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Oxidation States

Let's take an example of an atom that is made up of one proton and one electron, is called hydrogen. The proton and electron stay together because just like two magnets attracted from opposite poles. Have you ever thought why these two particles not colloid? The electron in an atom is constantly spinning in a certain energy level around the center of the atom, called the nucleus. Electrons are placed in fixed energy levels, known as orbitals. A compound is a pure substance consisting of two or more different elements.

Do you know, why elements combined to form element? We know that each element has a certain value of valency; means combining capacity of element with another element which is due to electrons present in the outermost shell of the element. These electrons are known as valence electrons. An atom always tries to make eight electrons in its outermost shell; that is called Octet rule. Hence, the valency of an element depends on valence electrons. The atoms of elements, having a completely filled outermost shell show little chemical activity. Or we could also say that their combining capacity or valency is zero. Like inert elements, the helium atom has two electrons in its outermost shell and all other elements have atoms with eight electrons in the outermost shell. They all commonly termed as inert elements. 

An atom having an incomplete outer shell converts in octet configuration by:
  • It can gain electrons from another atom.
  • It can lose electrons to another atom.
  • It can share one or more electron pairs with another atom.
Let's discuss about oxidation state of the elements.
Promotion of Electrons

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Oxidation State Calculator
 

Oxidation State Definition

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Oxidation state of any atom shows the ability of an atom to oxidize (to lose electrons) or to reduce (to gain electrons) other atoms or species in a chemical compound. It is related to the number of electrons which an atom loses, gains, or appears to use when bonding with another atom in a compound. It’s a difference between the number of electrons of the same atom in a compound, as compared with the number of electrons in an atom of the element. The oxidation state is the ionic charge on ions and the formal charge in the case of covalent compounds. For example, in NaCl the oxidation states of Na and Cl are +1 and -1 respectively and in CCl4 the oxidation state of Carbon is (+4) and (-1) for each chlorine.

Oxidation State of Transition Metals

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All the transition elements, except the first and last members of the series, exhibit a number of oxidation states. They show variable valency in their compounds. Some basic oxidation states of the first, second and third transition series elements are given below in tables.

Oxidation states of 3d series


Elements
Outer electronic
configuration
Oxidation states
Sc 3d14s2
+2, +3
Ti
3d34s2 +2, +3, +4
V
3d34s2
+2, +3, +4, +5
Cr
3d54s1
+2, +3, +4, +5, +6
Mn
3d54s2
+2, +3, +4, +5, +6, +7
Fe
3d64s2
+2, +3, +4, +5, +6
Co
3d74s2 +2, +3, +4
Ni
3d84s2 +2, +3, +4
Cu
3d104s1 +1, +2
Zn
3d104s2 +2

Oxidation states of 4d series


Elements
Oxidation states
Y +3
Zr
+3, +4
Nb
+2, +3, +4, +5
Mo
+2, +3, +4, +5, +6
Tc
+2, +4, +5, +7
Ru
+2, +3, +4, +5, +6, +7, +8
Rh
+2, +3, +4, +6
Pd
+2, +3, +4
Ag
+1, +2, +3
Cd
+2

Oxidation states of 5d series


Elements
Oxidation states
La
+3
Hf
+3, +4
Ta
+2, +3, +4, +5
W
+2, +3, +4, +5, +6
Re
+1, +2, +4, +5, +7
Os
+2, +3, +4, +6, +8
Ir
+2, +3, +4, +6
Pt
+2, +3, +4, +5, +6
Au
+1, +3
Hg
+1, +2


Cause for variable oxidation states


The valence electrons of the transition elements are in (n-1) d and ns orbitals which have little difference in energies. Both energy levels can be used in bond formation. They show the +2 oxidation state due to the 2 electrons in ns orbitals when the electrons of (n-1) d remain unaffected. 

