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Periodic Properties of the Elements Early versions of the Periodic table were constructed by Mendeleev and Meyer. We now know that the periodic properties.

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Presentation on theme: "Periodic Properties of the Elements Early versions of the Periodic table were constructed by Mendeleev and Meyer. We now know that the periodic properties."— Presentation transcript:

1 Periodic Properties of the Elements Early versions of the Periodic table were constructed by Mendeleev and Meyer. We now know that the periodic properties are due to the electronic structure of atoms. Electronic structure explains the observed trends in –Atomic size –Ionization Energy –Electron Affinity Al ZnGa CdIn Sn

2 Electron Shells Electrons in successive shells (n = 1,2,3...) produce maxima in the electron density at increasing distances from the nucleus. The inner, completed electron shells are called core shells. The outermost electron shell is called the valence shell. It may include subshells of inner shells. Example: The valence shell of 33 As includes 4s, 3d, and 4p subshells.

3 Atomic Size The radius of an atom is found from the distance between nuclei in a molecule.

4 Trends in Atomic Size Size increases going down a column of the periodic table. Size decreases from left to right in a row.

5 Trends in Atomic Size In a column, size increases with the addition of successive electron shells. Example: The alkali metals Cs 6s Rb 5s K 4s Na 3s Li 2s

6 Effective Nuclear Charge The size of an orbital decreases with increasing Z eff for the valence electrons. 2s 1s 2p Li +3

7 Effective Nuclear Charge The size of an orbital decreases with increasing Z eff for the valence electrons. Be +4

8 Effective Nuclear Charge The size of an orbital decreases with increasing Z eff for the valence electrons. B +5

9 Effective Nuclear Charge The size of an orbital decreases with increasing Z eff for the valence electrons. C +6

10 Effective Nuclear Charge The size of an orbital decreases with increasing Z eff for the valence electrons. O +8

11 Effective Nuclear Charge The size of an orbital decreases with increasing Z eff for the valence electrons. F +9

12 Ionization Energy The energy required to remove an electron from an atom in its ground level. Example: Hydrogen H(g) H + (g) + e I = E 0 E -R H 1s = 1312 kJ/mol

13 Ionization Energy Energies can be defined for successive ionizations. For Mg: Mg(g) Mg + (g) + e I 1 = 738 kJ Mg + (g) Mg +2 (g) + e I 2 = 1450 kJ Mg +2 (g) Mg +3 (g) + e I 3 = 7730 kJ For Al:

14 Trends in Ionization Energy Ionization energy increases with Z eff for the valence electrons. I 1 decreases going down a column of the periodic table. I 1 increases from left to right in a row. Alkali metals have the lowest ionization energies in a period. Rare gases have the highest.

15 Trends in Ionization Energy Arrange the following with increasing ionization energy: C, K, Mg, Na, Ne, Si K < Na < Mg < Si < C < Ne Why is the ionization energy of N greater than O? The 2p subshell of N has a slightly higher energy due to electron repulsion.

16 Electron Affinity The change in energy when an electron is added to a gaseous atom. Cl(g) + e Cl (g) E = -349 kJ/mol A large, negative E indicates strong attraction between the atom and the added electron. A positive E indicates the addition of an electron is unfavorable. Ne(g) + e Ne (g) E = 40 kJ/mol

17 Metals and Nonmetals Metals are characterized by low ionization energy; Nometals by high electron affinity.

18 Group 1A: Alkali Metals Alkali metals react to lose electrons. These metals are very reactive with reactivity increasing down the group. Reaction with H 2 O: Na(s) + H 2 O(l) NaOH(aq) + H 2 (g)

19 Groups 2A and 7A The alkaline earth metals have higher ion- ization energies than alkali metals and are less reactive. Ca + H 2 O Ca(OH) 2 + H 2 The halogens react to gain electrons. Reactivity decreases going down the group. Cl Br 2 Cl + Br 2


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