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PERIODICITY

PERIODICITY. SC4b. Compare and contrast trends in the chemical and physical properties of elements and their placement on the Periodic Table. The Periodic Table. Started by Dmitri Mendeleev – arranged by atomic mass Wrote out info known about 63 elements on cards

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PERIODICITY

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  1. PERIODICITY SC4b. Compare and contrast trends in the chemical and physical properties of elements and their placement on the Periodic Table

  2. The Periodic Table • Started by Dmitri Mendeleev – arranged by atomic mass • Wrote out info known about 63 elements on cards • Arranged by similar properties • Predicted existence of 3 unknown elements • w/I 4 years, 2 had been discovered • Since Mendeleev, evidence has supported that the arrangement of elements should be based on the elements atomic number (# of protons)

  3. Periodic Law • Periodic Law – the physical and chemical properties of the elements are periodic functions of their atomic numbers. • Periodic Law allows properties of elements to be predicted based on their position in the periodic table.

  4. The Periodic Table • Period – horizontal row across the PT • Elements in a row have the same number of major energy levels • i.e. A Row 4 element has 4 major energy levels • Group – vertical column in the PT • Elements in the same group have the same # of valence electrons • Valence e-’s are the outermost, bonding electrons (always in the S and P orbitals)

  5. By the NUMBERS Based on the Noble Gases He, Ne, Ar, Kr, Xe, and Rn, is there a repeatable pattern of atomic numbers? A: 2, 8, 8, 18, 18, 32

  6. The Periodic Table • Chemical and Physical properties • Vary across a period • Similar down a group • Elements are either solids or gases (only 2 liquid elements at room temperature, Hg & Br)

  7. METALS • Metals – majority of the PT (78 %) • Properties include high conductivity, high density, solid at room temp., malleable, ductile, and lusterous. • Primarily belong in sublevels S, D, and F.

  8. METALLOIDS • Metalloids – have properties b/w metals and non-metals; small portion (7 %)

  9. NON-METALS • Non-metals – make up right side of PT (15%) • Properties include being non-conductive, low density, mostly gases at room temp., brittle, non-lusterous • 7 non-metals form DIATOMIC (2 atoms) molecules when pure elements

  10. Quantum Numbers • The Principal Quantum Number (n) • Main energy level • Sublevels (s, p, d, f) • The principal quantum number is followed by the sublevel and the number of electrons within the sublevel

  11. Sublevels

  12. Electron Configuration What period are the follow elements in based on their electron configurations? 1) Nitrogen 2) Aluminum 3) Calcium

  13. Valence Electrons Valence Electrons • electrons in the highest energy level of atoms of an element • Each element in a group has same # of valence e-’s • Electrons available to be lost, gained, or shared in bonding

  14. Group Names • Alkali metals – group 1 of the PT • Alkaline Earth metals – group 2 of the PT • Transition Metals – groups 3-12 of the PT • Halogens – group 17 of the PT • Noble gases – group 18 of the PT • Lanthanides – elements 58 – 71 • Actinides – elements 90 - 103

  15. Groups 1 & 2 Alkali Metals Alkali Earth Metals Group 2 Reactive 2 valence electrons ns2 Soft Silvery • Group 1 • Very reactive • 1 valence electron • ns1 • Soft • Silvery

  16. Halogens • Group 17 • Most reactive group of elements • Why? They have 7 valence electrons and want 1 to have a full octet • Based on electronegativity values • Have an electron configuration of ns2 np5 • Fluorine has the greatest Electronegativity

  17. Noble Gases • Group 18 • 8 valence electrons • ns2 np6 • Not reactive • All gases

  18. PERIODICITY SC4a. Use the Periodic Table to predict periodic trends including atomic radii, ionic radii, ionization energy, and electronegativity of various elements.

  19. Periodic Trends • Atomic Radius • Electron Affinity • Ionic Radius • Ionization Energy • Electronegativity

  20. Atomic Radius • Atomic radius  ½ the distance b/w the nuclei of identical atoms that are bonded together • Radii decrease across a period • Radii increase down a group

  21. Electron Affinity • Electron affinity – the amount of energy absorbed when an electron is added to an isolated gaseous atom to form an ion with a 1- charge. Cl(g) + e-  Cl- (g) EA = -349 kJ/mol

  22. Electron Affinity

  23. Ionic Radius • Cation – a positively charged ion (i.e. Na1+) • Anion – a negatively charged ion (i.e. F1-) • Ionic radius increases from right to left across a period and increases from top to bottom down a group.

  24. Ionization Energy • Ionization Energy - energy needed to remove 1 e- from a neutral atom • Increases from bottom to top in a group. • Increases from left to right across a period.

  25. Electronegativity • Electronegativity is a measure of the relative tendency of an atom to attract electrons to itself when it is chemically combined with another atom. • Elements w/ HI Electronegativity’s (nonmetals) gain electrons to form anions. • Elements w/ LOW Electronegativity’s (metals) often lose electrons to form cations.

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