The Periodic Table

1. The Periodic Table

2. Guiding Questions Why is the periodic table so important? Why is the periodic table shaped the way it's shaped? Why do elements combine? Why do elements react? What other patterns are there in the world and how do they help us?

3. Mendeleev • “Father of Periodic Table” • organized elements based on increasing atomic mass. Found similarities in chemical properties and published his first periodic table in 1869.

4. Moseley • discovered while working with elements that they fit better into pattern when arranged by nuclear charge (number of protons-also known as atomic number).

5. Periodic Law states that physical and chemical properties of the elements are periodic functions of their atomic numbers. • This means that when elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals

6. Organization of the Periodic Table • Periods – rows across, horizontal • Indicates number of energy levels (Principle Quantum Number) • Groups/Families – down, columns • 18 • Indicate number of e- in outer most energy level (1-2, 13-18 main group elements) • Share chemical properities

7. Three Categories of Elements • Metals • Non-Metals • Metalloids

8. Three Categories of Elements • Metals - Groups 1-12 (except H) and under stair-step groups 13-15 - Form ionic and metallic bonds • Luster (shiny, reflects light) • Malleable - can be flattened into sheets • Ductile - can be drawn into thin wires • Good Conductors - heat and electricity can flow throughbecause the outer electrons are not held tightly to the nucleus and move freely - Most are solid at room temperature except Mercury

9. The Periodic Table The Metals are represented in the Periodic Table in blue.

10. Three Categories of Elements • Non-Metals • Dull • Not Malleable/Ductile • Poor Conductors - P block - Form ionic and covalent bonds - Most do not conduct heat or electricity -All, except H, are found on right of periodic table - At room temperature, most are solid or gaseous

11. The Periodic Table The Non-Metals are represented in the Periodic Table in yellow.

12. Three Categories of Elements • Non-Metals Some Examples of Non-Metals • Oxygen (O) • Helium (He) • Sulfur (S) • Chlorine (Cl) • Neon (Ne) • Nitrogen (N)

13. Three Categories of Elements • Metalloids • Have properties of both metals and non-metals. • Generally not shiny • Not malleable and ductile - Form ionic and covalent bonds • They conduct heat and electricity better than nonmetals, but less than metals.

14. Three Categories of Elements • Metalloids Here are the Metalloids • Boron (B) • Arsenic (As) • Tellurium (Te) • Silicon (Si) - Polonium (Po) • Antimony (Sb) • Germanium (Ge)

15. The Periodic Table The Metalloids are represented in the Periodic Table in green.

16. The Periodic Table What do you notice about the way the element groups are arranged in the Periodic Table?

17. Hydrogen • In a class by itself • Most common element in the universe (3/4) • Behaves unlike any other element because it consists of one proton and one electron.

18. Main Group Elements • S and p blocks • Groups 1,2, and 13-18 • Four groups have special names: Alkali metals (Group 1) Alkaline-earth metals (Group 2) Halogens (Group 17) Noble gases (Group 18)

19. Group 1 – Alkali Metals Characteristics • Soft, silver metals • Low melting points and densities • Highly reactive (esp. w/ water) • Do not occur in nature in elemental form • Stored in kerosene or mineral oil • Have one Valence electron

21. Group 2 – Alkaline-earth Metals Characteristics • Gray, metallic solids • Reactive; but, less than alkali • Not found as free elements in nature • Bright fireworks, aircraft • Harder, denser, and stronger than alkali metals • Higher melting points than alkali Have 2 valence electrons

23. Transition Elements • Groups 3-12 (elements in transition) D block • Vary in reactivity and can be found as free elements • Metal characteristics • Form colored compounds • Often occur in nature as uncombined elements • Hg – mercury – liquid metal

24. Valence Electrons for Transition Elements • Group 3: 3 valence electrons • Group 4: 2 to 4 valence electrons • Group 5: 2 to 5 valence electrons • Group 6: 2 to 6 valence electrons • Group 7: 2 to 7 valence electrons • Group 8: 2 or 3 valence electrons • Group 9: 2 or 3 valence electrons • Group 10: 2 or 3 valence electrons • Group 11: 1 or 2 valence electrons • Group 12: 2 valence electrons

26. Group 13 – Boron Group or Icosagens • Have 3 valence electrons • Boron is the only metalloid in the family • The rest are poor metals.

