5-3 Periodic Trends

1. 5-3 Periodic Trends

2. Atomic Radii · one-half the distance between the nuclei of identical atoms that are bonded together. • Atom size increases going down a group because electrons are being added to higher energy sublevels. (don’t worry about memorizing exceptions).

3. Atomic Radii

4. o Atom size decreases going across a periods left to right. • This is opposite what you would intuitively think. • Why? As you move across a period, the atom gains electrons as well as protons. Because the effective nuclear charge has increased, the electrons will therefore be pulled closer to the nucleus, causing the atom to become smaller.

5. Ionic Radii – the size of an ion. • Cation – a positive ion:Na+,Mg2+,Fe2+, Fe3+, • Formed by the loss of an electron • Cations are smaller than their neutral atom. • As an electron is lost, the effective nuclear charge is greater, causing the electron cloud to become pulled closer to the nucleus, making it smaller

6. Anion – a negative ion: Cl-, F-, N3-, … • Formed by the gain of an electron(s). • As an electron is added, the nucleus remains the same so the electrons are not drawn as close to the nucleus. Also, added electron causes repulsion among other electrons.

8. Ionization Energy • Energy required to remove one electron from an atom • 1st ionization energy, 2nd ionization energy, 3rd ionization energy, … A + energy  A+ + e-

9. Ion – an atom or group of bonded atoms that has a positive or negative charge. • Ex. Na+, Cl-, O2-, Fe3+, NH4+, CO32-, … • You have to memorize a lot of these ions…sorry • Monatomic ion – ion containing one atom • Polyatomic ion – ion containing atoms bonded together

10. Ionization Energy

11. Noble Gases have the highest ionization energies because they have a full octet

12. Group IA have the lowest ionization energies • Ionization energy increases going across a period left to right and decreases down a group • Decreases down a group because electrons in outer energy levels are easier to remove • Increases across period because more electrons in an energy level make it more difficult to get to the magic number of 8

13. It is possible to remove electrons from ions • Ex. Cu+ + energy  Cu2+ + e- • Each successive electron removal from an ion feels an increasingly stronger effective nuclear charge • Look at trend on Table 5-3 pg. 145

14. Electron Affinity • the energy change that occurs when an electron is acquired by a neutral atom

15. Most atoms release energy when they acquire an electron • A + e- → A- + energy • If energy is released, use negative sign • If energy is absorbed, use positive sign • Common unit is kJ/mol

16. Generally, the halogens gain electrons more readily; therefore, their values are most negative • This fact helps to explain their great reactivity

17. Notice electron affinities become more negative going across period • Generally, electron affinities become more negative going up a group • There are exceptions to this rule

18. It is also more difficult to add an electron to an already negative ion, so the second electron affinities are all positive. • Ex. Cl- takes less energy to form than Cl2-

20. Electronegativity • a measure of the ability of an atom in a chemical compound to attract electrons • In a chemical compound, the more electronegative atom will attract electrons more • Ex. HCl : H-Cl : electrons are pulled closer to Cl

21. The most electronegative atom on the Periodic Table is Fluorine • Electronegativity tends to increase across a period • Electronegativity tends to decrease down a group