Unit 5: Periodicity

1. Unit 5: Periodicity Cartoon courtesy of NearingZero.net

2. History • Russian scientist Dmitri Mendeleev taught chemistry in terms of properties. • Mid 1800 - molar masses of elements were known. • Wrote down the elements in order of increasing mass. • Found a pattern of repeating properties.

3. Mendeleev’s Table • Grouped elements in columns by similar properties in order of increasing atomic mass. • Found some inconsistencies - felt that the properties were more important than the mass, so switched order. • Found some gaps. • Must be undiscovered elements. • Predicted their properties before they were found.

4. Mendeleev’s Periodic Table Dmitri Mendeleev

5. The modern table • Elements are still grouped by properties. • Similar properties are in the same column. • Order is in increasing atomic number-Moseley • Added a column of elements Mendeleev didn’t know about. • The noble gases weren’t found because they didn’t react with anything.

6. Modern Russian Table

7. A Spiral Periodic Table

8. “Mayan” Periodic Table

9. Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius

10. Trends in Atomic Size • Influenced by two factors. • Energy Level • Higher energy level is further away. • Charge on nucleus • More charge pulls electrons in closer.

11. Group trends H • As we go down a group • Each atom has another energy level, • So the atoms get bigger. Li Na K Rb

12. Periodic Trends • As you go across a period the radius gets smaller. • Same energy level. • More nuclear charge. • Outermost electrons are closer. Na Mg Al Si P S Cl Ar

13. Table of Atomic Radii

14. Ionic Radii • Positively charged ions formed when • an atom of a metal loses one or • more electrons Cations • Smaller than the corresponding • atom • Negatively charged ions formed • when nonmetallic atoms gain one • or more electrons Anions • Larger than the corresponding • atom

15. Table of Ion Sizes

16. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativities tend to increase across • a period • Electronegativities tend to decrease down a • group or remain the same How fair it shares. Big electronegativity means it pulls the electron toward it. Atoms with large negative electron affinity have larger electronegativity.

17. Group Trend • The further down a group the farther the electron is away and the more electrons an atom has. • More willing to share. • Low electronegativity.

18. Periodic Trend • Metals are at the left end. • They let their electrons go easily • Low electronegativity • At the right end are the nonmetals. • They want more electrons. • Try to take them away. • High electronegativity.

19. Periodic Table of Electronegativities

20. Ionization Energy - the energy required to remove an electron from an atom • Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove

21. Tends to decrease down a group Outer electrons are farther from the nucleus • Increases for successive electrons • taken from the same atom

22. Table of 1st Ionization Energies

23. Ionization of Magnesium Mg + 738 kJ  Mg+ + e- Mg+ + 1451 kJ  Mg2+ + e- Mg2+ + 7733 kJ  Mg3+ + e-

24. Summation of Periodic Trends