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ORGANIZATION OF THE PERIODIC TABLE

This reading explains the organization of the periodic table, including periods, groups, valence electrons, and electron configuration. It also discusses the different ways the periodic table can be organized and why there are multiple versions.

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ORGANIZATION OF THE PERIODIC TABLE

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  1. ORGANIZATION OF THE PERIODIC TABLE

  2. Questions: Mendeleev Reading • What is a period? • What is a group? • List the ways the Periodic Table is organized:

  3. How is the periodic table arranged? • Periodic table is arranged by increasing atomic number • A new row is added so that elements are arranged with repeating properties in each column • The properties of the elements repeat periodically • Hence… the periodic table! • There are many different ways to organize the elements…

  4. Circular

  5. Long Form

  6. Spiral

  7. Pyramid

  8. Layers

  9. Why so many versions? • Everybody sees their own pattern in the elements • What patterns did we see? • Shells • Outer Electrons

  10. Dmitri Mendeleyev • Father of the Modern P.T.

  11. Periods and Group • Period – horizontal row on P.T. • Each period represents an energy level (think back to models of the atom) • Atoms in period 1 have 1 energy level, atoms in period 5 have 5 energy levels Group – vertical column P.T. Each group represents a certain number of valence electrons Also known as families

  12. Groups and Families • The columns in the periodic table are called groups, or families. • Each group has the same number of outer shell electrons Group 1: Alkali Metals Group 2: Alkaline Earth Metals Groups 3-12: Transition Metals Group 17:Halogens Group 18: Noble Gases

  13. Valence Electrons • Electrons exist within energy levels. • The electrons in the outermost energy level are called valence electrons! • Valence electrons determine how an atom behaves (they perform all bonding)

  14. Valence Electrons Ctd • Each group number represents the number of valence electrons elements in that group have • Ex: group 1 has 1 valence electron • No atom can have more than 8 v.e. • For group numbers 13 - 18 subtract ten to figure the number of valence electrons • Ex: Group 18 = 8 valence electrons • Groups 3-12, we assume to have 2 v.e. (not always the case)

  15. Valence Electrons

  16. 8 is great (sometimes 2) • The noble gases (group 18) are stable atoms, meaning they do not react! (this is good) • The noble gases are stable because they have 8 valence electrons (or 2 as in helium) • All other atoms will gain or lose electrons to become like noble gases (remember cations and anions)

  17. Metals Form Cations • Groups 1-13 are metals (except H) • Metals lose electrons to become like noble gases (less than 4 v.e., lose e-) • Therefore, they form cations (+ charge)!

  18. Non-Metals Form Anions • Groups 15 – 18 are non-metals • They gain electrons to be like noble gases! (more than 4 v.e., gain e-) • Therefore, they form anions (- charge)

  19. Practice • For each of the following elements, determine if they are a metal or non-metal and the charge they would form. Na O F Al Ba I S Cs P

  20. Typical Charge (Ion Formed) • You should remember that 8 is great! • Elements will either gain or lose electrons to try to have 8 • Elements with 3 or less valence electrons will LOSE electrons • Becoming _____________ • Elements with 4 or more valence electrons will GAIN electrons • Becoming _____________ • For Hydrogen and Helium, 2 is good too! • They can’t possibly hold 8 electrons, so they are full at 2

  21. Where are the electrons?-The Bohr Atom • Bohr Model • Based on Line Emission Spectrum of Hydrogen • Atoms consists of nucleus and energy Levels • Stated electrons followed specific circular paths called orbits

  22. Quantum Mechanical Model Consists of Energy levels, sublevels, and orbitals Key Points: • Electrons do not follow orbits, nor can location be known exactly • Electrons are located within orbitals (probable location of electron)

  23. Energy Levels • Each Period number corresponds to a principal energy level • Ex: elements in period 2 have 2 energy levels. • These energy levels are then broken into sublevels, based on which type of orbital found there

  24. Sublevel (also called subshell) • Found within energy levels • Designated by s, p, d, or f • Letter corresponds to orbital shapes found in sublevel “s” sublevel (1 orbital)

  25. The p sublevels (three orbitals) The d sublevel (five orbitals)

  26. The f sublevel (7 orbitals)

  27. Relative size of the 1s, 2s, 3s orbitals

  28. Orbitals • Generalized location of electron • You know I’m probably in this room all day, you just don’t know if I’m at my desk or in the storeroom or walking around • Does not have sharp edges • 1 orbital can contain a maximum of 2 electrons

  29. Electron Configuration • Electron configuration: description of what sublevels and orbitals are filled by electrons in any given atom (like a roadmap of the electrons in an atom) • Determined by the number of electrons the atom has • Governed by 3 rules!