28. Group 14 – Carbon group or Crystallogens • Have 4 valence electrons • Unique feature is that the elements can form different anions and cations Ex: C: 4- Si and Ge: 4+ Sn and Pb: 2+

30. Group 15 – Nitrogen Group or Pnictogens • Have 5 valence electrons • Able to form double and triple bonds

32. Group 16 – Oxygen group or Chalcogens • Have 6 valence electrons • From -2 ions • Physical properities vary dramatically Ex: oxygen – colorless gas Sulfur – yellow solid Tellurium – silver metalloid Selenium - black

34. Group 17 - Halogens • Most reactive non-metal • Irritating odor • Interact with alkali metals to form salts • 7 electrons in outer energy level • Easily pickup one electron • Bromine – only liquid nonmetal • Flourine and chlorine – gases at room temp • Iodine and astatine – solids at room temp

36. Group 18 – Noble Gases • No color or odor • Exist as individual gas atoms (monatomic) • Full outer energy level • Relatively un-reactive (inert gases)

38. Inner Transitional Metals • f-block, n-2 • ALL are radioactive and unstable • Includes lanthanides (atomic number 58-71) • Shiny metals, similar in reactivity to alkaline earth metals • and actinides (atomic number 90-103) • First 4 are found on earth, the remaining are synthetic

42. Diatomic Molecules Elements That Exist as Diatomic Molecules in Their Elemental Forms Element Present Elemental State at 25 oC Molecule hydrogen colorless gas H2 nitrogen colorless gas N2 oxygen pale blue gas O2 fluorine pale yellow gas F2 chlorine pale green gas Cl2 bromine reddish-brown liquid Br2 iodine lustrous, dark purple solid I2

43. How do we use an elements location on the periodic table to determine its ionic charge? • Atomic Radius (Angstrom) • ½ the distance from the nuclei to another • From the nucleus to edge of e- cloud • Going down a group, Atomic radius increases because of the increasing number of energy levels • Going across a period (left to right), atomic radius decreases because of the increase in positive charge in the nucleus

44. Atomic Radii

45. Coulombic attraction • Attraction of + and – charges • Two factors determine the strength: • Amount of charge • Distance between charges

46. Shielding Effect • Kernel electrons “shield” valence electrons from attractive force of the nucleus • Caused by kernel and valence electrons repelling each other • The more electron shells there are, the greater the shielding effect. • Explains why valence electrons are more easily removed

47. Valence Shielding Effect Kernel electrons block the attractive force of the nucleus from the valence electrons + - - nucleus - Electrons - Electron Shield “kernel” electrons

48. Why is cesium bigger than sodium? • Sodium has the electron configuration 1s22s22p63s1. There is one valence electron. • The attraction between this lone valence electron and the nucleus with 11 protons is shielded by the other 10 core electrons. • The electron configuration for cesium is 1s22s22p63s23p64s23d104p65s24d105p66s1. • While there are more protons in a cesium atom, there are also more electrons shielding the outer electron from the nucleus. The outermost electron, 6s1, therefore, is held very loosely. Because of shielding, the nucleus has less control over this 6s1 electron than it does over a 3s1 electron.

49. Ionization Energy • Ionization is a process that results in the formation of an ion • An ion is an atom or group of atoms that have a “+” or “-” charge • Change is created by gain or loss of e- • Losing an e creates a “+” charge (cation)

50. Ionization Energy • Losing the e- requires energy. Energy required to remove one e- from a neutral atom is called IONIZATION ENERGY (also known as 1st ionization energy)