Electronegativity & Bond Polarity in a Snap!
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The key points covered in this video include:
1. What is Electronegativity?
2. How do we measure Electronegativity?
3. Trends in Electronegativity - Why do we see these trends?
4. Polar Bonds
5. Dipole Moments
6. The Spectrum of Bonds
7. Polar Molecules
Electronegativity
Covalent bonds involve a shared pair of electrons. A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms. Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in the covalent bond.
How do we measure Electronegativity?
Electronegativity is measured on the Pauling Scale. This scale was invented by Chemist Linus Pauling.
Trends in Electronegativity
Electronegativity increases: Across a period, Up a group. In all directions towards Fluorine.
Why do we see these trends in Electronegativity?
Electronegativity increases across a period. The charge on the nucleus increases across a period. The number of protons in the nucleus increases, Increased attraction for the outer electrons. The bonding pair of electrons are attracted more strongly. Electronegativity increases up a group. Down the group, the bonding pair of electrons is held increasingly further away for the nucleus. The number of shells increases, Distance of the outer electrons from nucleus increases. The bonding pair of electrons are attracted less strongly.
Polar Bonds
The Chlorine is more electronegative than the Hydrogen. Cl has a greater attraction for the electrons than H. Electrons are closer to the Cl than the H. Chlorine: 3.0, Hydrogen: 2.1.
Dipoles
The differing attraction for the pair of electrons allows there to be a small charge difference between the atoms. This is a permanent dipole. The charge difference is always present.
Polar and Non-Polar Bonds
Non-Polar Bonds. If the two bonding atoms are identical, their attraction for the shared pair of electrons is equal. The electrons are equally distributed between the bonding atoms. The bond is perfectly covalent. Polar Bonds. If the two bonding atoms are different, their attraction for the shared pair of electrons is unequal. The bonding atom with a greater attraction for the shared pair of electrons is more electronegative. The bond is polarised.
Electron Density
The electron density relates to the probability of finding electrons at a particular position in space. It can be imagined as a cloud of electrons around the nucleus.
The Spectrum of Bonds
Rather bonds existing as discretely ionic and covalent they exist on a spectrum. Ionic Bonding - The difference in electronegativity is so great that one atom effectively takes the electron from the other. Polar-Covalent Bonding - The difference in electronegativity is small. The atoms share the electrons unequally. The bond is polarised. Covalent Bonding - There is no difference in electronegativity. The molecule is electronically symmetrical.
Polar and Non-Polar Molecules
Molecules containing polar bonds are not always polar. The symmetry of polar bonds can cancel the effect of any permanent dipole. Non-Symmetrical - A difference in charge exists across the molecule. There is an overall dipole. The molecule is polar. Symmetrical - The symmetry of the molecule means that the effect of any permanent dipoles is cancelled out. Linear, Trigonal Planar or Tetrahedral shape, All atoms attached to the central atom are identical. No Difference in charge exists across the molecule. The molecule is non-polar.
Summary: Polar and Non-Polar Bonds
Non-Polar Bond - The two bonding atoms are identical and the electrons are equally distributed between the bonding atoms. Example Hydrogen: 2.1. Polar Bond - The two bonding atoms are different and their attraction for the shared pair of electrons is unequal. Example: Hydrogen: 2.1, Chlorine: 3.0
Summary: Polar and Non-Polar Molecules
Polar Molecule - A difference in charge exists across the molecule. Example: Hydrogen: 2.1, Oxygen: 3.5. Non-Polar Molecule - The symmetry of the molecule means that the effect of any permanent dipoles is cancelled out and there is no difference in charge exists across the molecule. Oxygen: 3.5, Carbon: 2.5.
