Chapter 5

1. Chapter 5 The Periodic Law

2. Section 5-1History of the Periodic Table

3. History of the Periodic Table • Stanislao Cannizzaro In 1860 presented a method for accurately measuring the relative masses of atoms at the 1st International Congress of Chemists. • Dmitri Mendeleev heard about the new atomic masses and included the new values in a chemistry textbook he was writing.

4. Dmitri Mendeleev

5. Dmitri Mendeleev • Original table arranged elements in order of increasing atomic mass. • This “periodic table” was first published in 1869. • Left gaps in table where he predicted the existence of new elements • Credited as the discoverer of the Periodic Law. • “Periodic” means “having a repeating pattern.”

6. Mendeleev’s Periodic Table

7. 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?

8. Henry Moseley • In 1911, 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.

9. Modern Periodic Table • Arrangement of the elements in order of atomic number so that elements with similar properties fall in the same column or group.

10. New Groups in the Periodic Table • Noble Gases - In 1890s argon and helium were discovered first. Group 18 added, gases that are generally unreactive. • Lanthanides - Discovered in the 1900s. 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.

11. Periodicity • Periodicity with respect to atomic number can be observed in any group of elements. • Group 1 elements very similar with respect to their properties, but are different from each other and different from those in other groups. • Groups 1-2, and 13-18 follow a pattern of 8, 8, 18, 18, and 32 with the exception of hydrogen.

12. Section 5-2Electron Configuration and the Periodic Table

13. Electron Configuration & PTOE • Generally, 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 e- config Example: Arsenic = [Ar]3d104s24p3 Arsenic located in the 4th period of the periodic table.

14. Orbital filling table

15. Periodic Table with Group Names

16. Hydrogen and Helium • Hydrogen does not share the same properties as the elements of group 1 (1s1) • Helium hase- config of group 2, 1s2 - behaves like group 18 (Noble Gases) - chemically stable, nonreactive • First energy level only need 2 e- to be filled

17. Group 1: 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

18. Group 2: 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.

19. 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 are very unreactive (ex: gold, platinum

20. 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.

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

22. f-block Elements: Lanthanides & Actinides • Fill 4f sublevel • Total of 14 block f elements • Lanthanides – shiny metals, react similar to alkaline-earth metals. • Actinides – all radioactive; only the first 4 have been found naturally on Earth

23. Section 5-3Electron Configurations & Periodic Properties

24. Atomic Radii Atomic Radius Distance between nuclei Atomic radius – one-half the distance between the nuclei of identical atoms that are bonded together

25. Period Trends • Decreases across a period

26. Why? Protons are added to the nucleus moving across a period from left to right This increases the charge of the nucleus (effective nuclear charge – Zeff) As Zeff increases, electrons are pulled closer to the nucleus

27. + + + + + + + + + + Period Trends

28. Group Trends • Increase down a group

29. n=3 n=2 n=1 Why? • Electron-electron repulsion “plumps” up the atom • Zeff decreases the further the electrons are from the nucleus The addition of shells increases the electrons’ distance from the nucleus and the size of the atom

30. Atomic Radii Trends DECREASES DECREASES

31. Ion: an atom or group of bonded atoms that has a negative or positive charge. Ex: Na+Cl- Cr3+ Polyatomic ions SO42- OH- NH4+ Ionization: Any process that results in the formation of an ion. How do ions form? - Loss or gains of e- Metals lose e-Nonmetals gain e-

32. ALL Periodic Table Trends Influenced by three factors: 1. Energy Level Higher energy levels are further away from the nucleus. 2. Nuclear Charge (# protons) More charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect (e- blocking effect)

33. Ionization Energy • The energy required to remove one electron from a neutral atom of an element A + Energy  A+ + e-

34. Ionization Energy Trends Across Period • increases LR • increase in # of protons attracting the e- • becomes harder to remove the e- • requires more energy so IE increases Down Group • decreases down a group • valence e- farther away from p+ attracting them • easier to remove e- • requires less energy so IE decreases

35. Successive Ionization Energies IE increases for each successive e- removed from the same atom • 2nd ionization energy (IE2) is energy required to remove 2nd e- • Always greater than first IE • IE3 is energy required to remove 3rd e- • Greater than IE1 or IE2

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

37. Table of 1st Ionization Energies

38. Ionization Energy Trends

39. Electron Affinity • The change in energy that a neutral atom undergoes when an electron is acquired (the ability to attract an e -) • Most atoms release energy when they gain an e- A + e- A- + energy negative energy value (exothermic) • Halogens (Group 17) gain e- the easiest, very reactive

40. Electron Affinity Trends Period In general as e- add to the p block, • Electron affinity increases, becomes more negative • Going LR greater nuclear charge Down Groups • Tends to decrease down a group. • e- shielding causes decrease in nuclear attraction for e-

42. Summation of Periodic Trends

43. Electronegativity Is a measure of the ability of an atom in achemical compoundto attract electrons Valence electrons hold compounds together Uneven electrical charge is important in compound formation and other chemical properties

44. Electronegativity NOTE • Electronegativity is a property of atoms in compounds and differs from ionization energy and electron affinity, which are properties of isolated atoms

45. Electronegativity Trends Across Periods • Increases LR across a period. • Increases due to nuclear charge • Fluorine most electronegative Down Groups • Decreases or stays the same • Decrease due to e- shielding; greater nuclear charge • Noble gases don’t form cmpds, so aren’t assigned electronegativity #”s

46. Periodic Table of Electronegativities

47. Electronegativity Trends

48. Valence Electrons • Electrons available to be gained, lost, or shared when forming chemical compounds • Located in outer energy level - Main group elements = outer s and p sublevels