The document summarizes the development of the periodic table of elements from early classifications by Dobereiner and Newlands to Mendeleev's and Meyer's published periodic tables to Moseley's establishment of the ordering by atomic number. It explains key periodic properties including atomic structure and trends in atomic size, and provides an overview of the layout and information contained in the modern periodic table.

1. The Periodic Table of The Elements

2. Elements are like a collection As more and more elements were discovered it became more important to organize and classify them

3. Between the late 1700’s and mid 1800’s scientists, using mostly atomic spectroscopy, doubled the number of known elements.

4. In early 1800’s, German chemist J.W. Dobereiner observed that several of the elements could be classified into groups of three which he called triads

5. Ca, Sr, and Ba Li, Na, K Cl, Br, I Dobereiner based his triads on similar chemical properties

6. In addition: Many of the properties of the middle element in each triad are approximate averages of the properties of the first and third element

7. Element Atomic Density Mass (amu) Cl 35.5 1.56 g/L Br 79.9 3.12 g/L I 126.9 4.95 g/L Ca 40.1 1.55 g/cm 3 Sr 87.6 2.6 g/cm 3 Ba 137 3.5 g/cm 3

8. In 1865, English chemist J.A.R Newlands presented another way to classify and organize the 62 elements known at the time

9. Newlands placed the elements in order of increasing atomic mass He noticed that the properties of the eighth element were like those of the first, the ninth like those of the second, and so on….

10. He called this repeating pattern of every eight elements THE LAW OF OCTAVES After the eight notes of the musical scale

11. Because he linked chemistry to music, he was not taken seriously! It took 20 years for him to receive credit for recognizing periodicity

12. In 1869 Russian Chemist Dimitri Mendeleev and German chemist Lothar Meyer published nearly identical ways of classifying

13. But Mendeleev is generally more credited with the 1 st periodic table for 2 reasons: He published first He was better at explaining it than Meyer

14. Mendeleev also saw the “periods” Credited with publishing the first “periodic table”

16. Mendeleev got lots of credit because he left gaps for missing elements!

17. In 1913, English Chemist James Moseley Presented a way of organizing the elements that we still use today.

18. Moseley was the first to put the elements in order of increasing Atomic #

19. Atomic # represents the number of protons in the nucleus of each element

20. When he did this, he saw the same repeating periodic pattern, without the exceptions that Mendeleev had to switch around

21. Periodic Law When the elements are arranged in increasing order by their atomic numbers, their properties repeat periodically

22. The Modern Periodic Table

23. Each Square Tells about a different element H Element Symbol Hydrogen Element Name Not always there 1 Atomic Number Represents the number of PROTONS in each atom of this element 1.009 Atomic Mass Represents the number of PROTONS AND AVERAGE NUMBER OF NEUTRONS in each atom of this element

24. Atomic Number 1: 1 proton (positive charge particles) in nucleus of EVERY hydrogen atom Atomic Number 1: ALSO means 1 electron (negatively charged particle) OUTSIDE the nucleus Atoms are NEUTRAL: Same number of positives as negatives Atomic Mass of 1.009 means that there is an average of 1.009 PROTONS AND NEUTRONS in the nucleus… Since there is already 1 proton (AND THAT CAN NOT CHANGE AND STILL BE HYDROGEN) That means that the average atom of Hydrogen has 0 Neutrons! H Hydrogen 1 1.009

25. Let’s do another 4 Particles 2 MUST BE Protons 4-2 = 2 Neutrons in nucleus He Helium 2 4.003

26. Each block tells us about the element… The periodic table arranges the elements into rows and columns based on similarities

27. Vertical columns are called groups. Elements are placed in columns by similar properties. Also called families

28. Alkali Metals Group 1 ALL elements in Group 1 have 1 electron in their outer region

29. Alkali Earth Metals Group 2 ALL elements in Group 2 have 2 electrons in their outer region

30. Boron Family Group 13 ALL elements in Group 13 have 3 electrons in their outer region

31. Carbon Family Group 14 ALL elements in Group 14 have 4 electrons in their outer region

32. Nitrogen Family Group 15 ALL elements in Group 15 have 5 electrons in their outer region

33. Oxygen Family A.K.A. the Chalcogens Group 16 ALL elements in Group 16 have 6 electrons in their outer region

34. Halogens Group 17 ALL elements in Group 17 have 7 electrons in their outer region

35. Noble Gasses Group 18 ALL elements in Group 18 have 8 electrons in their outer region

36. Lanthanide Series Actinide Series

37. Horizontal rows are called “Periods” Each period also shows something in common

39. Period 1 Each element has one region of space for electrons around it

40. Period 2 Each element in this period has two regions of space around it for electrons

41. Period 3 Each element in this period has 3 regions of space around it for electrons

42. Period 4 Each element in this period has 4 regions of space around it for electrons

43. Period 5 Each element in this period has 5 regions of space around it for electrons

44. Period 6 Each element in this period has 6 regions of space around it for electrons The Lanthanides are part of period 6

45. Period 7 Each element in this period has 7 regions of space around it for electrons The Actinides are part of period 7

46. The periodic table can also show larger “Groups”

47. METALS Non-METALS METALLOIDS

48. Tall columns are collectively referred to as the “representative elements”

49. Short, center groups are collectively referred to as the “transition elements”

50. Two long rows on bottom are collectively referred to as the “inner transition metals”

51. Periodic Trends

52. Atomic Size First problem where do you start measuring. The electron cloud doesn’t have a definite edge. They get around this by measuring more than 1 atom at a time.

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

54. 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.

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

56. 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