Is an element which has half filled ‘d’ orbital.Transition Element are element which have at least one unpaired electron in the outer most sub energy level ‘d’. The transition element is an elements which have incompletely field d – orbital. The transition element is known as transition because have intermediate properties which differ from. ‘S’ block element and ‘P’ block element. Transition elements are d – block element, but not all d-block elements are transition elements. Most of transition element are metal element which called transition metal.
There are three series of transition metal:-
(i) Transition Metal of the first series.
(ii) Transition metal of the second series or LANTHANIDE METALS.
(iii) Transition Metal of third series or are sometime known as ACTINAMIDE METAL. (The strongest metal)
II. TRANSITION METALS OF THE FIRST SERIES
Are these which have half filled 3d – orbital are those which have at least ONE UNPAIRED ELECTRON IN SUB ENERGY LEVEL 3D.
Transition metals of the first series are given below
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GENERAL PROPERTIES OF TRANSITION METALS
Transition metals have the following general properties.
(i) Form colour (colour formation)
(ii) Are paramagnetic substance
(iii) Have variable oxidation states
(iv) Form complex compound
(v) Have catalytic action.
A. COLOUR FORMATION BY TRANSITION METALS
Transition metal appear coloured at room temperature of form colour when occur in ionic or combined state. But non – transition metal does not form colour at room temperature. The non – transition metal appear coloured when heated during flame test. The energy of flame excite electron which jump from low energy level. This cause atom to be unstable. In order to maintain the stability the effective nuclear force returns back the electron to the ground state. When the electron drop back emit the radiant energy which have a wavelength detected by a human eye. This radiant have definite colour.
But Magnesium does not produce colour during flame test instead remain colourless because energy of the flame is not enough to excite electron of Magnesium.
Transition metals colour formation can be explained by two theories.
(i)4s – 3d electron transition theory
(ii)d- Electron transition theory (crystal field theory)
(i) 4s – 3d ELECTRON TRANSITION THEORY
According to this theory colour formation is a result of movement of electron between sub energy level 4s and 3d. The energy difference between sub energy level 3d and 4s(E) is so small that normal radiant energy absorbed cause electron to move from 4s level to 3d level. Radiant energy absorbed from the sun makes 4s electrons to jump to 3d orbitals and the atoms becomes UNSTABLE. To maintain stability of the atoms excited electrons fall back to 4s orbital. During this process when electrons (excited) fall to their ground state heat energy is emitted and came out in form of radiation those wave have wavelength which are within the range that can be detected by human eyes with specified definite colour.
(ii) d – d ELECTRON TRANSITION THEORY OR CRYSTAL FIELD THEORY
Colour formation by transition metal can be explained by crystal field theory. This theory suggested that for transition metals to express colour there are two conditions which are necessary and must be fulfilled.
a.There must be at least one unpaired electrons in 3d – orbital
b.Presence of ligands.
PROCESS OF COLOUR FORMATION
In isolated Transition atoms all five 3d – orbitals are degenerate and have equal energy. In presence of ligands the d- orbitals split into two degenerates i.e.
“Double degenerate and Treble degenerate” and which have different energy. Ligands exert electric field (Repulsion force) to unpaired 3d electrons.Repulsion cause orbital to split down into two degenerates. Double and Treble degenerates.
The energy of separation (E) between the two degenerates is small. Such that the normal Radiant energy absorbed from the sun is enough to make electrons jump from treble degenerates orbital to double degenerates orbital and atoms excited, electrons fall back to their ground states, the process which is accompanied by emission of heat energy absorbed from the sun. The energy emitted by the falling electrons comes out in the form of radiation whose wave have specific definite colour.
Intensive of the colour whether to be faint or deep it depends on electric field and splitting power of the ligand. If the ligands exert weak electric field (Weak Repulsive forces) the d – orbital will split and separates to a small extent, The energy of separation E becomes very small.
If the energy of separation E between the two degenerates is small. Low energy will be required to excite electrons and make them jump to double degenerate orbital.
