uncle tom’s cabin play Electronegativity


Periodic Trends — Electronegativity

1A2A3A4A5A6A7A8A
(1)(2)(13)(14)(15)(16)(17)(18)
3B4B5B6B7B8B1B2B
(3)(4)(5)(6)(7)(8)(9)(10)(11)(12)
1

H

2.20

He

n.a.

2

Li

0.98

Be

1.57

B

2.04

C

2.55

N

3.04

O

3.44

F

3.98

Ne

n.a.

3

Na

0.93

Mg

1.31

Al

1.61

Si

1.90

P

2.19

S

2.58

Cl

3.16

Ar

n.a.

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

1.90

Zn

1.65

Ga

1.81

Ge

2.01

As

2.18

Se

2.55

Br

2.96

Kr

3.00

5

Rb

0.82

Sr

0.95

Y

1.22

Zr

1.33

Nb

1.60

Mo

2.16

Tc

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

2.60

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

1.62

Pb

2.33

Bi

2.02

Po

2.00

At

2.20

Rn

n.a.

7

Fr

0.70

Ra

0.89

Ac

1.10

 

Rf

n.a.

Db

n.a.

Sg

n.a.

Bh

n.a.

Hs

n.a.

Mt

n.a.


Ds

n.a.


Rg

n.a.

Uub

n.a.

Uuq

n.a.

6

 

Ce

1.12

Pr

1.13

Nd

1.14

Pm

1.13

Sm

1.17

Eu

1.20

Gd

1.20

Tb

1.10

Dy

1.22

Ho

1.23

Er

1.24

Tm

1.25

Yb

1.10

Lu

1.27

7

 

Th

1.30

Pa

1.50

U

1.38

Np

1.36

Pu

1.28

Am

1.30

Cm

1.30

Bk

1.30

Cf

1.30

Es

1.30

Fm

1.30

Md

1.30

No

1.30

Lr

1.30

Electronegativities reported in Pauling units

Data taken from John Emsley, The Elements, 3rd edition. 
Oxford:  Clarendon Press, 1998.

 

Electronegativity refers to the ability of an atom to attract shared
electrons in a covalent bond.  The higher the value of the
electronegativity, the more strongly that element attracts the
shared electrons.

The concept of electronegativity was introduced by Linus Pauling
in 1932; on the Pauling scale, fluorine is assigned an
electronegativity of 3.98, and the other elements are scaled
relative to that value.  Other electronegativity scales include the Mulliken scale, proposed by Robert S. Mulliken in 1934, in which the
first ionization energy and electron affinity are averaged together,
and the Allred-Rochow scale, which measures the electrostatic
attraction between the nucleus of an atom and its valence electrons.

Electronegativity varies in a predictable way across the periodic
table.  Electronegativity increases from bottom to top in
groups
, and increases from left to right across periods.  Thus,
fluorine is the most electronegative element, while francium is one
of the least electronegative. (Helium, neon, and argon are not
listed in the Pauling electronegativity scale, although in the
Allred-Rochow scale, helium has the highest electronegativity.)  The trends are not very smooth
among the transition metals and the inner transition metals, but are
fairly regular for the main group elements, and can be seen in the
charts below.

 

 

 

The difference in electronegativity between two bonded elements
determines what type of bond they will form.  When atoms with an electronegativity difference of greater
than two units are joined together, the bond that is formed is
an ionic bond, in which the more electronegative element
has a negative charge, and the less electronegative element has
a positive charge.  (As an analogy, you can think of it as
a game of tug-of-war in which one team is strong enough to pull
the rope away from the other team.)  For example, sodium
has an electronegativity of 0.93 and chlorine has an
electronegativity of 3.16, so when sodium and chlorine form an ionic
bond, in which the chlorine takes an electron away from sodium,
forming the sodium cation, Na+, and the chloride anion,
Cl.  Particular sodium and chloride ions are not
"tied" together, but they attract each other very strong because of
the opposite charges, and form a strong crystal lattice.

When atoms with an electronegativity difference of less than two
units are joined together, the bond that is formed is a covalent
bond
, in which the electrons are shared by both atoms.  When two of the same atom
share electrons in a covalent bond,
there is no electronegativity difference between them, and the
electrons in the covalent bond are shared equally — that
is, there is a symmetrical distribution of electrons
between the bonded atoms.  These bonds are nonpolar
covalent bonds
.  (As an analogy, you can think of it as
a game of tug-of-war between two equally strong teams, in which
the rope doesn’t move.)  For example, when two chlorine
atoms are joined by a covalent bond, the electrons spend just as
much time close to one chlorine atoms as they do to the other, and
the resulting molecule is nonpolar:

When the electronegativity difference is between 0 and 2,
the more electronegative element attracts the shared more
strongly, but not strongly enough to remove the electrons completely to
form an ionic compound.  The electrons are shared
unequally
— that is, there is an unsymmetrical
distribution of electrons between the bonded atoms.  These
bonds are called polar covalent bonds.  The more
electronegative atom has a partial negative charge,
d
, because the
electrons spend more time closer to that atom, while the less
electronegative atom has a partial positive charge,
d+
, because the
electrons are partly (but not completely) pulled away from that
atom.  For example, in the hydrogen chloride molecule,
chlorine is more electronegative than hydrogen by 0.96
electronegativity units.  The shared electrons spend more time
close to the chlorine atom, making the chlorine end of the molecule
very slightly negative (indicated in the figure below by the blue
shaded region), while the hydrogen end of the molecule is very
slightly positive (indicated by the red shaded region), and the
resulting molecule is polar:

For molecules with more than one covalent bond, the
three-dimensional shape of the molecule and how the polar bonds are
oriented with respect to each other, determines whether or not the
molecule is polar.  This polarity of molecules plays a large
role in determining the physical properties of compounds.