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The Periodic Table

The Periodic Table. The Periodic Table. Understand the rationale behind the periodic table; view the table as an ordered database of element properties. Explain how the periodic table reflects the quantum mechanical structure of the atom.

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The Periodic Table

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  1. The Periodic Table

  2. The Periodic Table • Understand the rationale behind the periodic table; • view the table as an ordered database of element • properties. • Explain how the periodic table reflects the quantum • mechanical structure of the atom. • Explain and use periodic trends in: • atomic radius • ionic radius • ionization energy • Explain the connection between ionization energy • and metallic character.

  3. Quick Review • Quantum Numbers (4) • Represent or describe the space orbital in which the e- moves or may be found. • Distance from nucleus • Shape of the orbital • Positionor location with respect to the 3 axes in space • Direction of spin of the e- in the orbital

  4. Quantum Numbers 1st Quantum Number: Principle Quantum # (n) – avg. distance of e- from nucleus. Positive whole # 1,2,3, etc. The main energy level designation equal to K,L,M, etc. shells

  5. Quantum # 2 Orbital Quantum # (l) – shape of orbital in which the e- moves. # of possible shapes = the value of the principal quantum #. 1st energy level = 1 shape 2nd energy level = 2 shapes 3rd energy level = 3 shapes, etc Letter designations for shapes are s,p,d,f. These are also sublevels or sub orbitals S – Sphere, P – 8 ( 3 - one on each axis x,y,z)

  6. 3rd Quantum # - Magnetic (m) – indicates the position about the three axes in space of the orbital. 1 position for s orbital 3 positions for p orbitals (pxpypz) 5 positions for d orbitals 7 positions for f orbitals

  7. 4th Quantum Number Spin (s) – right hand or left hand, clockwise or counter-clockwise Therefore 2 e- can occupy any space designated by the first 3 quantum #’s but they will have opposite spins. No 2 e- have the same quantum #’s No 2 e- have the same energy

  8. In 1859 Mendeleyev devised the periodic table of elements. • “Properties of the elements are in periodic dependence on their atomic weights.” • Ramsay discovered the noble gases: Ne, Ar, Kr, Xe. Noble or inert: will not react with others. • Periodic Law: The physical & chemical properties of the elements are periodic functions of their atomic #’s. 1. Period/Series – horizontal rows 2. Group/Family – vertical columns a. elements with similar properties have a similar arrangement of outer-shell e- (same group in the periodic table) or valence e-.

  9. Most active metal – lower left • Most active non-metal – upper right • Transition elements are metallic & have either 1 or 2 e- in their outer shell (generally) • Rare Earth elements are basically identical. The outer shell contains 2 e- • Zigzag line: separates metals from the non-metals. Those touching are metalloids, have both or either characteristics of metals and non-metals.

  10. Atomic Radii The radius of an atom does not increase with atomic #. 1. Atomic radius increases with atomic # in a particular group/family. (increased shells) 2. From group 1 to 18 (I to VIII) in a period, there is a general decrease in the atom radii. (increase in p+, increase attraction of the e-)

  11. Atomic radius

  12. Ionization Energy • Energy required to remove an e- from an atom. • normally, ionization removes valence electrons first • valence electrons are farthest from nucleus on average, so they feel the least attraction for the nucleus and are easiest to remove • end of valence electrons is marked by a big jump in ionization energies

  13. Ao (nuetral) + Energy = A + (positive ion) + e- • kJ – kilo joules or kcal – kilo calories may be the energy unit of measure used per mole of substance. • Ex: It takes 124 kcal to remove the outer shell e- from 6.02 X 1023 atoms of Li. • Mole: amount of the substance equal to a Avogadro Number: 6.02 X 1023atoms • Mole: may also equal that substances atomic weight in grams.

  14. Quick Check for understanding • normally, ionization removes valence electrons first • valence electrons are farthest from nucleus on average, so they feel the least attraction for the nucleus and are easiest to remove • end of valence electrons is marked by a big jump in ionization energies • Low ionization energy is characteristic of a metal. • High ionization energy is characteristic of a non-metal • Intermediate ionization energy for metalloids

  15. Within a group/family ionization energy generally decreases with increasing atomic # • Ionization energy does not vary uniformly within a period/series as a result of filled or unfilled sublevels. • To remove successive e- from the same atom becomes progressively more difficult due to the closeness of the remaining e- to the nucleus. • Na + Energy = Na+ + e- +496 kJ 1st ionization energy • Na + Energy = Na++ + 2e- • +4560 kJ 2nd ionization energy

  16. Electron Affinity • Energy released when an e- is added to a neutral atom. • A + e- = A- + energy (kcal) Note: E on Product side of the equation • Low electron affinity – weak bonding • High electron affinity – strong bonding • Electron affinities decrease in groups/families due to additional shells • Electronegativity: ability to attract e-

  17. Additional Terms • Ion – an atom which has become charged by either gaining or losing electrons. • Anion – negatively charged atom, characteristic of the non-metals. • Cation – positively charged atom, characteristic of the metals.

  18. Rare EarthElements Group or Family Names: 1 Alkali metals, 2 Alkaline earth metals, 3-12 transition metals, 13 Boron-Aluminum, 14 Carbon, 15 Nitrogen, 16 Chalcogen, 17 Halogens (salt formers), 18 Noble (Inert) or Group O

  19. Electron donor, lender Ionic Bonding Electron acceptor, borrower Covalent Bonding

  20. CHEMICAL FAMILIES: ALKALI METALS: GROUP IA (1) All alkali metals react with H2O to form an alkaline (Basic) Solution. ** Hydrogen is NOT considered an alkaline metal in this case. ALKALINE EARTH METALS: GROUP IIA (2) All alkali earth metals also react with H2O to form a basic solution. HALOGENS: GROUP VIIA (17) Salt Formers NOBLE GASES: GROUP VIIIA (18) Noble Gases have all their orbitals Filled. Rare Earth Elements – Inner Transition Elements Lanthanide Series: Fourteen elements beginning with lanthanum in which the highest energy electrons to be in the 4f sublevel. Actinide Series: Fourteen elements beginning with actinium in which the highest energy electrons to be in the 5f sublevel. The other groups are identified by the element at the top of the column. Example, GROUP IVA is the called the Carbon Family (or Group).

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