1. The Periodic Table

2. Dmitri Mendeleev (1834 – 1907) He organized elements into the first periodic table He arranged elements by increasing atomic mass

3. Henry Moseley (1913) He arranged elements according to atomic number rather than atomic mass The modern periodic table is arranged by atomic number

4. Periodic Law The periodic law states that there is periodic repetition of chemical and physical properties of elements

5. The Modern Periodic Table There are 18 groups (columns up and down) The group A’s (the tall columns) are called representative elements

6. The group B’s (the middle columns) are called transition metals

7. There are seven periods (rows across the periodic table)

8. Metals are to the LEFT of the zig-zag line (except hydrogen!) Metals in yellow

9. Nonmetals are to the RIGHT of the zig-zag line nonmetals in red

10. Metalloids Metalloids are those elements ON the zig-zag line Metalloids border the zig-zag line

11. Now...YOU fill in the chart using your book!

12. Metals solid at room temperature shiny (have luster) and smooth good conductors of heat and electricity

13. Metals malleable – “bendable” (can be pounded into sheets) ductile - can be pulled into wires

14. Metals react with acids mercury (Hg) is the only LIQUID metal

15. Nonmetals generally gases or brittle, dull looking solids at room temperature poor conductors of heat and electricity Bromine (Br) is the only LIQUID nonmetal

16. Metalloids sometimes called semimetals metalloids have properties of both metals and nonmetals

17. Metalloids silicon and germanium are two of the most important metalloids (they’re used in computer chips and solar cells)

18. Trends of the Periodic Table

19. Periodic Law If elements are organized according to atomic number, their properties will repeat periodically

20. Four Periodic Trends: 1.Atomic radii 2.Ionic radii 3.Electronegativity 4.Ionization energy

21. Atomic Radius The atomic radius basically tells you the size of the atom. It is half the distance between two nuclei of identical atoms bonded together. radius

22. The trend: atomic radii DECREASE across a period Why? Each time a positive proton is added to the nucleus, the negative electrons feel a greater attraction to the positively- charged nucleus and get “pulled in” tighter

23. Decreasing – getting smaller!

24. Trend: atomic radii INCREASE down a group Why? electrons are added to higher and higher energy levels as you go down

25. Atomic radii DECREASE down a group!

26. The farther the electrons from the nucleus, the larger the atomic radii!!!!!

27. Try these... 1.Which element has the larger atomic radius: C or F?  carbon 2.Which element has the smaller atomic radius: Ar or Kr?  argon

28. Ionic Radii Is basically the size of an ion or half the distance between the nuclei of two ions bonded together What is ion???

29. Ionic Radii Ion – an atom with a charge (+ or - ) An ion is formed when atoms lose or gain electrons

30. What happens to an atom if it LOSES an electron?  it loses a negative charge so it becomes POSITIVE Na +1

31. A positively charged ion is called a cation Na +1

32. Positive ions (cations) are smaller than the atoms they come from because they lose electrons making the atom smaller. Na Na +1

33. What happens to an atom if it GAINS an electron?  It gets more negative (so it has a negative charge) Cl -1

34. A negatively charged ion is called an anion. Cl -1

35. Negative ions (anions) are LARGER than the atoms they come from because they gain electrons – making the atom LARGER! Cl -1 Cl

36. The trend: See the board

37. Remember... all atoms want a full octet (8 valence electrons) atoms with 1 valence electron will give up that electron VERY QUICKLY to become stable

38. example: sodium has one valence electron: 1s 2 2s 2 2p 6 3s 1 if sodium gives it away, then the configuration will be: 1s 2 2s 2 2p 6 sodium will have a full octet

39. atoms with 7 valence electrons will hold on to those electrons VERY TIGHTLY they try to get one more and become stable

40. Ionization Energy The amount of energy needed to remove an electron think of it as: how tightly an atom holds on to its electrons

41. The trend: ionization energy INCREASES across a period

42. Why? the more valence electrons an element has, the more difficult it is to remove them!

43. The trend: Ionization energy DECREASES down a group

44. Valence electrons in higher energy levels are NOT held as tightly because they are farther from the nucleus Therefore, it is easier to remove an electron that is farther from the nucleus

45. Try these... Which has a higher ionization energy: Na or Cl  Chlorine Which has a lower ionization energy: Li or O  Lithium

46. Electronegativity The ability of an atom to attract electrons to itself

47. The Trend: electronegativity INCREASES across a period Why? atoms are trying harder to attract electrons to get a full octet

48. The trend : electronegativity DECREASES down a group Why? it is harder to hold on to the electrons that are farther away from the nucleus

50. Try these... Which element is more electronegative? F or Br  Fluorine Which element is more electronegative? B or Ca  Boron

51. Finished!