chem1612 - pharmacy week 7: oxidation numbers

CHEM1612 - PharmacyWeek 7: Oxidation NumbersDr. Siegbert Schmid

School of Chemistry, Rm 223Phone: 9351 4196E-mail: siegbert.schmid@sydney.edu.au

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

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, K19 protons, 19 neutrons19 electrons

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 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).

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 S

ilberberg, “Chem

istry”, McG

raw H

ill, 2006.

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

Lecture 21 -7

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

Lecture 21 -9

Rules for assigning O.N.1. The oxidation number for any free element (eg. K, Al, O in O2) 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. SO4

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

Lecture 21 -10

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:

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

H F

Oxidation numbers

HH-C-H

H

Lecture 21 -11

[Cr2O7]2

2(x) + 7(-2) = -2, x = +6, Cr(VI)

What is the oxidation number of Cr in the following?

CrO3

x + 3(-2) = 0, x = +6, Cr(VI)

Cr2O3

2(x) + 3(-2) = 0, x = +3, Cr(III)

Pop Quiz

Lecture 21 -12

ExamplesI2 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)

Pop Quiz

Lecture 21 -13

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

Lecture 21 -14

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

Transition Metals

Lecture 21 -15

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

Transition Metals

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.

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.

Lecture 21 -18

3d electrons

CommonO.N.

+III +IV +V +VI +VII +III +III +II +II +II +IV +III +IV +II +II +II +II

Lecture 21 -19

Mn = [Ar]4s23d5

7 valence electrons

Orbital Occupancy

Figure from S

ilberberg, “Chem

istry”, McG

raw H

ill, 2006.

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 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.

Influence of Oxidation State

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.

Lecture 21 -22

Oxidation states of NNV HNO3 / NO3

- Strong acid

NIV NO2, N2O4 Smog

NIII HNO2 / NO2- Weak acid / weak base

NII NO Smog + biology

NI N2O Greenhouse gas + laughing gas

N0 N2 Stable

N-I NH2OH Hydroxylamine

N-II N2H4 Hydrazine, rocket fuel

N-III NH3 / NH4+ Weak base / weak acid

Lecture 21 -23

Properties of N-compounds

HIGHLY VARIED!

Incredibly stable: N2

Extremely explosive:

nitroglycerine

N

N

CH3

NO

O

O

O

O

Otrinitrotoluene (TNT)

ON

OO

O

O

N

N

O

O

O

O

Strong acid HNO3

Weak base NH3

Photochemical smog: NO2

Biologically important: NO + amino acids

Lecture 21 -24

Nitrogen OxidesTable from

Silberberg, “C

hemistry”, M

cGraw

Hill, 2006.

Lecture 21 -25

Air pollution

Los Angeles

Sydney

The brown haze is largely NO2

Picture from www.consumercide.com

Picture from http://pdphoto.org

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