The Periodic Table

1. The Periodic Table

2. Pre-Periodic Table Chemistry … • …was a mess!!! • No organization of elements. • Imagine going to a grocery store with no organization!! • Difficult to find information. • Chemistry didn’t make sense.

3. Dmitri Mendeleev (1834-1907) • Presented the first valid periodic table of elements to the Academy of Sciences in 1869 • Grouped elements in ascending order by atomic mass and by similarities in properties • left gaps for elements not yet discovered • Element Mendelevium (element 101) was named to honor him

4. Mendeleev's Periodic Table(1871) Dmitri Mendeleev: Father of the Periodic Table HOW HIS WORKED… • Put elements in rows by increasing atomic weight. • Put elements in columns by the way they reacted. SOME PROBLEMS… • He left blank spaces for what he said were undiscovered elements. (Turned out he was right!) • He broke the pattern of increasing atomic weight to keep similar reacting elements together.

5. The Current Periodic Table • Mendeleev wasn’t too far off. • Now elements are put in rows by increasing ATOMIC NUMBER!! • The horizontal rows are called periods and are labeled from 1 to 7. • The vertical columns called groups are chemically similar and labeled from 1 to 18.

6. Across the Periodic Table • Periods: Are arranged horizontally across the periodic table • (rows 1-7) • These elements have the same number of valence shells (outer energy level). 2nd Period 6th Period

7. Quick quiz • 1. Name the two properties that Mendeleev used to arrange his Periodic Table? • 2. When was Mendeleev working on his Periodic Table? • 3. How did the answer of #1 change as we moved to the Modern Periodic Table?

8. Periodic TableElectron configuration from the Periodic Table • B • 2p1 H 1s1 He 1s2 F 2p5 Be 2s2 B 2p1 C 2p2 N 2p3 Ne 2p6 O 2p4 Li 2s1 Na 3s1 Mg 3s2 Cl 3p5 Si 3p2 S 3p4 Ar 3p6 Al 3p1 P 3p3 K 4s1 Ca 4s2 Zn 3d10 As 4p3 Be 4p5 V 3d3 Mn 3d5 Fe 3d6 Co 3d7 Sc 3d1 Ti 3d2 Ga 4p1 Ge 4p2 Se 4p4 Cr 4s13d5 Kr 4p6 Ni 3d8 Cu 4s13d10 Sr 5s2 Rb 5s1 Nb 4d3 Ru 4d6 Rh 4d7 Mo 5s14d5 Cd 4d10 Sn 5p2 I 5p5 Xe 5p6 Zr 4d2 Tc 4d5 Y 4d1 In 5p1 Sb 5p3 Te 5p4 Ni 4d8 Ag 5s14d10 Hf 5d2 Cs 6s1 Ta 5d3 Re 5d5 Os 5d6 Ir 5d7 W 6s15d5 La 5d1 Rn 6p6 At 6p5 Ni 5d8 Ba 6s2 Hg 5d10 Tl 6p1 Pb 6p2 Bi 6p3 Po 6p4 Au 6s15d10 Mt 6d7 Bh 6d5 Hs 6d6 Fr 7s1 Rf 6d2 Ra 7s2 Db 6d3 Sg 7s16d5 Ac 6d1

9. Valence Electrons • The valence electrons are the electrons in the last shell or outer energy level of an atom. They show a repeating or periodic pattern. The valence electrons increase in number as you go across a period.

10. Groups…Here’s Where the Periodic Table Gets Useful!! • Elements in the same group have similar chemical and physical properties!! • (Mendeleev did that on purpose.) • Why?? • They have the same number of valence electrons. • They will form the same kinds of ions and bonds.

11. Families on the Periodic Table • Columns are also grouped into families. • Families may be one column, or several columns put together. • Families have names rather than numbers. (Just like your family has a common last name.)

12. Alkali Family: 1 e- in the valence shell Halogen Family: 7 e- in the valence shell Down the Periodic Table • Family:Are arranged vertically down the periodic table • (columns or group, 1- 18 or 1-8 A,B) • These elements have the same number electrons in the outer most shells, the valence shell.

