1. Periodic Table Chapter 6

2. Periodic Table Many different versions of the Periodic Table exist All try to arrange the known elements into an organized table

3. Alternate Periodic Tables

10. Elements known since Ancient times

11. Elements Discovered in 1600’s

12. Elements Discovered in 1700’s

13. Elements Discovered 1800-1810

14. Elements Discovered 1810-1863

15. Elements Discovered 1875-1899

16. Elements Discovered 1900-1940

17. Elements Discovered 1944-1961

18. Elements Discovered 1966-1996

19. Elements Discovered since1999

20. History 1869 - Russian chemist and teacher, Dmitri Mendeleev proposed a table for organizing elements Mendeleev arranged the elements in a table based on increasing atomic mass.

21. History Mendeleev placed elements next to each other with similar chemical properties He would leave elements out of order based on atomic mass if they lined up better based on chemical properties

22. History Mendeleev left spaces for elements not yet discovered – He predicted properties of elements that would fit in those spots He predicted very closely the properties of Ge, Ga, Sc, and 5 others

23. History 1913 - British physicist, Henry Moseley, determined the atomic numbers for the elements The modern periodic table is arranged in order of increasing atomic number.

24. Periodic Table

25. Arrangement Columns are called Groups – Numbered 1-18 Rows are called Periods Elements in the same group have similar properties

26. Group Names Group 1 - Alkali Metals Group 2 - Alkaline earth metals Group 17 - Halogens Group 18 - Inert or Noble gases.

27. Group Names Groups 3-11 – Transition Metals Bottom 2 rows – Inner Transition

28. Phases at STP Most elements are solids at STP Hg and Br are liquids at STP H, N, O, F, Cl and Noble Gases are all gases at STP

30. Periodic Law Periodic Law – When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

31. Valence Electrons Electrons in outermost occupied energy level Valence Electrons are responsible for most chemical properties – Elements in the same group have similar properties because they have the same number of valence electrons

32. Classifying Elements Elements are classified into 3 groups based on their properties: Metals – Left and Middle Nonmetals – Right Metalloids - Staircase

33. Metals Good conductors of heat and electrical current High luster or sheen Many are ductile, meaning they can be drawn into wires Most are malleable, meaning they can be hammered into thin sheets

34. Metals Metallic Character increases as you move towards the lower left Most Metallic Element is Francium, Fr

35. Nonmetals Most are gases at room temperature, some are solids, and one is liquid Most are poor conductors Most solids are brittle

36. Nonmetals Non-Metallic Character increases as you move towards upper right Most nonmetallic element is Fluorine, F

37. Metalloids B, Si, Ge, As, Sb, Te Have properties of both metals and nonmetals, based on conditions Exceptions: – Al and Po are metals – At is a nonmetal

39. Group Characteristics Alkali Metals (Group 1) – H, Li, Na, K, Rb, Cs, Fr – All have 1 valence electron, tend to form +1 ions – Most reactive metals – Not found in nature by themselves, always combined with someone else – Have properties of metals but are softer and less dense

40. Group Characteristics (cont) Alkaline Earth Metals (Group 2) – Be, Mg, Ca, Sr, Ba, Ra – All have 2 valence electrons, tend to form +2 ions – Harder and more dense than alkali metals, but also have higher melting and boiling points – Highly reactive, but not as much as alkali metals – Not found by themselves in nature

41. Group Characteristics (cont) Halogens (Group 17) – F, Cl, Br, I, At – All have 7 valence electrons, tend to form -1 ions – Strongly non-metallic – Most active nonmetals – Have low melting and boiling points – Combine readily with metals to form salts

42. Group Characteristics (cont) Noble Gases (Group 18) – He, Ne, Ar, Kr, Xe, Rn – Colorless gases that are extremely non-reactive – Full valence shell, non-reactive – All are found in small amounts in our atmosphere

43. Group Characteristics (cont) Transition Metals (Groups 3-11) – Most are excellent heat and electrical conductors – Most have high melting points and are hard, except Hg – Less active than group 1 and 2 metals – Many combine with Oxygen to form oxides (Chemical property) – Many have more than one oxidation number – Form compounds that are colorful

44. Reminder STP Standard Temperature and Pressure – 1 atm, 0°C Reference Point for most measurements

45. Diatomics Eight elements are diatomic molecules when alone in nature (exist as two atoms bonded together) H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, At 2

46. Diatomics Hydrogen and the Magic 7

47. Coloring Color in the specific groups with your own color choices

48. Coloring Color in the different classifications with your own color choices Metals, Nonmetals, Metalloids

50. Orbital Blocks s - block p - block d - block f - block

51. Periodic Trends How a property changes either across a period or down a group – Atomic Number – Atomic Mass – Atomic Radius – Ionic Radius – Ionization Energy – Electronegativity

52. Trends Atomic number increases across a period. – Increasing number of protons Atomic number increases down a group – Increasing number of protons

53. Trends Atomic mass generally increases across a period. – Increasing protons, neutrons, and electrons. Atomic mass increases down a group. – Increasing protons, neutrons, and electrons.

54. Radius Atomic Radius – measure of the size of the atom – Half the distance between two nuclei Ionic Radius – measure of the size of an ion

55. Trends Atomic Radius decreases across a period – More protons to pull on the electrons Atomic Radius increases down a group – Increasing electrons into more energy levels (more shells)

56. Ions Atom, or group of atoms, that has gained or lost electrons Cation – positive ion Anion – negative ion

57. Ions When an atom loses an electron, it becomes positively charged – The radius becomes smaller – Metals tend to lose electrons

58. Ions When an atom gains an electron, it becomes negatively charged – The radius becomes larger – Nonmetals tend to gain electrons

59. Trends Ionic Radius decreases for positive ions across a period – More protons to pull on the electrons Ionic Radius decreases for negative ions across a period – More protons to pull on the electrons

60. Ionic Radius +1 +2 +3 +4 -3 -2

61. Trends Ionic Radius increases down a group – Increasing electrons into more energy levels (more shells)

63. Ionization Energy (IE) Amount of energy required to remove an electron from an atom – Ca  Ca + + e - 590kJ/mol First ionization energy is removing the first electron Second Ionization energy is removing the second electron after having the first removed – Ca +  Ca 2+ + e - 1145kJ/mol

64. IE Trends Ionization energy tends to increase across a period – More protons are able to hold on tighter to electrons Ionization energy tends to decrease down a group – Electrons are farther away from the protons (more shells)

65. Electronegativity (EN) Ability of an atom to attract an electron from another atom when in a compound. – Noble gases are usually omitted since they don’t form compounds – Fluorine, F, is the most electronegative element with a value of 4.0 – Francium, Fr, is the least electronegative element with a value of 0.7

66. EN Trends Electronegativity tends to increase across a period – More protons are able to attract electrons better Electronegativity tends to decrease down a group – Electrons are farther away from the protons (more shells)

67. Trends Summary PropertyPeriod (L  R)Group (T  B) Atomic Number Atomic Mass Atomic Radius Ionic Radius Ionization Energy Electronegativity

68. Reactivity Elements that are more reactive tend to either gain or lose electrons very easily Elements that lose electrons easily have low IE and low EN – Lower left, Fr Elements that gain electrons easily have high IE and high EN – Upper right, F