ELECTRONEGATIVITY
| Period | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |||||||||||||||||
| 1 | H 2.20 |
He 0 |
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| 2 | Li 0.98 |
Be 1.57 |
B 2.04 |
C 2.55 |
N 3.04 |
O 3.44 |
F 3.98 |
Ne 0 |
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| 3 | Na 0.93 |
Mg 1.31 |
Al 1.61 |
Si 1.90 |
P 2.19 |
S 2.58 |
Cl 3.16 |
Ar 0 |
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| 4 | K 0.82 |
Ca 1.00 |
Sc 1.36 |
Ti 1.54 |
V 1.63 |
Cr 1.66 |
Mn 1.55 |
Fe 1.83 |
Co 1.88 |
Ni 1.91 |
Cu 2.00 |
Zn 1.65 |
Ga 1.81 |
Ge 2.01 |
As 2.18 |
Se 2.55 |
Br 2.96 |
Kr 0 |
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| 5 | Rb 0.82 |
Sr 0.95 |
Y 1.22 |
Zr 1.33 |
Nb 1.60 |
Mo 2.16 |
Te 1.90 |
Ru 2.20 |
Rh 2.28 |
Pd 2.20 |
Ag 1.93 |
Cd 1.69 |
In 1.78 |
Sn 1.96 |
Sb 2.05 |
Te 2.10 |
I 2.66 |
Xe 2.6 |
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| 6 | Cs 0.79 |
Ba 0.89 |
La# 1.10 |
Hf 1.30 |
Ta 1.50 |
W 2.36 |
Re 1.90 |
Os 2.20 |
Ir 2.20 |
Pt 2.28 |
Au 2.54 |
Hg 2.00 |
Tl 2.04 |
Pb 2.33 |
Bi 2.02 |
Po 2.00 |
At 2.20 |
Rn 0 |
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| 7 | Fr 0.7 |
Ra 0.89 |
Ac* 1.1 |
Rf |
Db |
Sg |
Bh |
Hs |
Mt |
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| #Lanthanides | Ce 1.12 |
Pr 1.13 |
Nd 1.14 |
Pm 1.13 |
Sm 1.17 |
Eu 1.2 |
Gd 1.2 |
Tb 1.1 |
Dy 1.22 |
Ho 1.23 |
Er 1.24 |
Tm 1.25 |
Yb 1.1 |
Lu 1.27 |
|
| *Actinides | Th 1.3 |
Pa 1.5 |
U 1.38 |
Np 1.36 |
Pu 1.28 |
Am 1.3 |
Cm 1.3 |
Bk 1.3 |
Cf 1.3 |
Es 1.3 |
Fm 1.3 |
Md 1.3 |
No 1.3 |
Lr |
Electronegativity increases as you go across a period because, the next element along has an extra proton and an extra electron. The extra proton is in the nucleus, which means that the nucleus has a greater charge on it. The extra electron goes into the orbital 
with the lowest available energy, but because we are going across a period this means that the principle number of the orbital stays the same as for the last element. This means that the valence electrons do not experience any extra shielding from the nucleus, so they feel a greater effective nuclear charge and are more strongly attracted by the nucleus. This stronger attraction means the electrons are pulled closer to the nucleus, meaning the atom gets smaller (atomic radius decreases across a period). Thus, a shared pair of electrons will be more strongly attracted to the nucleus and so the atom withdraws more electron density. It is more electronegative.
Electronegativity decreases as you go down a group. This is because as you go down a group, the principal number of the valence orbital increases, meaning that there is an extra 'shell' of electrons between the valence electrons and the nucleus. This means that the valence electrons experience greater shielding from the nucleus. This factor is more important than the increased number of protons in the nucleus and the increased charge on the nucleus. So despite the extra protons the valence electrons are less strongly attracted by the nucleus, and the electrons are not held as close to the nucleus (atom radius increases down a group). Thus, a shared pair of electrons will be less strongly attracted to the nucleus, so the atom withdraws less electron density. It is less electronegative.
You may have noticed that the d block elements/transition metals do not seem to follow these trends quite as well as the s and p block elements, and there are indeed many exceptions to the trends in the d block. This is due to the valence electrons being in the d orbitals, these do not shield other electrons in the same way that s and p orbitals do. They do not shield electrons anywhere near as much as the s and p orbitals.
Move on to the next page: Detailed explanation of trends
Author: Alex Warren (document modification date - 10 May 2003)