Similarly, low amount of heat energy will be emitted in the form of radiations when the excited electrons fall to their ground state.Intensity of the waves in the radiations will be low and hence the resulting colour appears to be light or faint. Ligands which exert weak electric field and cause small separation of the d orbital resulting to formation of light or faint colour are said to be “Ligands of high spin” .Ligands of high spin include Oxygen containing Halogens E.g. OH–, CL–, Br etc.
On the other hand if ligands exert strong electric field (strong repulsive force) the d – orbitals split down and will separate to a large extent. The energy of separation (E) becomes large. In this case large heat energy will be required to excite electrons. Equally well large energy is emitted in the form of radiation, when the electrons fall back to their ground state. Intensity of the waves in the coming radiation will be high and hence the resulting colour appear deep.Ligands which exert strong electric field and cause large separation of the d- orbitals resulting to the formation of deep colours are said to be “ligands of low spin” Ligands of low spin include Nitrogen containing compound e.g NH3, CN–, CN etc.
1.Ligands of high spin
·Extent weak electric field
·d- orbitals split to small extent
· E energy of separation is small
·The colour is faint or light.
2.Ligands of low spin
·Extent strong electric field
·D – orbitals separates to a large extent
· E is large (energy of separation)
·The colour is deep.
It has been noted that colour formation by transition metals can be explained by two Theories.
(a)4s – 3d electron transition theory
(b)Crystal field theory.
Now which one is the TRUE THEORY between 4s – 3d electron transition theory and crystal field theory. The theory of crystal field is the True Theory and accounts better for colour formation by Transition metals
The following fact justify the statement
(i) Transition metals express their colours whenever they are in ionic form. In ionic form 4s – orbital is always empty. Since during ionization of the transition metallic atoms 4s electrons are given before 3d – electrons. This means that movement of electrons between 4s and 3d orbitals is not possible.
(ii) 4s – 3d electron transition theory cannot account for change in colour intensity. But crystal field theory accounts for change in intensity depending on the type of ligands weather are of high or low spin.
(iii) 4s – 3d electron transition theory cannot explain why SC 3+, Zn 2+ and Cu+ are colourless in aqueous solution but crystal field theory can account for this observation.
(iv)Observation has shown that Sc+ and Sc2+ are coloured in aqueous solution because the two ions have one unpaired electron in sub energy level 3d.Sc = (Ar) 4S2 3d1 has no electron in 3d orbitals. All the five orbitals in Sc3+ are empty. I.e. Sc3+ = [Ar] 4S0 – 3d0.
When ligands come with their lone pairs there will be no repulsion since d – orbital have no electrons.So no splitting of the d – orbital and no d – electron transition in Sc3+ Cu+ and Zn2+ are colourless because all the five 3d orbitals on these ions are full occupied by two electrons ie [sub energy level 3d in Cu+ and Zn2+ have Ten electrons]. No unpaired electron in 3d – orbital of Zn2+ and Cu+.
Despite the fact crystal field theory explains better colour formation by transition metals than 4s -3d electron transition theory yet it has some weakness. Crystal field theory fails to account for colour formation in Manganese (ii) ion (Mn2+). Manganese (iii) ion has unpaired electrons in its 3d – orbitals but in aqueous solution it is colourless
Manganese in the permanganate ion [MnO4–] has no unpaired electrons in the 3d – orbitals. Manganese in the permanganate ion has an oxidation state of +7. Oxidation state of +7 is attained after losing all the 4s and 3d electrons. However permanganate, ions in aqueous solution are coloured, make the aqueous solution PURPLE.
[B] COMPLEX COMPOUND
Is a compound which contain central atom and several ligand. The complex compound consists of central atom attached to the several atoms or group of atoms .
Is a metallic ion which accepts or accommodate pair of electrons during formation of complex compound. Most of central atoms are transition element.
The characteristics features which result into a transition metal or another metal to form a complex compound include the following:-
· They have vacant orbit or empty orbital which accommodate pair of electrons.
· They have high nuclear charge which exerts nuclear attractive force.
· They have small atomic size which exerts a strong nuclear attractive force.
These central atom or metallic ions have constant total number of ligand accommodate during complex compound formation. This is according to the number of vacant orbitals and atomic size of metallic ions. If the central atoms accommodate ligand above those required result the compound to be unstable. The following include the central
atom together with their total number of ligand.