  30. e- configuration rules • Aufbau Principle: an electron occupies the lowest energy level & orbital available • Pauli Exclusion Principle: only two electrons can occupy any orbital, and they must have opposite spins • Hund’s Rule: Each orbital in a given sublevel (s, p, d, or f orbital) must have 1 electron before any can have two

  31. Color in or outline the sections on your blank periodic table to match this diagram

  32. Orbital Diagrams • Using the periodic table from the previous slide, we can start to understand how electron configurations work. (example below) • Ex: Write the Orbital Diagram Magnesium. 1: Determine the atomic number of the element from the Periodic Table • This gives the number of protons and electrons in a neutral atom of that element Mg, Z = 12, so Mg has 12 protons and 12 electrons

  33. 1s 2s 2p 3s 3p Orbital Diagrams ctd • 2. Draw 9 boxes to represent the first 3 energy levels s and p orbitals

  34. 1s 2s 2p 3s 3p Orbital Diagrams ctd 3. Add one electron to each box in a set, then pair the electrons before going to the next set until you use all the electrons • When pairing, put in opposite arrows

  35. Orbital Diagram Practice • Write orbital diagrams for each of the following elements: • Ne • Be • Al

  36. standard e- configuration • Similar to orbital diagram, yet we don’t draw in arrows. • Instead we write the number of electrons in each sublevel as a exponent • Ex: Oxygen 1s2 2s2 2p4 (standard configuration)

  37. Electron Configurations how many electrons in that orbital • Nitrogen: 1s22s22p3 energy level orbital (atomic number = 7) Tro's Introductory Chemistry, Chapter 9

  38. Standard e- configuration practice • For each of the following, write out the standard electron configuration: • Mg • Ne • S • Fe

  39. Compare these 2 e- configurations • Ar (#18) 1s2 2s2 2p6 3s2 3p6 • Ca (#20) 1s2 2s2 2p6 3s2 3p6 4s2 • They are exactly the same until the outer most sub level

  40. Noble Gas (abbreviated) configuration • We use the elements in group 18, the ones with a full octet, the noble gases to make a short cut. (remember they do not gain or lose e-) • [Ar] 4s2 tells us that Ca has the same e- configuration as Ar, except for the valence electrons, which are in 4s2.

  41. Noble Gas (abbreviated) practice • For each of the following, list the Noble Gas configuration • Li • Pb • Fe • I • Kr

  42. Exceptions to Rules • Groups 6 & 11 do not follow all rules exactly! • They steal one electron from the previous “s” sublevel to stabilize their unfilled “d” orbital • Ex:

  43. Charges & E- Config • Determine number of electrons and write e-config to go along with number of e- • Metals form Cations (lose electrons) • Na+1 • Ca+2 • Non-Metals form Anions (gain electrons) • F-1 • S-2

  44. Quiz next class (Fri 11/16) • Ions (how to determine charge) • Three rules governing E config • Aufbau • Pauli • Hunds • Three types of e config • Orbital • Standard • Abbreviated (noble gas) • E config for Ions/Exceptions

  45. Flame Test Pre-Lab (due next class) • Cover Page • Date/Name/Lab Name • Introduction • Discuss topics covered in class (energy levels) • Northern Lights/Neon Signs • 2 different citations/sources • Establish objectives/purpose of the lab • Data Table

  46. Flame Test – Final Report (due 11/20) • Title Page • Intro • Data Table • Post Lab Questions (just 1-5) • Type Reponses in complete sentences • Conclusion • What was learned (examples of flame colors for ions)? • 2 Possible Errors and Impact on observations • Identification of unknown with references to data table to support

  47. Quiz Practice For each atom below: • Determine # of Valence e- • Determine if metal or non-metal • Determine charge it forms • Al • Br • S • P When done, pg 107 #24, 26, 27

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