> See also: > - Reference # Electronegativity **Electronegativity** is a measure of the tendency of an atom to *attract electrons* (or areas of electron density) towards itself. The general trend for electronegativity is an increase towards the upper right corner of the periodic table (*in the direction of fluorine*). > [!important]+ Electronegativity Chart > ![[Electronegativity Chart.png]] ### Determining Bond Type We can predict a compound’s **bonding mode** by looking at the difference in electronegativity. | Bond Type | Electronegativity Difference | | --- | --- | | Non-Polar (Pure) Covalent | $ x < 0.4 $ | | Polar Covalent | $ 0.4 \le x \le 1.8 $ | | Ionic | $ x > 1.8 $ | Polar refers to the fact that one atom is pulling more than the other. This difference in positive and negative regions of a molecule forms [[Dipoles]]. ![[Electronegativity Bond Types.png]] Using the different levels of electronegativity, we can say that: 1) Covalent bonds are typically between two nonmetals 2) Ionic bonds are typically between a metal and a nonmetal Ionic bonds are stronger than covalent bonds - The electrons are essentially "popping off" the original atom and onto the new one due to just how different their electronegativities are --- > As the electronegativity difference increases between two atoms, the bond becomes more ionic. --- ### Partial Charges $𝛿^−$ $𝛿^+$ ## Factors Influencing Electronegativity > [“Why Do Acids Burn?” - Buy Why? (Youtube)](https://www.youtube.com/watch?v=Y3oY3vbuDR8) It can be thought of as the ratio between the protons in the nucleus and the electron shells underneath the valence shell.