13. Halogen Alkali Noble Gas Alkaline (earth) Transition Metals Infamous Families of the Periodic Table • Notable families of the Periodic Table

14. Metals vs Non-Metals • Metals are found to the left of the staircase • Lose e-’s • Good conductors of heat and electricity • Malleable and ductile • Shiny & hard • High MP/BP • Non-metals are right of the staircase • Gain e-’s • Poor conductors of heat and electricity • Brittle • Not shiny & softer • Low MP/BP

15. Periodic Table: Metallic arrangement • Layout of the Periodic Table: Metals vs. nonmetals Nonmetals Metals

16. Reading the Periodic Table: Classification • Nonmetals, Metals, Metalloids, Noble gases

17. Hydrogen • Hydrogen belongs to a family of its own. • Hydrogen is a diatomic, reactive gas. • Hydrogen was involved in the explosion of the Hindenburg in 1937. • Hydrogen is promising as an alternative fuel source for the future • 1 valence e- • http://www.bing.com/videos/search?q=hindenburg+disaster&qpvt=hindenburg+disaster&FORM=VDRE#view=detail&mid=41F914E72CBA1C33237741F914E72CBA1C332377

18. Alkali Metals • 1st column on the periodic table (Group 1) not including Hydrogen. • Very reactive metals, always combined with something else in nature (like in salt, NaCl). Never found as individual atoms. • Metals explode in water http://www.sciencekids.co.nz/videos/experiments/metalexplosion.html • Soft enough to cut with a butter knife. Low melting point and boiling point • 1 valence e-

19. Quick Quiz • 1. Why is Hydrogen not included in the Alkali Family? • Although hydrogen has 1 valence electron, it is a non-metal • 2. Are Alkali Metals known for their lack of chemical activity? • No, alkali metals are highly reactive • 3. Are Alkali Metals stored underwater so they cannot come in contact with the Oxygen in the air and begin to oxidize? • No, alkali metals explode in water! • 4. If Hydrogen becomes a substitute for petroleum fuels, will cars become much less safe than today’s vehicles? • No, hydrogen fuel does not produce hot ash or radiant heat as gasoline, and hydrogen leaks disperse rapidly in the atmosphere resulting in less time to burn; Issues with tank material, strength, and weight

20. Alkaline Earth Metals • Second column on the periodic table. (Group 2) • Reactive metals that are always combined with nonmetals in nature. • Harder, denser, and stronger than alkali metals • Several of these elements are important mineral nutrients (such as Mg in bones and Ca in teeth and bones. • 2 valence e-’s

21. Transition Metals • Elements in groups 3-12 • Less reactive harder metals • Gold and platinum among the least reactive of all elements • High luster; strong but deformable • Good electric and heat conductors • Includes metals used in jewelry and construction • Metals used “as metal” • Similar properties due to same number of electrons in outer energy level

22. Boron Family • Elements in group 13 • Aluminum metal was once rare and expensive as Gold. Not a “disposable metal.” • Has now been discovered in a very common ore called Bauxite • Family has 3 valence e-’s

23. Rare Earth Elements • Lanthanide Series • Atomic Numbers 57-71 • Not periodic • Unpredictable • Shiny metals • Similar reactivity to Group 2 Alkaline Earth Metals • Actinide Series • Atomic Numbers 89-103 • Uranium is heaviest natural element • Heavier elements are man-made ONLY • All element above 92 are Radioactive

24. Carbon Family • Elements in group 14 • Contains CARBON which is important to all living things • Carbon is the basis for Organic Chemistry. • Silicon and Germanium are important in computers as semiconductors. • Form the most, strongest bonds • 4 valence e-’s

25. Nitrogen Family • Elements in group 15 • Nitrogen makes up 78% of the atmosphere. • Nitrogen and phosphorus are both important in living things (proteins and amino acids). • Most of the world’s nitrogen is not available to living things. • Nitrogen is found in explosives (TNT) and fertilizers • The red stuff on match heads is Phosphorus • 5 valence e-

26. Oxygen Family • Elements in group 16 • Oxygen is necessary for respiration and combustion. • Many things that stink, contain sulfur (rotten eggs, garlic, skunks,etc.) • 6 valence e-