Ag+ – 2
Hg+2 – 4
Al+3 – 4
Zn+3 – 4
Pb+2 – 4
Cu+2 – 4
Cr+3 – 6
Mn+2 – 5/6
Fe+2/+3 – 4/6
Ni+2 – 4/6
CO+2 – 4/5
Pb+2 – 4/6
Is a non metallic ion or molecules which donate pair of electrons to the central metal atom/ion during complex compound / ion formation. The non metallic ions or molecules have pair of electrons in the valency orbital. These pair of electrons called Ione pair.
There are two kind of ligand :
(i) Neutral ligand is a ligand which is electrically neutral. These are not charged ligand Such ligands are generally the molecular species having one or more lone-pair of electrons.
Ethane – 1, 2 – diaminie [en] NH2 – CH2 – NH2
(ii) Anionic ligands: Are ligand which are electrically charged.
These are ligands which carry negative charge on them.These charged ligand include the following:-
Anionic ligands generally form anionic complex ions. For example [Ag]-, [Pb C]2-, [AL 2- , [Ag(CN]
Ligands can also be classified based on the mode of attachment.
(i)Monodentate (Unidentate) ligands: There are ligands which can attach to the central metal atom/ion only through one point.
(ii)Bidentate ligands:These are ligands which get attached to the metal ion through two points.
(iii)Tridentate ligands:These ligands get attached to the metal ion through 3 points.
FORMATION OF COMPLEX COMPOUND
The central atom provide vacant orbit and ligand provide pair of electrons. In order for a central atom to form complex compound,first should lose electrons equivalent to their valency which become metallic ions then attract ligand. For transition element which form complex compound use the following vacant orbitals. For central atom which accommodate four ligand use ns and np vacant orbitals which result nsp3 – hybridization.
The central atom form complex compound if the ligand occurs in excess or high concentration Al+3 in a dilute or low amount of NaOH form simple compound Al+3 + NaOH → Al[OH]3 + Na+
But when the ligand occur in excess or high concentration cause Al3+ to form complex compound.
TYPES OF COMPLEXES
1. Cationic complexes are complex compound which are electrically positively charged. The total charge of central atom or metallic ions and all ligand result into the compound to be positively charged. Example of cationic complexes [ Cr Cl (H2O)3(NH3)2]+
2. Anionic complexes are complexes compound which are negatively charged. The total charge of metallic ions and all ligand result into the compound to be negatively charged. The total charge of metallic ions and all charge ligand result into the compound to have negative charge (Fe (CN)6]-6.
3. Neutral complexes are complex compound which are electrically neutral and has no net charge. The total number charge of metallic ions and all ligand form the neutral compound. [ Al(NH3) (OH)3].
RULES FOR NAMING INORGANIC COMPLEXES
The system adopted is that developed by the international union of pure and applied chemistry ( IUPAC).
i) Cations are always named before anions
The oxidation states of the central metal atom or ion is shown in Roman numerals in brackets immediately after it’s name.
Eg. [ Fe (H20) 6] Cl3 Hexaaquairon (III)Chloride.
ii) Within complex ,ligands are named first followed by the central metal ion . Ligands are named in alphabetical order.
Eg. [Co(NH3)4Cl2) Tetraamminedichlocobalt (III) ion
iii) The number of particularly are ligands present must be specified by using the following prefixes.
di – 2 ligands
tri – 3 ligands
tetra – 4 ligands etc.
iv) Names of negative ligands end up with suffix O. eg. CN– cyano ,Cl– chloro etc.
Neutral ligands usually retain their normal names except for the special cases : H20(aqua),NH3(ammine) etc
v) Anionic complexes end up in the suffix – ate often appended to the to the Latin or English name of the metal
Eg, [Co (CN)6]3- Hexacyanocobaltate (III) ion
[Fe (CN)6]3- Hexacyanoferrate (III) ion
vi) Cationic complexes use up their English name unchanged for the central metal ion.
Eg,[Cr(NH30)4cl2]+ Tetraamminedichlorochromium (III) ion.
