University:
Carleton University
Course:
CHEM 1001 | General Chemistry I
Academic year:

2024
Views:

222

Pages:

2
Author:

ali

PERIODIC TRENDS LECTURE 1 Periodic Table: Periodic trends CHARGE: Nuclear Charge (Z): no. of protons in the nucleus Effective Nuclear Charge (Zeff): positive charge felt by an electron in a multi-electron atom Electrons in the outer orbitals are partially shielded from the nuclear charge resulting in a weaker attraction and a lower Zeff Electrons closest to the nucleus feel a charge closer to Z Further out, < Z ATOMIC RADIUS: Across a period, Z eff ↑ Electrons are attracted more strongly by the nucleus Smaller radius Down a group Electrons are added to orbitals that are further from the nucleus Larger radius ANIONIC RADIUS: When forming an anion, electrons are usually added to the same orbital Z is unchanged, Zeff is unchanged More electrons = more electronic repulsion For the same reasons, anions with -2 charge are larger than anions with -1 charge CATIONIC RADIUS: Valence orbital is closer to the nucleus Cation is smaller than the atom Ionization Energy (E) E1 is the amount of energy required to remove an electron (e-) completely from an atom (in the gaseous state) Ionization energy is always positive Large Zeff: e- is more tightly held to the nucleus Harder to remove e- Higher E1 Across a period, Zeff ↑ e- are held more tightly ↑ E1 Down a group Weaker nuclear attraction ↓ E1 The first E1 is the energy required to remove the first e- from a neutral atom. Electron affinity (Ea) is the energy change when an e- is added to an atom in the gas phase. If energy is released when an e- is added, the atom has an affinity for e- (negative affinity) Metals have a low E1 (readily lose e-) Non-metals have an affinity for e- (readily gain e-)