Chemistry Chapter 5

1. Chemistry Chapter 5 The Periodic Law

2. History of the Periodic Table • Stanislao Cannizzaro In 1860 he presented a method of accurately measuring the relative mass of an atom at the 1st International Congress of Chemists. • John Newlands In 1864 he noted that when elements were arranged by increasing atomic mass, their properties repeated every 8 elements. (Law of Octaves) • Dmitri Mendeleev In 1869 he heard about the new atomic masses and decided to put them in a chemistry book he was writing.

3. Dmitri Mendeleev • Arranged elements in order of increasing atomic mass. • Predicted the existence of elements that would fill three of the spaces. • Within 15 years, those elements had been discovered. (scandium, gallium, and germanium) • The success of Mendeleev’s predictions earned him credit as the discoverer of Periodic Law.

4. Mendeleev’s Periodic Table Dmitri Mendeleev

5. Two Questions Remained! • Why could most of the elements be arranged in the order of increasing atomic mass but a few could not? • What was the reason for chemical periodicity?

6. Henry Moseley • Forty years later Henry Moseley discovered a previously unrecognized pattern. • Moseley arranged the elements in increasing order according to atomic number or the number of protons in the nucleus. • Periodic Law – The physical and chemical properties of the elements are periodic functions of their atomic numbers.

7. Modern Periodic Table • An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group. • Periodicity- Can be observed in any group (column) in the periodic table. The difference in atomic number down the noble gasses is 8, 8, 18, 18, 32. The same for the group 1-2 elements and group 13-17.

8. New Groups in the Periodic Table • Noble Gases - In 1890’s Argon and Helium were discovered first and in order to fit them into the periodic table a new group was added. • Lanthanides - Discovered in the 1900’s. They are 14 elements with atomic numbers from 58-71. These elements are very similar in properties. • Actinides - They are 14 elements with atomic numbers 90-103.

9. Electron Configuration and the Periodic Table • Generally the electron configuration of an atom’s highest occupied energy level governs the atom’s chemical properties. • The period of an element can be determined from the element’s electron configuration. • For example: Arsenic = [Ar]3d104s24p3. Arsenic is located in the 4th period of the periodic table.

10. Orbital filling table

11. Periodic Table with Group Names

12. Hydrogen and Helium • Hydrogen does not share the same properties as the elements of group 1. • Helium has the electron configuration of group 2 elements however it behaves like group 18 (Noble Gases)

13. The Properties of Group 1: the Alkali Metals • Li, Na, K, Rb, Cs, Fr • Silvery appearance • Soft (cut with a knife) • React violently with water • React with halogens to form salts • Usually stored in kerosene

14. The Properties of Group 2: the Alkaline Earth Metals Be, Mg, Ca, Sr, Ba, Ra • Harder , denser, and stronger than alkali metals. • Higher melting points • Less reactive but still too reactive to be found in nature as free elements.

15. The Properties of Groups 3-12: Transitions Elementsd-block elements • good conductors of electricity • high luster • less reactive than the alkali metals and the alkaline earth metals. • Some deviations from orderly d sub-level filling occur in groups 4-11

16. The Properties of Groups 13-18: Main Group Elementsp-block elements • Properties vary greatly • The right-hand end of the p block includes all the nonmetals except H and He. • Contains all six of the metalloids (B, Si, Ge, As, Sb, and Te) • At the left-hand end of the p block there are 8 metals.

17. The Properties of Group 17: Halogens F, Cl, Br, I, and At • Halogens are the most reactive of the nonmetals. • React vigorously with most metals to form salts. • Their reactivity is due to the seven electrons in the outer energy level. (one short of noble gas notation)

18. The Properties of the f-block Elements: Lanthanides & Actinides • Fill 4f sublevel • Total of 14 block f elements • Lanthanides – shiny metals • Actinides – radioactive; only the 1st four are found naturally

19. Determination of Atomic Radius: Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius • Radius decreases across a period Increased effective nuclear charge due to decreased shielding • Radius increases down a group Addition of principal quantum levels

20. Table of Atomic Radii

21. Ionization Energy - the energy required to remove an electron from an atom • Increases for successive electrons taken from • the same atom • Tends to increase across a period • Tends to decrease down a group Outer electrons are farther from the nucleus

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

23. Table of 1st Ionization Energies

24. Another Way to Look at Ionization Energy

25. Electron Affinity - the energy change associated with the addition of an electron • Affinity tends to increase across a period • Affinity tends to decrease as you go down • in a period Electrons farther from the nucleus experience less nuclear attraction

26. Ionic Radii Cations • Positively charged ions • Smaller than the corresponding • atom Anions • Negatively charged ions • Larger than the corresponding • atom

27. Summation of Periodic Trends

28. Table of Ion Sizes

29. 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

30. Periodic Table of Electronegativities