Name the following complexes
Potassium hexacyanoFerric (ii)
Tetra aminebromonitrocabalt (iii) sulphate
Hexaminechlomium (iii) trioxalatochromium (iii)
(d) [Al(H2O) (OH)5]-2
(1) + Cr + (-4) + (0) + 5(0) = 0
Cr = +3
(f) K [CrCL4(H2O) NH3] 5H2)
(1) + Cr + (-4) + (0) + (0) (5(0) = )Cr = +3.
Pentahydrate potassium amine aquatetrachlorochromate (iii)
(a) Complex compound [CO(NH3) 5(Br)SO4
(i) Name the complex compound above
(ii) What the coordinate number of central atom
(iii) If all ligand placed by chloride ligand what is the charge of complex compound.
(iv) Write isomers of compound
(b) Using hybridization principles prove the following.
Fe(CN)6 -4 d2 sp3 hybridized
CO F6 -3 d2sp3 hybridized
Ni(CN)S-3 dsp3 hybridized
(i) Diaminebromocobalt (iii) Sulphate
(ii) Co – ordinate number = 3
(v) [CO(NH3)Br]+2[COCl3] the complex compound so neutral
(vi) Isomers of complex compound
Isomers of ligand obtain through changing the number of each ligand
(b) [Fe(CN)6]-4 d2 sp3 hybridization
Provide the IUPAC name of the following compound
(a) [Pt(NH3)3 Cl3] Cl
(b) Na (Al(CN)2]
(c) K6 (C02)(en)10 4H2O
(d) [Ni(H2O)4] [Ni(CN)4]
(e) [Mn(NH3)6] Cl3
(f) Na3(Cr) CN)6]Cl3
Consider the complex compound
(a) What is the name of central atom
(b) List neutral ligand in the compound
(c) Name the charged ligand in the compound
(d) What is the name of complex compound
(e) What is the co – ordinate number of central atom
(a) Cr = Chromium
(b) NH3 = Amine
(c) Cl-1 = Chloro and C2O4-2
(d) Potassiumdiamminedichrolooxalatochromium (III)
(e) Co – ordinate number = 6
Write the formula of the following complex compounds
(i) Dichlorotetra ammine cobalt (iii) chloride Tetraammine aquacopper
(iii) Chloropentaaquacobalt (III) sulphate
(iv) Diaquatetraammine manganese (II) bromide
(v) Calcium di – iodo oxolatoferrate (III)
(vi)Trichlorotriammine chromium (III)
(vii) Trichlorotroammine planum (IV) chloride
(viii)Potassium disulphat atotetra aquachromate (III)
(ix)Tetrammine copper (II) sulphate monohydrate
(x) Potassium hepta axodichromate (vi)
(xi)Trichlorotriammine platinum (IV) chloride.
P =  +  +  = 1
[COCl2 (NH3)2] Cl
(ii) [ Cu(NH3)4H2O]+2SO4-2
 +  +  = 2
Reaction of the following salt with dilute NaOH
|Salt solution||Small quantity||Excess|
|Ag (OH) ppt
Zn (OH)2 ppt
Ca (OH)2 ppt
Soluble – Na2 (Zno)2)
Soluble – Na2 [pbo]2
Soluble – Na[ACO2]
NOTE: Precipitates of Zn, Pb and Al are able to dissolve in excess alkali because they form soluble complexes.
Reaction of the following salt with dil NH3 solution.
|Salt solution||Small quantity||Excess|
|Ag (OH) ppt + NH4 OH
Zn (OH)2 ppt + NH4 OH
Ca (OH) ppt + NH4 OH
Pb (OH)2 ppt + NH4 OH
Fe (OH)2 ppt + NH4 OH
Cu (OH)2 ppt + NH4 OH
Mg (OH)2 ppt + NH4 OH
|[Ag (NH3)2] NO3CO
Pb (OH)2 – Insoluble ppt
[Cu (NH3)4] (NO3)2
NOTE: All complex compounds are SOLUBLE
Bp is insoluble in excess NH3 (It form ppt)
Complete the following equations.
(i)Cu (OH)2(s) + NH3(ag) [Cu (NH3)4] (NO3)2
(ii)Ag Cl (s) + NH3 No reaction
(iii)Fe2 SO4 + KCN [Fe (CN)6] SO4
GEOMETRIC SHAPE OF COMPLEX COMPOUND
The number of ligands direct coordinated to the metal ion is the coordination number,of that atom in the complex.
