Periodic Trends - Chemistry

Effective Nuclear Charge (Z/eff)

The net positive charge a valence e- can actually feel, explaining how strong the nucleus can attract valence e- toward it

Calculate Effective Nuclear Charge

Effective Nuclear Charge = Number of protons - Number of electrons in inner ELs

Z - S = Z/eff

Effective Nuclear Charge Trend for Periods

As you move to the right of PT, Z/eff increases because # of p+ and ve- also increases, causing there to be higher diff btwn the two and thus higher Z/eff

Effective Nuclear Charge Trend for Groups

As you move down on PT, Z/eff decreases because there are more inner ELs, making ve- far away from the nucleus and thus the e- are not able to feel the attraction from the nucleus as much

Trend for Reactivity in Periods

As you move to the right of a period (towards Fluorine), Metal reactivity decreases and Nonmetal reactivity increases because Fluorine has the most potential to gain electrons and is thus the most reactive nonmetal.

Trend for Reactivity in Groups

As you move down a group (towards Francium), Metal reactivity increases and Nonmetal reactivity decreases because Francium has the most potential to lose electrons, thus metals close to Francium are highly reactive.

Trend for Metallic Character

Metallic Character increases the more down left of the PT (to Fr) because Francium is the most reactive metal

Atomic Radius

The distance between an atom’s nucleus and valence electrons (measured as ½ the distance between 2 nuclei of like atoms)

Trend for Atomic Radius

Increases going down a group because the # of ELs increases as well.

Decreases going from left to right of a period because Z/Eff increases and pulls e- closer to the nucleus.

Ionic Radius

Radius of Ions of Elements

Trend for Ionic Radius

Increases going down a group because there are more inner ELs

Decreases going from left to right of a period because Z/Eff increases

Cation < Anion, Cation < Atom, Anion > Atom: because metals lose e- = lose EL and get smaller (cation), nonmetals gain e- = many e- repel and ion gets big (anion)

Ionization Energy

The energy needed to remove 1 electron from an atom

Trend for Ionization Energy

Decreases going down a group because as the # of ELs increases going down a group the e- are held less tightly because they are further from the nucleus, thus requiring less energy to be removed.

Increases going from left to right of a period (with many exceptions) because the # of p+ increases, making the attraction of e- to p+ (Z/Eff) stronger/higher, making removal of e- more difficult.

Increases going to top right corner of the Periodic Table (towards nonmetals)

Metals have low, Nonmetals have high

Electron Shielding

When valence electrons become blocked from the nucleus by inner shell electrons

Electronegativity

The ability of an atom to attract electrons in a bond

Trend for Electronegativity

Decreases going down a group because the more reactive metals get, the worse their ability to attract e-. Also the p+ are further away from e- due to high atomic radius, causing less attraction.

Increases going from left to right of a period because left is metals who lose e- so their ability to attract e- is low, while right is nonmetals who gain e- so their ability to attract e- is higher. Also Zeff is higher meaning higher attraction.

Increases going to top right corner of the Periodic Table (towards nonmetals)

Metals have low, Nonmetals have high