Periodic Table

1. Periodic Table

2. Periodic Table • Placed in order of their atomic numbers • The similar elements are placed in columns, known as groups or families • The rows of the table are called periods

3. How did Mendeleev’s periodic table differ from the modern periodic table? • The elements in Mendeleev’s table were listed in order of atomic weights rather than atomic numbers. Atomic numbers were unknown in 1871. None of the internal structure of the atom was known in 1871. • Noble gases are completely absent--they were not known in 1871. The first noble gas discovered in 1898 was argon.

4. Representative Elements • Main Group Elements – Outer subshell is an s and p in the highest energy level • Group 1a – Alkali • Group 2a – Alkaline Earth • Group 3a – Boron Family • Group 4a – Carbon Family • Group 5a – Nitrogen Family • Group 6a – Oxygen Family • Group 7a – Halogens • Group 8a – Noble Gases • Write the final energy level for each group

5. Groups 1b to 8b • Transition Metals • Incompletely filled d subshell

6. Valence Electrons • Outer electrons of an atom, which are involved in bonding

7. Ionization Energy • Ionization energy is defined as the energy required to remove an electron from a gaseous atom. X(g) + energy → X+(g) + e-

8. For every row of the periodic table, the ionization energy increases from left to right (with some exceptions) reaching a maximum on the noble gases. It then starts low again for the next alkali metal (Group IA) and gradually increases again reaching a maximum on the next noble gas.

9. The periodic trends are: Horizontal: IE increases across a row of the periodic table (from left to right) Vertical: IE decreases down a group (column) of the periodic table (from top to bottom)

10. What are the reasons for these trends? • Horizontal: As you go across from left to right in a row of the periodic table, you are adding electrons to the same shell but with increasing nuclear charge. The increasing number of protons (higher Z) attracts the electrons more, making it harder to remove an electron from the atom--hence a higher IE. • Vertical: As you go down a group from top to bottom, you always have the same valence shell configuration. However, each succeeding shell is further from the nucleus, and is shielded from the nuclear pull by inner electrons. It is thus easier to remove electrons from outer shells which are less attracted to the nucleus.

11. Atomic Radius • Atomic radii measure the size of the atom. Although the atom does not have a distinct boundary, there are several ways to estimate atomic radii based on distances between atoms in crystals or molecules.

12. Atomic radii show distinct trends. • Horizontal: atomic radius decreases across a row of the periodic table (from left to right) • Vertical: atomic radius increases down a group (column) of the periodic table (from top to bottom)

13. What are the reasons for these trends? • Horizontal: As you go across from left to right in a row of the periodic table, you are adding electrons to the same shell but with increasing nuclear charge. The increasing number of protons (higher Z) attracts the electrons more, making it the electron cloud closer to the nucleus; hence a smaller atomic radius.

14. Vertical: As you go down a group from top to bottom, you always have the same valence shell configuration. However, each succeeding shell is further from the nucleus, and is shielded from the nuclear pull by inner electrons. So the outer shell logically is larger there are more shells (higher principal quantum number).

15. Most atoms form ions. Which is larger, the ion or the atom from which it is formed? • Positive ions are formed when the atom gives off an electron. The resultant ion is always smaller than the corresponding atom, since the resultant positive charge (more protons than electrons) causes the remaining electron cloud to be pulled toward the nucleus.

16. Negative ions are formed when the atom gains an electron. The resultant ion is always larger than the corresponding atom, since the resultant negative charge (more electrons than protons) causes the electron cloud to be repelled away from the nucleus.

17. Electron Affinity • When an electron is added to a gaseous atom, forming a negative ion, energy may be either released or absorbed. • The ΔH for this process is defined as the electron affinity. X(g) + e- → X- (g) ΔH = electron affinity of X

18. What are the trends in electron affinity? • Horizontal: Going from left to right across a row, the electron affinity gets more negative (more attraction for electrons), with the halogens (not the noble gases) having the most negative electron affinity (most attraction for electrons). • Vertical: Going from top to bottom down a group, the electron affinity gets less negative (less attraction for electrons.

19. What are the reasons for these trends? • Horizontal: As you go across from left to right in a row of the periodic table, you are adding electrons to the same shell but with increasing nuclear charge. The increasing number of protons (higher Z) attracts the electrons more; a higher attraction for electrons means a more negative electron affinity.

20. Vertical: As you go down a group from top to bottom, you always have the same valence shell configuration. However, each succeeding shell is further from the nucleus, and is shielded from the nuclear pull by inner electrons. Less attraction by the nucleus means a less negative electron affinity.