The higher oxidation state from +3 to +7 is due to the use of all 4s and 3d electrons in the transition series of elements. In the excited state, the (n-1) d electrons become bonding and give the variable states to the atom. Thus the variable oxidation state is due to the participation of both ns and (n-1) d orbitals in bonding. 

Some important features of oxidation state of transition metals are,
  1. The most common oxidation state of 3d series is +2 except scandium, due to the loss of two ns electrons. This shows that d orbitals are more stable than s orbitals after scandium.
  2. The ionic bonds are generally formed in +2 and +3 state while the covalent bonds are formed in higher oxidation states. Covalent bonds are formed by the sharing of d-electrons. For example- permanganate ion MnO4, all bonds formed between manganese and oxygen are covalent.
  3. The highest oxidation state increases with increasing atomic number of element, reaches maximum in the middle and then starts decreasing as shown in the table. For example, iron shows the common oxidation state of + 2 and + 3, but ruthenium and osmium in the same group form compound in the +4, + 6 and + 8 oxidation states.
  4. The elements in the beginning of the series exhibit less oxidation state because of having small number of electrons to lose or contribute. The elements in the middle of the series show the greatest number of oxidation. For example; Mn shows all the oxidation states from +2 to +7. The highest oxidation state shown by any transition metal is eight which is shown by Ru and Os.
  5. The maximum oxidation state of reasonable stability in 3d series is equal to the sum of s and d electrons up to Mn followed by decrease in the stability of higher oxidation states.
  6. The variability of oxidation states arises in such a way that successive states differ in unity.
  7. The higher oxidation states are more stable in heavier elements. Like in Mo (VI), W (VI) are more stable than Cr (VI).
  8. They also form compounds in low oxidation states such as +1 and 0 having ligands like CO. For example, nickel in nickel tetra carbonyl, Ni(CO)4 and Fe in Fe(CO)5 has zero oxidation state. Cu is in +1 state in CuCl.
  9. The oxidation state of a metal in solvent depends on the nature of the solvent. For example, Cu in +1 state in water is unstable as it may undergo oxidation while Cu in +2 states is stable.
  10. The oxidation state of transition metals depends on the nature of the combining atoms. So the compounds of metal with fluorine and oxygen exhibit the highest oxidation state due to their small size and the high electro negativity of fluorine and oxygen.
  11. The relative stabilities of elements which exist in more than one oxidation state can be known from the standard electrode potential.

Determining Oxidation States

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The process of determining the oxidation states is based on knowing the elements which can have only one oxidation state other than in the elemental state and which elements are able to exist in more than one oxidation state other than in the elemental state. These are some of the "rules" for assigning the oxidation states.

Rules for Assigning Oxidation States

  • The oxidation state of free elements is always zero. For example, Fe(s), O2(g), O3(g), H2(g), Hg(l), Hg(g), S(s) etc.
  • The oxidation state of metals is taken as the charge of the ion. For example, Group IA, Group IIA, Group IIIA (except Tl), Zn2+, Cd2+ forms only one ion.
  • In the case of monatomic ions, either anions or cations, the charge on the ion is considered as its oxidation state. For example, Cl- (-1), Fe2+ (+2), Fe3+(+3), S2- (-2), H+ (+1) etc.
  • In the case of compounds having polyatomic ions, the overall charge of the polyatomic ion is used to determine the charge of the ion.
  • Some elements have common oxidation states.
  1. Alkali metals (Li+, Na+, K+) are always in +1 oxidation state while alkali earth metals (Mg2+, Ca2+, Sr2+, Ba2+) are in +2 state.
  2. Hydrogen is in +1 state (except in metal hydride compounds such as LiH).
  3. Oxygen is in (-2) state (except in peroxides such as H2O2).
  4. Halogens (F-, Cl-, Br-, I-) are generally in -1 state.
  • The total sum of the oxidation states in a molecule is always zero. For example; In CH2O, the sum of the oxidation state is (0) + 2(+1) + (-2) = 0. Similarly, in CH3OH, sum of the oxidation states= (+2) + 3(+1) + (-2) + (+1) = 0
For example, determining the oxidation states of the elements in the compound KMnO4. By applying rule, the oxidation state of K = +1 and four oxygen is in -2 state. 

