Periodicity-trends of the elements on the periodic table

1. Periodicity-trends of the elements on the periodic table Mr. Guerrero LFHS

2. The 2 reasons for most trends: • 1) Effective Nuclear Charge(Zeff)-the ability(pull) of the nucleus to attract its valence electrons. This increases with increasing numbers of protons in the nucleus. • The Zeff is explained by Colomb’s Law: Where K = 8.99 x 109 N m2/C2 • 2) Shielding Effect-the interference on effective nuclear charge by core electrons. This effect increases as energy levels increase. Shielding has no effect along a period, only within groups! • 1) Effective Nuclear Charge(Zeff)-the ability(pull) of the nucleus to attract its valence electrons. This increases with increasing numbers of protons in the nucleus. • The Zeff is explained by Colomb’s Law: • 2) Shielding Effect-the interference on effective nuclear charge by core electrons. This effect increases as energy levels increase.

3. Periodic Trend #1: • Atomic Radius- ½ the distance between the nuclei of two identical bonded atoms. • Trend Atomic Radius increases toward He.

4. Ionic Radius-size of cation or anion LEO/GER

5. Arrange by increasing atomic radius • Al, Sr, F, P, Ba • K+, K, N3-, Mg, Mg2+, Ca2+, Na, Na+

6. Why do d-orbitals have n-1 • Penetration effect- the effect whereby a valence electron penetrates the core electrons, thus reducing the shielding effect and increasing the pull from the nucleus. (The electron is temporarily closer to the nucleus than normal.) • most penetration ns>np>nd>nfleast penetration This means that the s electrons “penetrate” or come closer to the core more often than do p, d, or f electrons. • Electrons fill orbitals in order of increasing energy. Because of the penetration effect, electrons fill (n+1)s before nd. ((n+1)s has lower energy). For example, 4s fills before 3d.

7. Electronegativity- the attraction of a nucleus for electrons. Trend: electronegativities increase toward Fluorine • Arrange by decreasing electronegativity: Cs, Si, F, Ca, Ga

8. Electron Affinity (electron loving) • Electron Affinity- the energy change associated when an atom, in the gaseous state gains an electron. Electron affinity values are always negative. • Rxn: X(g) + e- X-(g) DE = ???? • Trend: same as electronegativity

9. First Ionization Energy • First Ionization Energy – The energy required to remove the first electron from a neutral atom, in the gaseous state. • Rxn: X(g)  X+(g) + e-DE = ???? Trend:Valence electrons in smaller atoms are harder to remove, because they are closer to the nucleus. Therefore: He has the highest Ionization energy!

11. Middle of the p-block exception: Exceptions to the rule Ionization energy and electron affinity have several examples of exceptions. This includes exceptions within the p-block that you are responsible for knowing. Atomic radius and electronegativity do not have as many exceptions, and none are notable enough that you must know.

12. p-block exceptions: e- affinity & 1stIoniz. E.-As e-’s fill the p-block, the 4th e- must enter the first occupied orbital. The previous e- causes a large e-/e- repulsion, which causes an anomaly in ionization energy and electron affinity trends.

13. Ionization energy actually drops between group2-13(Be-B) AND groups 15-16(N-O)

14. Use arrows to show the increase in: electronegativityatomic radiuse- affinityfirst ionization energy

15. Second Ionization energy(IE2)- the energy required to remove the second electron from a neutral atom, in the gas phase. • After the removal of each electron, the electrons get closer to the nucleus, due to less e-/e- repulsion. • The energy required to remove the next electron is always greater than the energy required to remove the previous electron. IE1 < IE2 <IE3 < IE4….

16. Ionization Energies in kJ/mol

17. Work out AP released 2007, form B, #6

18. Polarizability-the distortion of a nonpolar electron cloud to a temporary dipole, by the presence of a near by atom. • As the total number of electrons increases, polarizability also increases.