27. Halogens • Elements in group 17 F, Cl, Br, I • Very reactive, volatile, diatomic, nonmetals • Very TOXIC • Always found combined with other elements in nature . • Used to kill bacteria and to strengthen teeth. • 7 valence e-’s

28. Quick Quiz • 1. What element is most often associated with explosives? • Nitrogen • 2. What element is frequently associated with bad smelling products? • Sulfur • 3. If your toothpaste has Sodium Fluoride, what is it doing for you? • Kill bacteria and strengthen teeth • 4. What properties of Carbon cause it to be the “Basis of Life”? • Four valence electrons for bonding; can form long chains of stable molecules; single double and triple bonds

29. The Noble Gases Helium, Neon, Argon, Krypton, Xenon, Radon and Ununoctium

30. The Noble Gases • Elements in group 18 • VERY UNREACTIVE monatomic gases • Used in lighted “neon” signs • Helium is used to fix the Hindenburg problem and balloons • Have a full valence shell with usually 8 valence e-’s.

31. Metalloids • B, Si, Ge, As, Sb, Te • Semiconductors • Mostly brittle solids • Some properties of metals and some of non-metals

32. Summary • Periodic Table: Map of the Building block of matter • Type: Metal, Metalloid and Nonmetal • Groupings: Representative or main, Transition and Lanthanides/Actanides • Family: Elements in the same column have similar chemical property because of similar valence electrons • Alkali, Alkaline, Halogens, Noble gases • Period: Elements in the same row have valence electrons in the same shell (energy level)

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

34. What do they influence? • Energy levels and Shielding have an effect on elements within a GROUP • Nuclear charge has an effect on elements within a PERIOD

35. Atomic Size } • Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule. Radius

36. Atomic Radius - Group trends H Li • As we increase the atomic number (or go down a group). . . • each atom has another energy level, • so the atoms get bigger. Na K Rb

37. Atomic Radius - Period Trends • Going from left to right across a period, the size getssmaller. • Electrons are in the same energy level. • But, there is more nuclear charge. • Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

38. Trend in Atomic Radius

40. Ions • Some compounds are composed of particles called “ions” • An ion is an atom (or group of atoms) that has a positive or negative charge • Atoms are neutral because the number of protons equals electrons • Positive and negative ions are formed when electrons are transferred (lost or gained) between atoms

41. Ions • Metals tend to LOSE electrons, from their outer energy level • Sodium loses one: there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” • The charge is written as a number followed by a plus sign: Na1+ • Now named a “sodium ion”

42. Ions • Nonmetals tend to GAIN one or more electrons • Chlorine will gain one electron • Protons (17) no longer equals the electrons (18), so a charge of -1 • Cl1- is re-named a “chloride ion” • Negative ions are called “anions”

43. Trends in Ionic Size: Cations • Cations form by losing electrons. • Cations are smaller than the atom they came from – not only do they lose electrons, they lose an entire energy level. • Metals form cations. • Cations of representative elements have the noble gas configuration before them.

44. Ionic size: Anions • Anions form by gaining electrons. • Anions are bigger than the atom they came from – have the same energy level, but a greater area the nuclear charge needs to cover • Nonmetals form anions. • Anions of representative elements have the noble gas configuration after them.

45. Ion Group trends • Each step down a group is adding an energy level • Ions therefore get bigger as you go down, because of the additional energy level. Li1+ Na1+ K1+ Rb1+ Cs1+

46. Ion Period Trends • Across the period from left to right, the nuclear charge increases - so they get smaller. • Notice the energy level changes between anions and cations. N3- O2- F1- B3+ Li1+ Be2+ C4+

47. Configuration of Ions • Ions always have noble gas configurations ( = a full outer level) • Na atom is: 1s22s22p63s1 • Forms a 1+ sodium ion: 1s22s22p6 • Same configuration as neon. • Metals form ions with the configuration of the noble gas before them - they lose electrons.

48. Configuration of Ions • Non-metals form ions by gaining electrons to achieve noble gas configuration. • They end up with the configuration of the noble gas after them.

49. Trends in Ionization Energy • Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom). • Removing one electron makes a 1+ ion. • The energy required to remove only the first electron is called the first ionization energy.

50. Ionization Energy • The second ionization energy is the energy required to remove the second electron. • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.