Periodic Table Properties Handout

HANDOUTModule: Tutor: Chemistry SL/HL Dr. Liakatas Topic: Date given: Properties of the periodic table Radius increases Ionization energy decreases Electronegativity decreases Group # → # of valence eRadius increases Ionization energy decreases Electronegativity decreases Period # → # of shells IONS − + F=kq+q-/r2 but same protons in nucleus → same attraction of each valence e.in a covalent bond → smaller electronegativity Negative ion: ► gain of valence e.farther from nucleus → larger atomic radius ► more shells → valence e.by nucleus but larger e-e repulsion → larger ionic radius Positive ion: either ► loss of all valence e.but same protons in nucleus → same attraction of each valence e.farther from nucleus → weaker attraction of an e.by nucleus but smaller e-e repulsion → smaller ionic radius .→ smaller ioniz.in same energy level but less protons in nucleus → weaker attraction of each valence e. energy ► valence e.HANDOUT EXPLANATION OF THE PROPERTIES In same GROUP going down: ► more shells → valence e.in same energy level but less protons in nucleus → weaker attraction of each valence e.in a covalent bond → smaller electronegativity In same PERIOD going left: ► valence e.by nucleus → larger atomic radius ► valence e.farther from nucleus → weaker attraction from nucleus → smaller energy to remove an e. energy ► more shells → valence e.by nucleus → larger atomic radius → weaker attraction of an e.by nucleus → smaller energy to remove an e.→ smaller ioniz.→ one shell less → smaller ionic radius Or ► loss of some valence e.in same energy level but less protons in nucleus → weaker attraction of each valence e. p.p.p. decreases…[4] Period # → # of shells [1]… because valence e.p. no free electrons at all [6]… because for metals.forming free electron cloud trying to keep cations together → weaker forces → smaller m.p. less electrons in valence shell to form free electron cloud but then for non-metals. more valence e. covalent bond weaker than metallic ./b. decreases…[2] Group 7: Reactivity increases…[3] Melting point/b.p./b.HANDOUT Period 3: Electrical conductivity increases and then decreases …[5] Melting point/b.farther from nucleus → easier to be lost to form cation → easier to react with an anion → larger reactivity [2]… because bigger nucleus but same valence e. [5]… because for metals.p. [3]… because larger electronegativity → easier to attract electron and become anion → easier to react with a cation → larger reactivity [4]… because lower molar mass → weaker Van der Waals forces → smaller m.forming free electron cloud keeping cations together but then for non-metals. increases and then decreases…[6] Group # → # of valence eGroup 1: Reactivity increases…[1] Melting point/b. HANDOUT Chemical Properties of Period 3 Chlorides of Period 3 ionic Na compounds Mg Al Si covalent compounds P S Cl NaCl MgCl2 Al2Cl6 (AlCl3) + H2O + H2O + 3 H2O + 4 H2O + 3 H2O + H2O → → → → → ↔ Na+ Mg2+ Al2O3 (base) (acid) (oxide) + Cl+ 2 Cl+ 6 HCl + 4 HCl + 3 HCl + HCl dissociation (conductive when molten) SiCl4 PCl3 Cl2 Si(OH)4 H3PO4 HClO acidic solutions (not conductive) Oxides of Period 3 ionic compounds giant covalent compounds Na Mg Al Na2O MgO Al2O3 Al2O3 SiO2 P4O10 SO3 Cl2O7 + H2O → 2 Na+ + 2 OH→ Mg(OH)2(s) + H2O + 6 HCl → 2 AlCl3 + 3 H2O (base) + 2 NaOH 2 NaAl(OH)4 + 3 H2O (acid) + 2 NaOH → Na2SiO3 + H2O + 6 H2O → 4 H3PO4 + H2O → H2SO4 + H2O → 2 HClO4 Extra ½ O for each successive oxide form bases is amphoteric is weak acid form acids Si P covalent S compounds Cl Dr. Liakatas . Cl) forms a dative (coordinate) bond H2O. NH3. Colored because the d orbitals split and electrons can absorb light and rise to a higher energy orbital Sc & Zn ions are colorless 4. Ni.Liakatas .HANDOUT Properties of d-block elements 1. Formation of complex ions (mainly Co. Multiple oxidation numbers because the 3d and 4s orbitals have similar energy. NH3 → octahedral. Cu) when a ligand (usually H2O. Catalysts: Decomposition of hydrogen peroxide Hydrogenation of alkenes Haber process Contact process 2 H2O2 → 2 H2O + O2 C2H4 + H2 → C2H6 N2 + 3 H2 ↔ 2 NH3 2 SO2 + O2 → 2 SO3 (catalyst = MnO2) (catalyst = Ni) (catalyst = Fe) (catalyst = V2O5) Dr. 2. Cl → tetrahedral 3. O.→ Br2 + 2ClX I Solid Colorless Black/purple Brown AgI → yellow precipitate Cl2 + 2I.p. increasing m. of hydrogen halides (HX) compared to noble gases Hydrogen bond: ► between a H attached to a F. of straight chain isomers Dipole-dipole: ► electrostatic attraction between polar (asymmetric) molecules ► explains: increased m. of halogens. higher m. of water.p.p. N ► explains: high m.p.p. N and a free electron pair of a F. of alcohols and organic acids Strength increases ©IL 6/6 .→ I2 + 2ClBr2 + 2I. O.→ I2 + 2Br- Intermolecular forces Van der Waals: ► attraction of temporary dipoles due to random motion of valence electrons ► in all species. polar and non-polar ► proportional to molar mass and molecule’s surface ► explains: forces between halogen molecules. increased m.HANDOUT F State at room temperature Color of ions Color at normal state Color in solution + Ag + Cl2 + Br2 X X X Gas Colorless Pale yellow Cl Gas Colorless Yellow/green Green AgCl → white precipitate (turns black in sunlight) X X Br Liquid Colorless Red/brown Yellow/orange/brown (depending on concentration) AgBr → creamy precipitate Cl2 + 2Br.