1. Reactions Oxidation-reductionof
Oxidation-reduction
prof. V. Paulauskas

2. Oxidation-reduction is a chemical process wherein the oxidation number of an element is changed
The abbreviated term used for oxidation-reduction is redox
Redox reactions
Redox process involves the complete transfer of electrons in the formation of ionic bonds, or partial transfer or shift of electrons in the formation of covalent bonds
prof. V. Paulauskas

3. Ag+(aq) + Cl– (aq) → Ag+Cl–(s)
AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)
IONIC REACTION
Ag+(aq) + Cl– (aq) → Ag+Cl–(s)
Element have the same oxidation number
2Na (s) + Cl2 (g) → NaCl (s)
REDOX REACTION
2Na0(s) + Cl20(g) → 2Na+Cl–(s)
Oxidation number of the element has changed
prof. V. Paulauskas

4. One electron is transferred from a neutral sodium atom to a neutral chlorine atom to form Na+ and Cl- ions that have filled-shell configuration ▫▫ ▫▫
Na + Cl → Na+ Cl – ▫ ▫▫ ▫ ▫▫ ▫▫ ▫▫ ▫▫
Sodium (alkali metals) has only one electron in its valence shell
Chlorine (halogens) is missing just one electron from having a complete shell configuration
Valence shell – the outermost shell of an atom
prof. V. Paulauskas

5. Oxidation numbers (ON) are positive and negative numbers assigned to the atoms of the elements in a compound –
they result from the transfer of electrons from one atom to another in ionic cmpds or sharing between in covalent cmpds
ON are not real charges - they are the results of an accounting method whereby we can keep track of electrons during a chemical reaction
Thus ON are defined using a premise, that the bonds in compounds for which you wish to assign oxidation numbers are assumed to be 100% ionic in nature
prof. V. Paulauskas

6. RULES for determination of ON:
1. Free element has always ON=0, e.g. ON of solid, metallic Cu is 0
2. Hydrogen in almost all compounds has ON=+1
3. Oxygen in compounds is always ON=-2, except with F or in peroxides
4. Alkali metals in compounds with non-metals have ON=+1
5. Alkaline earth metals in compounds have ON=+2
6. Halogens (except in cmpnds. with O or one another) have ON=-1
7. Charge of the ion or molecule equals sum of ON of all atoms
8. Algebraic sum of ON in a molecular compound must equal 0
prof. V. Paulauskas

7. Oxidation Is Losing electrons Reduction Is Gaining electrons
In oxidation, the ON of an element increases in a positive direction as a result of losing electrons
Oxidation
───────────────────────►
◄───────────────────────
Reduction
In reduction, the ON of an element decreases, becoming less positive and more negative, as a result of gaining electrons
OIL RIG
Oxidation Is Losing electrons 
Reduction Is Gaining electrons
prof. V. Paulauskas

8. the oxidation half and the reduction half
The loss and gain of electrons is a characteristic feature of all redox reactions +4 -2
S + O2 → SO2
Redox reaction may be written
as two half-reactions –
the oxidation half and the reduction half
prof. V. Paulauskas

9. O0 + 2e → O-2 S0 – 4e → S+4 reduction process oxidation process
oxidising agent
S0 – 4e → S+4
oxidation process
reducing agent
Substance that gains electrons (does the oxidising) – is called the oxidising agent (oxidator)
Typical oxidators have high ON – N+5
Substance that loses electrons (does the reducing) – is called the reducing agent (reductor)
Typical reductors have low ON – N-3
prof. V. Paulauskas

10. _________________________
Oxidation-reduction – a joint process, so in redox reactions:
the number of electrons lost by the reducing agent
must be equal
to the number of electrons gained by the oxidizing agent 3
│ S e → S-2 (oxidator)
│ N+2 – 3e → N+5 (reductor) 8
_________________________
±24e
3S N+2 → 3S N+5
prof. V. Paulauskas

11. Natural redox processes – burning, respiration, photosynthesys, soil formation processes
Industrial – metal extraction from ores, production of nitrogen fertilisers, fuel combustion, metal plating, chemical industries
Daily life – diesel combustion, galvanic processes (galvanic elements, lead batteries), metal corrosion
prof. V. Paulauskas

12. Types of Redox Reactions
In combination reactions two elements are combined
In a decomposition reaction, one element is broken down into its constitutive parts
In displacement reactions, one or more atoms is swapped out
In combustion reactions, a compound is combined with oxygen to produce heat
In disproportion reactions, a molecule may be simultaneously reduced and oxidized
prof. V. Paulauskas

