1. Periodic Properties of the Elements Chapter 7 AP Chemistry

2. I History and Development Elements in pure form like Gold and Silver were known thousands of years ago 31 elements were known in 1800 by 1865 the number had more than doubled to 63. The increase lead to the grouping and classification of the elements

3. First Periodic Law Properties of elements are a periodic function of atomic weight Although Meyer also grouped elements like Mendeleev Mendeleev was given more credit because he left blanks for un- discovered elements Mendeleev

4. Modern periodic law In 1913, by using x-ray spectra obtained by diffraction in crystals, he found a systematic relation between wavelength and atomic number,1913x-ray spectra diffractioncrystals wavelengthatomic number Using Atomic number Mosley rearranged the elements noting that the elements were a periodic function of atomic number Henry Mosley

5. Design of the Table Rows and Columns Horizontal rows – series or periods – numbered 1-7 Notes the shell or principle QN Vertical columns – group or a family – numbered 1-18 or 1A-8A, 1B-8B Review electron configuration notation

6. Periodic Trends Size of Atoms Within each group the atomic radius tends to increase as you go down the group – the principle QN increases – size of the cloud gets larger Within a period the atomic radius decreases as you go from left to right – the effective nuclear charge increases – the number of protons increases while the number of core electrons stays the same – the outer electrons don’t shied each other well

7. Ionization Energy (I) I – the energy required to remove the outermost electron from a gaseous atom or ion I 1 – the energy to remove one electron – I 2 – the second electron Metals tend to give up e- to obtain a Nobel gas configuration

8. Ionization Energy Trends Ist Ionization energy generally follows effective nuclear charge – increases from left to right in periods and from bottom to top in groups Except from group 2A to 3A there is a slight increase – the second e- in the S sub-shell is harder to remove that the 1 st P e- GR 5A to 6A there is a slight increase due to repulsion of paired e-s in the P4 configuration Every element shows large increase in IE when e’s are removed from a Nobel gas core

9. Electron Affinities Is the energy change of the reaction of adding an electron to a gaseous atom or ion. Tends to be an exothermic processes (Delta H is neg) although in some cases it is positive or endothermic In general electron affinity tends to decrease ( become more negative) from right to left in periods. There is very little change going down a group in the value.

10. Properties of Metal, Non-Metals & Metalloids Metals – conduct heat and electricity are lustrous, malleable, and ductile Tend to lose e’s to become cations Form basic oxides – metal oxides react with water to produce basic or alkaline solutions Ex 7.7,7.8 in text

11. Metal Properties Cont. Metal oxides + an acid yield a salt and water Ex 7.9,7.10 in the text

12. Non-Metal Properties Poor conductors of heat and electricity are often dull in color and shatter if forged Nonmetals tend to gain electrons form other sub and become anions Non-metals from acidic oxides – non-metal oxide react with water to form acidic solutions Ex 7-12, 7-13 in text Non-metal oxides + a base yield a salt plus water ex 7.14, 715 in text.

13. Mettaloids Boundary between metals and non-metals Mettaloids can either gain or lose electrons Al and Po are metals not mettaloids Properties between metals and non-metals Several are semiconductors and re important because of their use in circuits and computer chips Metallic character – increases from left to right and from top to bottom in a group

14. Group Trends Alkali Metals Soft, gray metals Low Ionization energies – lose e’s easily – form +1 cations – very reactive Electrolysis is a technique to force an electron back on a cation to produce a neutral metal – the metals can be obtained

15. Alkali Metals by passing an electrical current through a molten salt. The metals combine with hydrogen to from hydrides ex. 7.19, or with Sulfur to form Sulfides ex. 7.20 or with chlorine to form chlorides. ex 7.21 The alkali metals react with water to produce hydrogen gas and hydroxides ex 7.22. See examples of reactions with oxygen ex 7.23-7.25

16. Alkaline Earth Metals Slightly harder and more dense than Alkali metals 1 st ionization energies are low but not a s low as alkali metals form 2+ cations Reactivity tends to increase as you progress down a group. See examples of reactions 7.26 – 7-29 Got Milk?

17. Trends in Selected Non-metals Hydrogen Usually placed in grp 1A because of its 1s electron config. Although a non-metal It reacts with other non-metals to form molecular cmps ex 7.30-32 Reacts with metals to form hydrides 7.31-32

18. Oxygen Family Non-metallic – top of family – metallic at bottom O, S, Se are typical non-metals – O and S are allotropes – molecules of elements with different structures ( O 2, O 3, S 2, S 4, S 6, S 8 ) Polonium is a rare and radioactive element See examples of reactions 7.33-36

19. The Halogens All non-metals (astatine is radioactive – seldom discussed) All are diatomic in gas phase Bromine normally a liquid – iodine is normally a solid See reactions 7.37-42

20. Nobel Gases Mostly non-reactive Some compounds of Xenon and Krypton are known All monatomic gases Highest ionization energies of any family xenon tetrafluoride