(a) Coordination Number 6:
Complexes belonging to this coordination number have octahedral structure’
Eg, [Fe (CN)6]4- hexacyanoferrate (II) ion
This is an inner orbital complex because the cyano – ligands are so strong that they push unpaired electrons inwards and force them to pair up.
(b) [COF6]3- hexafluorocobaltate (iii) ion.
This is an outer orbital complex because the fluoro – ligands are weak. They cannot force electrons to pair inwards.
The six fluoro – ligands manage to occupy the orbitals in the 4th quantum shell.
(a) Octahedral structure of [Fe (CN) 6] 4-
(b) Octahedral structure of [COF 6] 3-
(b) Coordination number 5
Complexes with coordination number have trigonal bipyramidal structure .
eg. [Ni (CN) 5]3-
(C) Coordination number 4.
The most common complexes with this coordination number have square planar structure
eg. [Cu (NH3)4]2+
The shape 0f square planar complexes is flat
Write the geometrical shape of the following complex.
(c) [Fe(H2O)2 (CN)4]–
With exception to zinc, transition metals have catalytic action. They are used as catalysts in most chemical reactions catalytic power of catalyst is explained by their ability exist into stable oxidation states. In their catalytic ion, transition metals act as electron carries electrons in electron donors and pass them to electron acceptors.
Example iron and Nickel are catalyst during hydrogenation of alkene. Example V2O5, are used as catalyst during contact process. Example Manganese (iv)oxide used as catalyst during preparation of oxygen,iron used as a catalyst during Harber process.
(D) PARAMAGNETIC PROPERTY
Is a substance which can be induced with magnetism and can be attracted by magnetic bar. For a substance to be a paramagnetic it must possesses unpaired electrons in its electronic structure. Transitional metals with exception to zinc are paramagnetic. This is because they have unpaired electrons in 3d – orbital. Unpaired electrons form weak magnetic field and magnetic forces arise due to weak electric current so formed in the orbitals as a result of spinning of electrons in the orbitals.When a magnetic bar is brought near to a paramagnetic object there is a tendency of domains or unpaired electrons to shift from a side of weak magnetic field [from the object] to the side of strong magnetic field which is strong and the whole object appear to be attracted towards the [Magnetic field] magnet.
When electrons in the orbitals are paired there is interference and cancellation of magnetic fields occur as the magnetic field of two paired electrons move towards exactly opposite to each other (Vector) . All the magnetic fields cancel each other. Such substances cannot be induced with magnetic and cannot be attracted by a magnetic bar. Substances which cannot be induced with magnetic and cannot be attracted by a magnetic bar substances which cannot be attracted by an external magnetic field are said to be DIAMAGNETIC SUBSTANCES.
Transition metals are paramagnetic in pure metal or electrons are present in the delectronic structure of the transition metallic ion, for example Ferrous suphate [FeSO4] is paramagnetic just because Fe2+ has unpaired electrons in 3d orbitals.
This is because ligands of high spin when approaches the d – orbitals and cause small separation of the d – orbitals.The energy of separation ( E) between double degenerate and treble is small, electrons will spread in accordance to Hund’s Rule in which each electron occupy its orbital singly before pairing.
In that case the central transition metallic ion remains with unpaired electrons in the electronic structure. Hence the complex combined will be paramagnetic. Consider Fe2+ ion when complexes with ligands of high spin like water [H2O]
This complex [Fe(H2O4)4]2+ is a 4sp3 – hybrid complex and is paramagnetic. Since Fe2+ ion in the complex contains unpaired electrons in the d – electrons. If on the other hand Fe2+ complexes with ligands of low spin the paramagnetic property will be destroyed. The complex will be diamagnetic. Why?
This is because ligands of low spin split the d – orbitals and cause large separation.The energy of separation [ E] between the two degenerates becomes large.If the energy of separation is large electrons fill in the orbital in accordance to AUFBAU’S PRINCIPLE in which electron pair themselves in Treble degenerate orbitals, which have low energy according to AUFBAU’S PRINCIPLE. Consider Fe2+ complexes with ammonia which is ligand of low spin.This complex is 3d24sp3 hybrid complex and is diamagnetic since there are no unpaired electrons in electronic structure of Fe2+.All electrons Fe2+ are paired in Treble degenerate orbitals.
