polar bonds

ΔE

Metalic bonds	many metals have a low value of E
Ionic bonds	ΔE: > ±1,6
Covalent bonds - covalent, non polar - covalent, polar	0 < ΔE < ±1,6 0 < ΔE: < ±0,4 ΔE: > ±0,4

Att.
In natural sciences we use lots of symbols from the Greek alphabeth, like:
The Greek capital Delta Δ: represents the exact difference between two values.
the Greek small lettre delta δ: represents that there is only a slight difference.

In a covalent bond can exist a difference between de E values (electronegativity) of the two participating atoms.
Then there isa ΔE, that means: a calculatable difference between the electronegativities of those two atoms.
There will be some dislocation of the electrons between the atoms, i.e. of the shared electrons.
The side to where those electons shift will become a bit more negative, because electrons are negative.
The other side of the bond, with the atom that has the bigger attraction (the biggest ΔE), will have a bit of a negative charge that we indicate as: (δ^-).
At the other side of the is an atom with the smaller electronegativity. That atom must let go (a bit) the electrons, and that atom will become a bit postive = (δ⁺).

Covalent bonds can have a complete NON POLAR character if there is no δ^- and no δ⁺.

If inside the covalent bond there is a certain difference, if the shared electrons stay not exactly in the middle, if there is a certain disequilibrium of the charges δ⁺ and δ^-, then we call that bond: POLAR.

Let's look at it in another way:
The bigger the difference of electronegativity E of the elements of a bond, the more ionic will be the character of that bond between those elements.

Thus a substance like Al₂O₃ has a 70% ionic character and a 30% covalent character.

Dipoles / dipole molecules

The following image you might have seen before:

Molecules can contain one or more polar bonds, and they might be, or not, a dipole.
That depends on the symmetria of the molecule (consider the examples).

Examples:

CS₂ (ΔE = ±0)	CO₂ (ΔE = ±1.0)	H₂O (ΔE = ±1.3)
Covalent bonds	Covalent bonds	Covalent bonds
Non polar molecules	Non polar molecules	Polar molecules
There is no dipole	There is no dipole	There is a dipole
There are no δ⁺ and δ^-	There are δ⁺ and δ^- of which the central points do overlap	The central points of δ⁺ and δ^- do not overlap each other (remain at a distance)
S=C=S	O = C = O δ^- δ⁺ δ^-	δ⁺ δ⁺ H H \ / O δ^-