Course Outline


The Periodic Table



 

After completing this unit you should be able to:

The Main Group Elements

When the elements are arranged in order of increasing atomic number it is observed that certain properties seem to repeat in a regular way.  We start a new row (a period) so that elements with similar properties fall in the same column (a group).  This arrangement is called the periodic table.  The second and third periods have eight elements each (you should have already learned their names and symbols), and on the basis of their properties they define eight main groups, designated 1A to 8A.  (An alternative numbering in common use is shown in parentheses).  The fourth and subsequent periods also have main group, or representative, elements:

 

1A
(1)

2A
(2)

B groups
(3 - 12)

3A
(13)

4A
(14)

5A
(15)

6A
(16)

7A
(17)

8A
(18)

1

1
H
2
He

2

3
Li
4
Be
  5
B
6
C
7
N
8
O
9
F
10
Ne

3

11
Na
12
Mg
  13
Al
14
Si
15
P
16
S
17
Cl
18
Ar

4

19
K
20
Ca
21 - 30 31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr

5

37
Rb
38
Sr
39 - 48 49
In
50
Sn
51
Sb
52
Te
53
I
54
Xe

6

55
Cs
56
Ba
57 - 71 72 - 80 81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Rn

7

87
Fr
88
Ra
89 - 103 104 - 112            

Some of the groups have common names.  Group 1A elements are the alkali metals and group 2A are the alkaline earths.  Group 7A are the halogens and group 8A are the noble gases, so called because they are extremely unreactive.  You should learn the names and symbols of the elements rubidium (Rb) and cesium (Cs) from group 1A, strontium (Sr) and barium (Ba) from group 2A, tin (Sn) and lead (Pb) from group 4A, arsenic (As) from 5A, selenium (Se) from 6A, and bromine (Br) and iodine (I) from 7A.

The Transition Elements

The fourth period has ten more elements squeezed in between the main groups 2A and 3A.  These define the 'B' groups.  Since they are all metals, these elements are referred to as the transition metals.

 

3B
(3)

4B
(4)

5B
(5)

6B
(6)

7B
(7)

8B
 (8)        (9)        (10)

1B
(11)

2B
(12)

4

21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn

5

39
Y
40
Zr
41
Nb
42
Mo
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
Cd

6

57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
80
Hg

7

89
Ac
                 

You should learn the names and symbols of the elements 21 - 30: scandium (Sc), titanium (Ti), vanadium (V), chromium (Cr), manganese (Mn), iron (Fe), cobalt (Co), nickel (Ni), copper (Cu), and zinc (Zn).  You should also know the symbols for a few other common transition metals: silver (Ag), gold (Au), tungsten (W) and mercury (Hg).

Squeezed into the sixth period after lanthanum (Z = 57) are another 14 elements, with an equivalent set in the seventh period.  On most periodic tables these elements are drawn in two separate rows at the bottom so that the chart does not become unmanageably wide.  These are the inner transition elements, also called the rare earths, of which the best known is uranium (U, Z = 92).  All of the known elements beyond uranium have been artificially created and are radioactive.

For more information about any particular element, try one of these online periodic tables:
http://periodic.lanl.gov/
http://www.periodicvideos.com/

Electron Distribution

Electrons are distributed around the nucleus in shells, with the outermost shell having the same number (n) as the period that the element is in.  For example silicon, in the third period, has its outermost electrons in the shell with n = 3.  As n  increases, the volume of the shell increases, and electrons are on average further from the nucleus.  Each shell is made up of n sub-shells, so the first shell has just one sub-shell, the second shell has two sub-shells, and so on.  For historical reasons, the sub-shells are named s, p, d, f, etc, and each letter represents a different shape for the region in space occupied by an electron in that sub-shell.  Each kind of sub-shell can hold a different number of electrons as shown in the following table:

Shell
n
Electrons
allowed
s p d f g Electrons in noble gas
at end of period n
1 2 2         He 2
2 8 2 6       Ne 2.8
3 18 2 6 10     Ar 2.8.8
4 32 2 6 10 14   Kr 2.8.18.8
5 (50) 2 6 10 14 (18) Xe 2.8.18.18.8
6 (50+) 2 6 10 (14) (18) Rn  2.8.18.32.18.8
7 (50+) 2 (6) (10) (14) (18)    

Numbers in parentheses are never reached because only 114 elements have been observed so far.

Each sub-shell in an atom has an associated potential energy.  Electrons in atoms occupy sub-shells with the lowest potential energy available.  A hydrogen atom has one electron in its 1s subshell, and a lithium atom has two electrons in its 1s subshell and 1 electron in its 2s subshell.  Within a shell, for atoms with more than one electron, the energy of sub-shells increases in the order s < p < d < f.  An isolated oxygen atom has two electrons in 1s, two electrons in 2s, and four in 2p, written as 1s22s22p4.

For larger atoms with many electrons the order of energy of different sub-shells follows the pattern

ns < (n-2)f < (n-1)d < np                (for example:  4s < 3d < 4p)

If we think of building up an atom of an element from the one before it in the periodic table, we arrive at a noble gas after we have finished putting six electrons into a p sub-shell.  The next two elements have electrons in the next s sub-shell, then skipped (n-2)f  and (n-1)d sub-shells (if there are any) are filled before filling the outermost p sub-shell.  This pattern of filling of sub-shells can be traced through the periodic table and leads us to refer to groups 1A and 2A as the s-block, groups 3A - 8A as the p-block, the B groups as the d-block, and the inner transition elements as the f-block.

The electron distribution in atoms of the noble gases (each of which has eight electrons in its outermost shell) is particularly stable, which is why they exhibit so little chemical reactivity.  The chemical properties of an element are determined by the number of electrons in its outermost shell, also called the valence shell.  Within each group, atoms of each element have the same number of valence electrons, which is why they have similar chemical properties.  Main group elements have the same number of electrons in their valence shell as their group number (or the units digit of their group number in the 1 - 18 system), so we can see the number of valence electrons in an atom by looking at the periodic table.

 

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