13. Types of Redox Reactions
1. Intermolecular
2. Intramolecular
3. Dismutation
4. Comutation
prof. V. Paulauskas

14. ELECTROLYSIS
Process where electrical energy is used to bring about a chemical change
Application:
Manufacturing of sodium, chlorine
Production of pure hydrogen and oxygen
Purification of metals (silver)
Electroplating of metals
Charging of batteries
prof. V. Paulauskas

15. Metal chemical properties
Neutral metals are good reducing agents – in the course of reaction, free Me atom loses electrons to form Me2+ ion:
Meo – ne  Men+ reductor/oxidation
Cuo – 2e  Cu2+ reductor/oxidation
The more negative the standard reduction potential of the Me, the stronger are reductive (metalic) properties
prof. V. Paulauskas

16. Metalic properties are determined by electrones in the outer shell that can be easily lost:
Meo – ne  Men+ oxidation
With an increasing atom diameter, the metallic properties become stronger, as element is is much more likely to lose electrons:
Fr  Cs  Rb  K  Na  Li
prof. V. Paulauskas

17. Activity Series of the Metals
(Electromotive Series)
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
, , , ,0V +0, ,7
 METAL ACTIVITY DECREASING
 REDUCTIVE PROPERTIES DECREASING
It is a listing of metals in decreasing order of their reactivity with hydrogen-ion sources such as water and acids
Metal is oxidized to a metal ion, and the hydrogen ion is reduced to H2
prof. V. Paulauskas

18. Activity of Metals
Elements toward the bottom left corner of the periodic table are the metals that are the most active in the sense of being the most reactive
Lithium, sodium, and potassium all react with water.
The rate of this reaction increases as we go down this column, because these elements become more active as they become more metallic
prof. V. Paulauskas

19. Activity Series of the Metals
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
, , , ,0V +0, ,7 ↔
Very active Me
Non-reactive Me
Strong metallic properties
Weak metallic properties
It is related to the standard reduction potential of a Me cation
The more positive the standard reduction potential of the cation, the more difficult it is to oxidize the metal to a metal cation, and the later that metal falls in the series
prof. V. Paulauskas

20. Activity of the Metals Zn + 2HCl  ZnCl2 + H2 Cu + HCl  no reaction
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
All Me with a negative reduction potential will react with acids releasing hydrogen:
Zn + 2HCl  ZnCl2 + H2
Cu + HCl  no reaction
Very active metals (alkaline) react even with water to release hydrogen:
2Na + 2H2O  2NaOH + H2
prof. V. Paulauskas

21. Metal Displacement Reactions
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
A metal will displace (take the place of) a less reactive metal in a metal salt solution:
Zn + CuSO4  ZnSO4 + Cu
Cu + ZnSO4  no reaction
Mg + ZnSO4  MgSO4 + Zn
The greater the difference between metal potentials – the greater is the reaction rate
prof. V. Paulauskas

22. Metal Displacement Reactions
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
2AgNO3(aq) + Cu(s) → 2Ag(s) + Cu(NO3)2(aq)
Copper displaces silver in solution when a copper wire is dipped in a silver nitrate solution – solid silver precipitates out
prof. V. Paulauskas

23. Metal reaction with acids
When metal reacts with non-oxidizing acid (notably HCl and diluted H2SO4), a salt and hydrogen are produced:
metal + acid → salt + hydrogen
Fe + H2SO4  FeSO4 + H2
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
Hydrogen is included in the table to separate those metals that will not react with acids (the noble metals) from those that will
Metals higher than Hydrogen will always replace it in an acid
prof. V. Paulauskas

24. Metal reaction with acids
Li-K-Ca-Mg-Al-Zn-Cr-Fe-Cd-Pb-H-Cu-Hg-Ag-Pt-Au
Oxidising acids (notably concentrated sulfuric and nitric acids), react with metal to give much more interesting products
For example, concentrated nitric acid react with copper metal to give salt, water AND a particularly unfriendly gas, NO2
Cu + HNO3  Cu(NO3)2 + NO2 + H2O
Metals lower than Hydrogen (the noble metals) will also react with oxidising acids
Fe + H2SO4  Fe2(SO4)3 + H2S + H2O
prof. V. Paulauskas

25. Metal reactions with acids
Reduced products with oxidizing acids (HNO3, concentrated H2SO4) depends upon Me activity:
HNO3(N+5)  NO2(N+4), NO (N+2), N2 (N0), NH3(N-3) 
Cu …..… Pb……. Zn…….. Mg
prof. V. Paulauskas