIODIC TABLE OF THE ELEMENTS
1A                                                                                                                                      8A
1                                                                                                                                       2
H                                                                                                                                       He
1.008   2A                                                                                      3A      4A      5A      6A      7A      4.003
3        4                                                                                       5       6       7       8       9      10

Generally    Li
6.941
Be
9.012
B
10.81
C
12.01
N
14.01
O
16.00
F
19.00
Ne
20.18
11      12                                                                                      13      14      15      16      17      18
Increasing   Na
22.99
Mg
24.31   3B      4B      5B      6B      7B              8B              1B      2B
Al
26.98
Si
28.09
P
30.97
S
32.07
Cl
35.45
Ar
39.95

Electron     19      20      21      22      23      24      25      26      27      28      29      30      31      32      33      34      35      36
K       Ca      Sc      Ti      V       Cr      Mn      Fe      Co      Ni      Cu      Zn      Ga      Ge      As      Se      Br      Kr
Affinity     39.10
37
40.08
38
44.96
39
47.88
40
50.94
41
52.00
42
54.94
43
55.85
44
58.93
45
58.69
46
63.55
47
65.39
48
69.72
49
72.61
50
74.92
51
78.96
52
79.90
53
83.80
54
Rb      Sr      Y       Zr      Nb      Mo      Tc      Ru      Rh      Pd      Ag      Cd      In      Sn      Sb      Te       I      Xe
85.47   87.62   88.91   91.22   92.91   95.94   (98)    101.1   102.9   106.4   107.9   112.4   114.8   118.7   121.8   127.6   126.9   131.3
55      56      57      72      73      74      75      76      77      78      79      80      81      82      83      84      85      86
Cs      Ba      La      Hf      Ta      W       Re      Os      Ir      Pt      Au      Hg      Tl      Pb      Bi      Po      At      Rn
132.9   137.3   138.9   178.5   181.0   183.8   186.2   190.2   192.2   195.1   197.0   200.6   204.4   207.2   209.0   (209)   (210)   (222)
87      88      89      104     105     106     107     108     109
Fr      Ra      Ac      Unq     Unp     Unh     Uns     Uno     Une
(223)   226.0   227.0   (261)   (262)   (263)   (262)   (265)   (266)

Generally Increasing Electron Affinity

40
41
Periodic Chemical Properties
The physical and chemical properties of the main-group

Group IA – valence shell of ns1, know as the alkali metals Li,
Na, K, Rb, Cs

All react with water forming H2(g).

2Na(s) + 2H2O(l) => 2NaOH(aq) + H2(g)

Li, K, Rb, & Cs will react in the same way and same
stoichiometry

42
Let R = alkali metal, all will form ROH when reacted
with H2O.

2R(s) + 2H2O(l) => 2ROH(aq) + H2(g)

All alkali metals with form RO when reacted with O2(g)

R(s) + ½O2(g) => RO(s)

All Group IA elements will react with Cl2(g) to RCl:

2Na(s) + Cl2(g) => 2NaCl(s)

Note the ionic state of the alkali metal in the products of those
+
reactions. All reactions produced R .
43
Now let’s consider the electron configurations of the valence
electrons of the reactant, Na and the product, Na+, in the
following half reaction.

Na, 3s1 => Na+, 3s0 + e-

♦ Recall, that s2p6 valance configuration is very stable. This is
essentially achieved by Na the loss of one e- to the 2s22p6,
pseudo-noble gas core or 3s0.

♦ All alkali metals seek this trend to assume the pseudo-noble
1
gas configuration by the loss of the one electron of its ns
valence shell.

44
Quick review of redox reactions:
Find the oxidation state for each element in the following
reaction:

2K(s) + Cl2(g) => 2KCl(s)
oxidation numbers:       ___    ___         __ __

half reactions:       2K        => 2K+ + 2e-
Cl2 + 2e- => 2Cl-

Reducing agent: _________

Oxidizing agent:__________
45
Summary of Group IA,
1
Alkali Metals, ns

♦   React violently with water to form H2
♦   Readily oxidized to 1+, ns 0

♦   Metallic character, oxidized in air
♦   R2O in most cases

46
2
Group IIA ns – alkali earths – Be,
Mg, Ca, Sr, Ba
All follow the same trend in their reaction with O2(g),

2Ca(s) + O2(g) => 2CaO(s)

2R(s) + O2(g) => 2RO(s)

This is driven by:
2        2+   0
R (ns ) => R (ns ) + 2e-

Ca (4s2) => Ca2+ (4s0) + 2e-

47
Group IIA – Trends

♦ React with water to form H2
2       6
♦ Closed s shell configuration, (n-1)s (n-1)p
♦ Metallic

48
2     1-10
Transition Metals – ns (n-1)d

• May have several oxidation states
• Metallic
• Reactive with acids

49
2    1
Group III A, ns np

Metals (except for boron)
Several oxidation states (commonly 3+)
2    1      0   0
ns np => ns np + 3e-

Oxide formation trend: B2O3, Al2O3….R2O3

50
2   2
Group IV A ns np
Oxidation numbers vary between 4+ and 4-
2   2            2   6
ns np + 4e- => ns np

ns2 np2 => ns0 np0 + 4e-

C (nonmetal)
Si, Ge (semimetals, or metalloids)
Sn, Pb (metals)

51
2      3
Group V A, ns np
Form anions generally(1-, 2-, 3-), though positive
oxidation states are possible

Oxide formation varies, hard to predict

Nonmetals – N, P
Semimetal – As
Metals – Sb, Bi

52
2     4
Group VI A (Chalcogens), ns np
Form 2- anions generally, though positive
oxidation states are possible

ns2 np4 + 2e- => ns2 np6

React vigorously with alkali and alkali earth metals
Example: 2Na(s) + S(s) => Na2S(s)

Nonmetals

53
2   5
Group VIIA (Halogens), ns np
Form mono-anions
2   5         2  6
ns np + e- => ns np

High electronegativity (electron affinity)

Diatomic gases:
F2, Cl2, Br2, I2, At2

Most reactive of nonmetals.

54
2     6
Noble Gases, ns np

Minimal reactivity filled s+p subshells, ns2 np6

Monatomic gases He, Ne, Ar, Kr, Xe, Rn
2     6
Closed shell, ns np

55
Operational Skills & Highlights
• Applying the Pauli exclusion principle.

• Determining the configuration of an atom using the Aufbau
principle.

• Determining the configuration of an atom using the period
and group numbers.

•   Applying Hund’s rule.

•   Applying periodic trends.

56

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