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:
Blackman, Bottle, Schmid, Mocerino and Wille,
Chemistry, John Wiley & Sons Australia, Ltd. 2008
ISBN: 9 78047081 0866

Textbook: Blackman, Bottle, Schmid, Mocerino & Wille,
“ Chemistry”,
John Wiley
& Sons Australia, Ltd., 2008.
Today’s lecture is in
Potassium atom, K
19 protons, 19 neutrons
19 electrons
 Section 4.6, 4.8
 Section 12.1
 Section 13.1, 13.2
Lecture 21 -3


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 ).
Lecture 21 -4

 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+).
Lecture 21 -5


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).
Lecture 21 -6

Lecture 21 -7

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
Lecture 21 -8

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
Lecture 21 -9




Molecules and polyatomic ions: shared electrons are assigned to the more electronegative atom.
Examples: HF
CO
2
CH
NO
4
3
-
F -I
O -II
H I
C +IV
H +I C -IV
O=C=O
-1 charge on anion
= 3 x O -II + N V
H F
Determining an atom’s oxidation number:
H
H-C-H
H
2.
3.
4.
1.
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
Lecture 21 -10

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)
[Cr
2
O
7
] 2

2( x ) + 7(-2) = -2, x = +6, Cr(VI)
Lecture 21 -11

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

 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
Lecture 21 -13

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

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

In general, the ( n -1)d orbitals are filled between the n s and n p orbitals.
Lecture 21 -16
Blackman Figure 4.29



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 -17

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

Mn = [Ar]4s 2 3d 5
7 valence electrons
Orbital Occupancy
Lecture 21 -19

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 solutions.
, or slowly by O
2 in alkaline
Lecture 21 -20


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 -21

N V 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
N I
N 0
N
2
O
N
2
Greenhouse gas + laughing gas
Stable
N -I Hydroxylamine
N -II
N -III
NH
2
OH
N
2
H
4
NH
3
/ NH
4
+
Hydrazine, rocket fuel
Weak base / weak acid
Lecture 21 -22

Incredibly stable:
Extremely explosive:
Strong acid
Weak base
Photochemical smog:
Biologically important:
O
O
N
CH
3
O
N
O
N
2
N
O
O trinitrotoluene (TNT)
HNO
3
NH
NO
3
2
NO + amino acids
O N
O
O O
O
O
N O nitroglycerine
Lecture 21 -23
N
O
O

Lecture 21 -24

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






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
Lecture 21 -26