NOTES ON PERIODIC TRENDS 6.1 Effective Nuclear Charge ( Z eff )

1. NOTES ON PERIODIC TRENDS 6.1 Effective Nuclear Charge (Zeff) • In a many-electron atom, each electron is attracted to the positively charged nucleus and repelled by the other negatively charged electrons. • Zeff takes both of these factors into account and represents an estimate of the net electric field experienced by an electron.

2. Effective Nuclear Charge (Zeff) • The formula we use is : Zeff = Z – S • Z = number of protons in the nucleus. • S = screening constant; the number of core electrons that screen the outer electrons from the positive charge in the nucleus.

3. Estimate the effective nuclear charge experienced by the outer electron in each of the following atoms: • Potassium • Fluorine • Silicon

4. Increases (b/c number of protons increases) Effective nuclear charge

5. Decreases Atomic Radius– represents one half the distance between the nuclei of atoms held together by a bond. Increases

6. As we move from left to right across a period, the atomic radius decreases. • Why does this occur? • as we move from left to right, atomic number increases. • Since there are more protons in the nucleus, the effective nuclear charge increases. • This increases the attractive force felt by the outer e- making the atom smaller.

8. As we move from the top of a group to the bottom of a group, the atomic radius increases. • Why does this occur? • as we move down a group, the principal quantum number (n) increases. • This means that the outer e- are more likely to be found in higher energy orbitals further from the nucleus.

9. Ionic Radii • When atoms gain an electron(s), they become negatively charged anions. Anions are larger than the corresponding neutral atom. X– will be larger than X because: • Negatively charged ions have more electrons in the electron cloud. • Each electron repels the other negatively charged electrons, causing the electron cloud to “spread out”. • More electron-electron repulsions leads to a larger radius than a neutral atom would have.

10. Ionic Radii • When atoms lose an electron(s), they become positively charged cations. Cations are smaller than the corresponding neutral atom. M+ will be smaller than M because: • Positively charged ions have less electrons in the electron cloud. • Less electrons means there will be fewer electron-electron repulsions, making the ion smaller than a neutral tom would be.

11. Increases Ionization Energy– the energy required to remove an electron from the ground state of a gaseous atom or ion. Decreases

12. As we move from left to right across a period, the ionization energy increases. • Why does this occur? • as we move from left to right, atomic number increases. • Since there are more protons in the nucleus, the effective nuclear charge increases. • This increases the attractive force felt by the outer e- making it harder to remove.

13. As we move from the top of a group to the bottom of a group, the ionization energy decreases. • Why does this occur? • as we move down a group, the principal quantum number (n) increases. • This means that the outer e- are more likely to be found in higher energy orbitals further from the nucleus. • This makes them less attracted to the nucleus so they take less energy to remove.

14. Successive Ionization Energies • Removing a second (or third, etc.) electron always requires even more energy than removing the first electron. • This occurs because the atom gets smaller each time an electron is removed and this makes the outer electrons closer to the nucleus (and more strongly attracted to it).

15. The ionization energies for aluminum are summarized in the table below. • Why is there such a large jump in the amount of energy required for I4? • Once aluminum has lost 3e- it is Al3+ and has an octet of 8 valence electrons. • According to the octet rule this is the most stable state for an atom, and removing one of it’s octet of valence electrons would require a very large amount of energy.

16. Electron Affinity– the energy change that occurs when an electron is added to a gaseous atom.

17. A negative affinity indicates a high attraction for an electron (energy is released) • A positive affinity indicates that an electron is not likely to be gained by the atom. • In general, electron affinity increases moving from left to right across a period…however it shows a lot more variation than the other trends. • Therefore, it is more helpful to look at the trends for different groups of elements.

18. Group Trends for Electron Affinity • Halogens have highly negative electron affinities because they have the greatest attraction for electrons since they only need to gain one electron to reach a full outer shell. • Noble Gases have positive electron affinities because they already have an octet of valence electrons and would have to add an electron to an unoccupied higher energy level. • Group 5A sees a decrease in affinity because the p-sub-shell is half-full and adding another electron will result in increased electron-electron repulsions in the atom.

19. Electronegativity – the ability of an atom in a molecule to attract electrons to itself.

20. Electronegativity tends to increase as we move from left to right across a period and tends to decrease as we move down a group. • The most electronegative elements have highly negative electron affinities (attraction for additional electrons) and high ionization energies (tendency to hold onto their won electrons).