Periodic Table: Trends

Atomic Radius pg. 151  The radius of an atom (size of an atom)  Determined by the energy levels (periods on PT) & proton/electron attraction Lithium 3e - Francium 87 e -

Atomic Radius  Down a group:  AR ↑ because # of energy levels ↑  Across a period:  Right to left  AR ↑ because of ↓ in p + attraction to surrounding e - Increasing

Increasing atomic radius 16 p + in nucleus 16 e - in 3 energy levels Same Period on Periodic Table 17 p + in nucleus 17 e - in 3 energy levels 18 p + in nucleus 18 e - in 3 energy levels SOME attraction between p + in nucleus and e - Large atom MORE attraction between p + in nucleus and e - Smaller atom MOST attraction between p + in nucleus and e - Smallest atom

Trend in Atomic Radius BIGGEST

Atomic Radius Examples  Which element is larger? Explain. Silicon or Sulfur  Which element is smaller? Explain. Barium or Zirconium Zr has 5 energy levels, Ba has 6, so Zr is smaller Silicon’s p + don’t attract the e - as close as Sulfur’s p + do

Na Metal Valence e - : 1 Cl Nonmetal Valence e - : 7 Ionic radius vs Atomic Radius Now Na has a +1 charge… smaller radius And Cl has a -1 charge… larger radius Cation < Atom < Anion For radius:

Ionization Energy (IE) pg. 153  The energy required to remove one valence electron from an atom to make a cation. FKr Be C

Ionization Energy (IE)  Up a group:  IE ↑ because # of energy levels ↓ (more p + and e - attraction)  Across a period:  IE ↑ because of # of valence e - increases Ionization Energy Increasing

Trend in Ionization Energy HIGHEST

Ionization Energy: Examples  Which element has a higher ionization energy? Explain. Silicon or Sulfur  Which element has a lower ionization energy? Explain. Barium or Zirconium S requires more E to remove an e - because has 6 val e -, close to desired 8 Ba has more energy levels, so easier to take an e - away than Zr

Multiple Ionization Energies  The energy required to remove more valence e - from a positive ion 3 Be Al

 The energy change that occurs when an electron is added to an atom  Release energy (want to gain an electron)  Absorb energy (forced to gain an electron) Electron Affinity (EA) pg. 157 F Be Down a group:  EA ↑ because # of energy levels ↑ (less p + and e - attraction)  Across a period:  EA ↓ because of # of valence e - increases

Trend in Electron Affinity HIGHEST DO NOT include Noble Gases (don’t bond)

 The ability for an atom to attract electrons in a chemical bond (atoms strength)  Related to # of valence electrons  Related to p + and e - attraction Electronegativity (EN) pg. 161 F H

Electronegativity (EN) Increasing  Up a group:  EN ↑ because # of energy levels ↓ (more p + and e - attraction)  Across a period:  EN ↑ because of # of valence e - increases

Trend in Electronegativity HIGHEST DO NOT include Noble Gases (don’t bond)

Electronegativity Examples  Which element is more electronegative? Explain. Silicon or Sulfur  Which element is less electronegative? Explain. Barium or Zirconium S wants more e - because has 6 val e - only needs 2 more Ba has 1 more energy level, so less attraction with highest energy e -