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 -