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Transcript of Electronegativity
Bonding - AS Chemistry
By Samiul Oliver Khan
What is Electronegativity?
Electronegativity is defined as
"The relative tendency of an atom to pull electron density towards itself from a covalent bond"
The chemical property was first proposed by the American scientist Linus Pauling in 1932. He devised calculations to give relative values of electronegativity for the elements in the periodic table.
These relative values should range from 0.7 to 4.0.
4.0 is the highest electronegative value an element can have. Fluorine has an electronegativity value of 4.0. As a result Fluorine has the greatest power to pull electron density towards itself from a covalent bond.
Electron density is the area where there is the greatest probability of finding electrons. This area is negatively charged.
Factors affect Electronegativity
Electronegativity depends on:
the nuclear charge
the distance between the nucleus and outer shell electrons
shielding of the nuclear charge by electrons in inner shells
Trends in the Periodic Table
increases across a period
. As we move across the period, there is the same number of energy levels in atoms of the element. Therefore the
effect of shielding is negligible
. However, the
nuclear charge increases
across the group and thus the valence electrons are pulled closer to the nucleus. As a result the
distance between the nucleus and the outer shell electrons is reduced
. The ability to pull electron density towards itself is much stronger.
decreases down the group
. As we move down the group, despite the nuclear charge increasing, there are
more energy levels in the atom
. As a result,
the distance between the nucleus and the outer electron shells increases
. Due to there being more inner shells as we move down the group there is
more shielding of the nuclear attraction
from the nucleus with the valence electrons. The ability to pull electron density towards itself has decreased since the attraction is much weaker.
An introduction to Polarity
Polarity is caused by the unequal sharing of the electrons between atoms that are bonded together covalently.
Compounds can either be ionic or covalent. However there are some compounds such as Beryllium Chloride which have ionic character. We call compounds like Beryllium Chloride "polar".
We can determine whether a compound is ionic, covalent or polar by considering the difference of electronegativity between the atoms in a compound. It is important to realise that the electron density in a compound will be distributed differently depending on the type of compound.
You will be given a table showing the electronegativity values of atoms in the exam.
The difference in electronegativity between atoms in an ionic compound is larger than 1.7 (or very large).
Electron density is completely transferred from one atom to another.
For example: Sodium Chloride
Electronegativity of Na-0.9
Electronegativity of Cl – 3.0
Difference = 2.1 pull effect by the Cl atom
Ionic bonds are generally formed between reactive metals and non metals where the difference in electronegativity is large.
The difference in electronegativity between atoms in a covalent compound is virtually 0 (less than 0.3/ very small).
E.g - Chlorine Molecule
Electronegativity of Chlorine: 3.0
Difference ( Cl-Cl ) = 3.0 – 3.0 = 0
Note that the bonds between identical non metal atoms will always be non polar as the electronegativity difference will be 0
The Difference in electronegativity between atoms in a polar compound is between 0.3 to 1.7
Electronegativity of H: 2.1
Electronegativity of Cl: 3.0
Electronegativity difference: 3.0 – 2.1 = 0.9
The Chlorine atom will pull the electron density towards itself.
"Polar covalent bonds are covalent bonds with greater electron density around one of the two atoms. As a result the atom pulling the electron will gain a slightly negative charge and the atom that the electron pair is being pulled away from will gain a slightly positive charge"
ALWAYS USE THE TERM: SLIGHTLY
δ means that there is a slight charge.
Since the Fluorine atom is more electronegative it pulls the electron density towards itself. As a result greater negative charge surrounds the atom. Overall a slightly positive charge surrounds the hydrogen atom. This forms a slightly positive hydrogen end and a slightly negative fluorine end.
Special Cases: Considering Shape
A symmetrical molecule will not be polar since the charges cancel out. There is no net dipole moment.
However a non-symmetrical molecule will be polar as the individual bonds are polar:
AMMONIA – Polar N-H bonds (notice the lone pair)
ClCH3 - Polar C-Cl bond – different bonds involved
Water – Polar O-H bonds – lone pairs involved
I have used a couple of examples written by Neil Goalby in his powerpoint notes.
THANKS FOR WATCHING