Let the oxidation state of Mn = x as manganese can be in several oxidation states.

As the overall charge of the compound is zero. So the algebraic expression can be written as

1 + x - 8 = 0

So the oxidation state of manganese is x - 7 = 0 or x = +7.

Similarly the oxidation state of chromium in dichromate ion, (Cr2O72-).

Since the oxidation state for chromium is not known, so taken it as 2x for chromium
Cr2O7 2-
2x 7(-2)

Since the overall charge of the ion is -2, so the algebraic equation for x
2x + 7(-2) = -2

After solving it for x:
2x - 14 = -2
2x = 12
x = +6, so each chromium is +6 oxidation state.

Iron Oxidation States

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Iron has atomic number 26 and the electronic configuration is [Ar] 3d6 4s2.Iron has 4 unpaired electrons in 3d orbital and 2 paired electrons in 4s. So the possible oxidation states are + 2, + 3, + 4, + 6 but generally it shows +2 and +3 oxidation state.

Iron Oxidation States
  1. Iron (II) (Ferrous ion) - Iron in the +2 oxidation state is known as the ferrous ion which is pale green in color. Ferrous ion is easily oxidized to the ferric ion in the presence of a small amount of oxygen. So the solutions of ferrous ion are sometimes used as reducing agents. The ferrous ion complexes usually have octahedral geometry. For example, hexaaquo and hexacyano complexes of Fe+2. These ions have a particular affinity for amine ligands.
  2. Iron (III) (Ferric ion) - Iron in the +3 oxidation state is known as the ferric ion. The solutions containing the ferric ion are usually yellow or yellow-brown and acidic in nature due to the formation of species such as [Fe(H2O)5(OH)]2+. They form a blood-red complex ion with the thiocyanate (SCN-) ion which is used for testing the presence of the ferric ion.

Oxidation State of Chromium

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Chromium has atomic number 24 and electronic configuration is [Ar] 3d5 4s1 due to more stable half filled orbitals. They show the oxidation state + 2, + 3, + 4, + 5, + 6. Generally, Chromium forms compounds with oxidation state +2, +3 and +6. The +2 state of chromium is easily oxidized in aqueous solution to the +3 state. The +3 and +6 states are more stable in chromium.
  • Chromium (III) – They can be obtained by dissolving the element chromium in acids like HCl or H2SO4. Generally they form the octahedral complexes. For example-the dark green complex chromium(III) chloride hydrate [CrCl2(H2O)4]Cl, Chromium(III) hydroxide (Cr(OH)3), chromium(III) oxide (Cr2O3), Anhydrous chromium(III) chloride (CrCl3) etc.
  • Chromium (VI)- These compounds are powerful oxidants at low or neutral pH. For example- chromate anion (CrO42-) and dichromate (Cr2O72-) anions. Their halides are also known like hexafluoride CrF6, chromyl chloride (CrO2Cl2). Other examples are Chromic acid (H2CrO4), chromium (VI) oxide CrO3 which is the acid anhydride of chromic acid.
  • Chromium (V) – Only few compounds are known. For examples-peroxochromate (V), chromium (V) fluoride (CrF5) etc.
  • Chromium (IV) - They are more common than Cr(V). For example- tetra halides like CrF4, CrCl4 and CrBr4 etc.
  • Chromium (II) – For example- CrCl2, chromos acetate (Cr2(O2CCH3)4) etc.

Oxidation State of Silver

The electronic configuration of silver is [Kr] 4d10 5s1. The most common oxidation state of silver is +1 and +2 state. For example- AgNO3 silver(I) nitrate, Silver(II) fluoride (AgF2), potassium tetra fluoro argentate K[AgF4] for +3 state of silver etc.
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