1. CHEM1612 - Pharmacy Week 7: Oxidation Numbers Dr. Siegbert Schmid 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. Lecture 21 -3 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 Oxidation numbers Potassium atom, K 19 protons, 19 neutrons 19 electrons

4. Lecture 21 -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 Cl 2 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. Lecture 21 -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. Lecture 21 -6 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). Electronegativity

7. Lecture 21 -7

8. Lecture 21 -8 Electronegativity and the Periodic Table Blackman Figure 5.5 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

9. Lecture 21 -9 Rules for assigning O.N. 1.The oxidation number for any free element (eg. K, Al, O in O 2 ) is zero. 2.The oxidation number for a simple, monatomic ion is equal to the charge on that ion (eg. Na + has oxidation number + I ) 3.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. SO 4 2- ). 4.In all its compounds fluorine has oxidation number – I. 5.In most of its compounds hydrogen has oxidation number + I. 6.In most of its compounds oxygen has oxidation number - II. Blackman pg. 464

10. Lecture 21 -10 Molecules and polyatomic ions: shared electrons are assigned to the more electronegative atom. Examples: HF F -I H I CO 2 O -II C +IV O=C=O CH 4 H +I C -IV NO 3 - -1 charge on anion = 3 x O -II + N V Determining an atom’s oxidation number: 1. The more electronegative atom in a bond is assigned all the shared electrons; the less electronegative atom is assigned none. 2. Each atom in a bond is assigned all of its unshared electrons. 3. The oxidation number is give by: 4. O.N. = no. of valence e - - (no. of shared e - + no. of unshared e - ) For F, O.N. = 7 – (2 + 6) = -1 Oxidation numbers H H-C-H H

11. Lecture 21 -11 [Cr 2 O 7 ] 2  2(x) + 7(-2) = -2, x = +6, Cr(VI) What is the oxidation number of Cr in the following? CrO 3 x + 3(-2) = 0, x = +6, Cr(VI) Cr 2 O 3 2(x) + 3(-2) = 0, x = +3, Cr(III) Pop Quiz

12. Lecture 21 -12 Examples I 2 O.N.=0 (elemental form) Zn in ZnCl 2 O.N.=+2 (Cl=-1, sum of O.N.s =0) Al 3+ O.N.=+3 (ON of monatomic ion=charge) N in HNO 3 O.N.=+5 (O=-2, H=+1, sum of ONs=0) S in SO 4 2- O.N.=+6 (O=-2, sum of O.N.s=charge on ion) N in NH 3 O.N.= -3 (H=+1, sum of O.N.s = 0) N in NH 4 + O.N.= -3 (H=+1, sum of O.N.s =charge on ion) Pop Quiz

13. Lecture 21 -13 Demo: Oxidation states of V Zn (s) + 2 VO 3 - (aq) + 8H + (aq) → 2VO 2+ (aq) + Zn 2+ (aq) + 4 H 2 O +5, vanadate, yellow +4, vanadyl, green Zn (s) + 2 VO 2+ (aq) + 4 H + → 2 V 3+ (aq) + Zn 2+ (aq) + 2 H 2 O +4, vanadyl, green +3, blue Zn (s) + 2 V 3+ (aq) → 2 V 2+ (aq) + Zn 2+ (aq) blue+2, violet

14. Lecture 21 -14 Multiple oxidation numbers – ns and (n-1)d electrons are used for bonds. Transition Metals

15. Lecture 21 -15 Multiple oxidation numbers – ns and (n-1)d electrons are used for bonds. Transition Metals

16. Lecture 21 -16 Filling of Atomic Orbitals (Aufbau) Blackman Figure 4.29 In general, the (n-1)d orbitals are filled between the ns and np orbitals.

17. Lecture 21 -17 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.

18. Lecture 21 -18 3d electrons Common O.N. +III +IV +V +VI +VII +III +III +II +II +II +IV +III +IV +II +II +II +II

19. Lecture 21 -19 Mn = [Ar]4s 2 3d 5 7 valence electrons Orbital Occupancy Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

20. Lecture 21 -20 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 Fe 2+ 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 MnO 2, or slowly by O 2 in alkaline solutions. Influence of Oxidation State

21. Lecture 21 -21 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.

22. Lecture 21 -22 Oxidation states of N NVNV HNO 3 / NO 3 - Strong acid N IV NO 2, N 2 O 4 Smog N III HNO 2 / NO 2 - Weak acid / weak base N II NO Smog + biology NINI N2ON2O Greenhouse gas + laughing gas N0N0 N2N2 Stable N -I NH 2 OHHydroxylamine N -II N2H4N2H4 Hydrazine, rocket fuel N -III NH 3 / NH 4 + Weak base / weak acid

23. Lecture 21 -23 Properties of N-compounds HIGHLY VARIED! Incredibly stable: N 2 Extremely explosive: nitroglycerine trinitrotoluene (TNT) Strong acid HNO 3 Weak base NH 3 Photochemical smog:NO 2 Biologically important:NO + amino acids

24. Lecture 21 -24 Nitrogen Oxides Table from Silberberg, “Chemistry”, McGraw Hill, 2006.

25. Lecture 21 -25 Air pollution  Los Angeles Sydney The brown haze is largely NO 2 Picture from www.consumercide.com Picture from http://pdphoto.org

26. Lecture 21 -26 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