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CHEM1612 - Pharmacy Week 7: Oxidation Numbers

CHEM1612 - Pharmacy Week 7: Oxidation Numbers. Dr. Siegbert Schmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au. Unless otherwise stated, all images in this file have been reproduced from:

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CHEM1612 - Pharmacy Week 7: Oxidation Numbers

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  1. CHEM1612 - PharmacyWeek 7: Oxidation Numbers Dr. SiegbertSchmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au

  2. Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille,Chemistry, John Wiley & Sons Australia, Ltd. 2008      ISBN: 9 78047081 0866

  3. Oxidation numbers Textbook: Blackman, Bottle, Schmid, Mocerino & Wille, “Chemistry”,John Wiley & Sons Australia, Ltd., 2008. Today’s lecture is in • Section 4.6, 4.8 • Section 12.1 • Section 13.1, 13.2 Potassium atom, K 19 protons, 19 neutrons 19 electrons

  4. Oxidation numbers: definition • Each atom in a molecule is assigned an OXIDATION NUMBER (O.N.). • The oxidation number is the charge the atom would have if the electrons in a bond were not shared but transferred completely to the more electronegative atom. Electrons shared equally as both Cl atoms in Cl2 have the same electronegativity. Oxidation number = 0. Unequal sharing of electrons, F has higher electronegativity than H. Therefore oxidation number of H will be positive (+I), and F will be negative (-I).

  5. Oxidation numbers (states) • USE OF OXIDATION NUMBERS • Naming compounds • Properties of compounds • Identifying redox reactions • In a binary ionic compound O.N.= its ionic charge. • In a covalent compound O.N. ≠ a charge. • O.N. is written as • a roman numeral (I, II, III, etc.) • a number preceded by the sign (+2) • Ionic charge has the sign after the number (2+). Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

  6. Electronegativity • Definition: Ability of a bonded atom to attract the shared electrons. (Different from electron affinity, which refers to the ability of an isolated atom in the gas phase to gain an electron and form a gaseous anion). • Electronegativity is inversely related to atomic size. • Atomic size: increases down group (electrons in outer shells) decreases across period (electrons in same shell) • Electronegativity is directly related to ionization energy (energy required to remove an electron from atom).

  7. Electronegativity and the Periodic Table • Linus Pauling defined electronegativity in arbitrary units 0.7 to 4.0 • smallest at lower left Periodic Table - Cs cesium • greatest at upper right - F fluorine Blackman Figure 5.5

  8. Rules for assigning O.N. • The oxidation number for any free element (eg. K, Al, O in O2) is zero. • The oxidation number for a simple, monatomic ion is equal to the charge on that ion (eg. Na+ has oxidation number +I) • The sum of all the oxidation numbers of the atoms in a neutral compound must equal zero (e.g. NaCl). The sum of all the oxidation numbers of all the atoms in a polyatomic ion must equal the charge on that ion (e.g. SO42-). • In all its compounds fluorine has oxidation number –I. • In most of its compounds hydrogen has oxidation number +I. • In most of its compounds oxygen has oxidation number -II. Blackman pg. 464

  9. H H-C-H H Oxidation numbers • Molecules and polyatomic ions: shared electrons are assigned to the more electronegative atom. • Examples: HF F-I H I CO2 O-II C+IV O=C=O CH4 H +I C-IV NO3- -1 charge on anion = 3 x O-II + NV • Determining an atom’s oxidation number: • The more electronegative atom in a bond is assigned all the shared electrons; the less electronegative atom is assigned none. • Each atom in a bond is assigned all of its unshared electrons. • The oxidation number is give by: • O.N. = no. of valence e- - (no. of shared e- + no. of unshared e-) For F, O.N. = 7 – (2 + 6) = -1

  10. Pop Quiz What is the oxidation number of Cr in the following? [Cr2O7]2 2(x) + 7(-2) = -2, x = +6, Cr(VI) CrO3 x + 3(-2) = 0, x = +6, Cr(VI) Cr2O3 2(x) + 3(-2) = 0, x = +3, Cr(III)

  11. Pop Quiz • Examples I2 O.N.=0 (elemental form) Zn in ZnCl2 O.N.=+2 (Cl=-1, sum of O.N.s =0) Al3+ O.N.=+3 (ON of monatomic ion=charge) N in HNO3 O.N.=+5 (O=-2, H=+1, sum of ONs=0) S in SO42- O.N.=+6 (O=-2, sum of O.N.s=charge on ion) N in NH3 O.N.= -3 (H=+1, sum of O.N.s = 0) N in NH4+ O.N.= -3 (H=+1, sum of O.N.s =charge on ion)

  12. Demo: Oxidation states of V • Zn (s) + 2 VO3-(aq) + 8H+ (aq) → 2VO2+ (aq) + Zn2+ (aq) + 4 H2O +5, vanadate, yellow +4, vanadyl, green • Zn (s) + 2 VO2+ (aq) + 4 H+ → 2 V3+(aq) + Zn2+(aq) + 2 H2O +4, vanadyl, green +3, blue • Zn (s) + 2 V3+(aq) → 2 V2+ (aq) + Zn2+ (aq) blue +2, violet

  13. Transition Metals Multiple oxidation numbers – ns and (n-1)d electrons are used for bonds.

  14. Transition Metals Multiple oxidation numbers – ns and (n-1)d electrons are used for bonds.

  15. Filling of Atomic Orbitals (Aufbau) In general, the (n-1)d orbitals are filled between the ns and np orbitals. Blackman Figure 4.29

  16. Transition Metals – Ion Formation • Period 4 Transition Metals: as the d orbitals fill, the 3d orbital becomes more stable than the 4s. • In the formation of Period 4 transition metal ions, the 4s electrons are lost before the 3d electrons. • The 4s orbital and the 3d orbitals have very similar energies  variable oxidation states.

  17. 3d electrons Common O.N. +III +IV +V +VI +VII +III +III +II +II +II +IV +III +IV +II +II +II +II

  18. Mn = [Ar]4s23d5 7 valence electrons Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. Orbital Occupancy

  19. Influence of Oxidation State Hexavalent Chromium • Cr(VI) is classified as “carcinogenic to humans” • Cr(VI) compounds are soluble in water & may have a harmful effect on the environment. • Cr(VI) is readily reduced by Fe2+ and dissolved sulfides. Trivalent Chromium • Cr(III) is considered an essential nutrient. • Most naturally occuring Cr(III) compounds are insoluble and it is generally believed that Cr(III) does not constitute a danger to health. • Cr(III) is rapidly oxidised by excess MnO2, or slowly by O2 in alkaline solutions.

  20. Properties of N-compounds • Some non-metals like sulphur or nitrogen or chlorine also have a very wide range of oxidation states in their compounds. • N-compounds have a very wide range of properties. • N has an intermediate electronegativity and has an odd number (5) of valence electrons. N has one of the widest ranges of common oxidation states of any element.

  21. Oxidation states of N

  22. Properties of N-compounds HIGHLY VARIED! Incredibly stable: N2 Extremely explosive: trinitrotoluene (TNT) Strong acid HNO3 Weak base NH3 Photochemical smog: NO2 Biologically important: NO + amino acids nitroglycerine

  23. Nitrogen Oxides Table from Silberberg, “Chemistry”, McGraw Hill, 2006.

  24. Air pollution Picture from www.consumercide.com Picture from http://pdphoto.org Sydney The brown haze is largely NO2  Los Angeles

  25. Summary • Rules for assigning oxidation numbers • Trends in electronegativity • Electron configuration of elements and ions • Aufbau – rule for filling atomic orbitals • Electron configuration of transitions metals

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