NECTA 2001 P2 Question 6c
Use the configuration of 3d – orbital electron on cobalt (iii) ion to explain why [COf6]3- is paramagnetic while [COCN6] is not paramagnetic? [10%]
Consider electronic structure of CO and CO3+ ions given below
Complex [CoF6] 3- is paramagnetic because fluoride ion is a ligand of high spin hence cause small magnetic field 3d – orbit of cobalt (iii) ion hence the energy separation is small and filling of electrons in according to Hund’s Rule, thus cobalt (iii) ion contains unpaired electrons in the 3d – orbitals. Hence a paramagnetic substance.
This is 4sp3 hybrid complex. And it is paramagnetic due to the in 3d – orbital
The complex [COCN6]3- is not paramagnetic due to absences of unpaired electrons in 3d – orbital in cobalt (iii) ion in the complex caused by complexing with ligands of low spin the cyano. This ligands exert strong magnetic field to the unpaired 3d – orbital hence large separation which cause excitation of be difficult hence filling of electrons is according to AUFBAU’S PRINCIPLE.
This complex is 3d24sp3 hybrid complex and is diamagnetic since there are no un paired electrons in electronic structure of Fe2+. All electrons Fe2+ are paired in treble degenerate orbitals.
NECTA 2001 P2 Question.6c
Use the configuration of 3d – orbital electron an cobalt (iii) ion to explain why [COf6]3- is paramagnetic while [COCN6] is not paramagnetic? [10%]
Consider electronic structure of CO and CO3+ ions given below.
Complex [CoF6] 3- is paramagnetic because fluoride ion is a ligand of high spin hence cause small magnetic field 3d – orbit of cobalt (iii) ion hence splits separates the hence the energy separation is small and filling of electrons in according to Hund’s Rule, thus cobalt (iii) ion contains unpaired electrons in the 3d – orbitals. Hence a paramagnetic substance.
This is 4sp3 hybrid complex. And it is paramagnetic due to the in 3d – orbital
The complex [COCN6]3- is not paramagnetic due to absences of compared electrons in 3d – orbital in cobalt (iii) ion in the complex caused by complexing with ligands of low spin the cyano. This ligands exert strong magnetic field to the unpaired 3d – orbital hence large separation which cause excitation of be difficult hence filling of electrons is according to AUFBAU’S PRINCIPLE.
Paramagnetic property of transition metals can be destroyed in two ways
a. Temperature/ oxidation process
b. Complex compound formation
a. TEMPERATURE/ OXIDATION
Raise in temperature destroy magnetic property of the elements. Raise in temperature causes excitations of electrons and ionization of the atoms by losing electron.
Ionization to an extend of losing all unpaired electrons destroys magnetic property of the substance.
b. COMPLEX COMPOUND FORMATION
Formation of complex compounds by coordination with ligand may also destroy magnetic property of transition metals. However this will depends on oxidation state of the metals and nature of the ligands involved in company compound formation whether are ligand of high or low spin.
When the complex compound formed by ligand of high spin such as C2O4-2, OH-, F, Br, 1-, Cl- energy separation between treble and degenerate. Electron filled according to the Hund’s rule this result electron to be unpaired in the double and treble degenerate the complex remain paramagnetic example COF6-3 is paramagnetic substance. The F- filled in 4s, 4p and 4d vacant orbital the 3d remain with unpaired electron which result paramagnetic
a. Predict the coordination number of Ni2+ and state whether the complex will be paramagnetic or diamagnetic it Ni2+ complex with
i) Bromine ions, Br – [Ligands of high]
ii) Ammonia molecules NH3 [ligands of low spin]
b. Fe3+ complexing with NH3 of low spin
C) VARIABILITY IN OXIDATION STATES
With exception to zinc, transition metals form more than one stable oxidation state. They have variable oxidation state. Variability in oxidation state in transition metals is explained by the ability to ionize by losing electron from both sub energy levels 4S and 3d.The energy present between 4s and 3d is very small. The gap existing between 4s and 3d is small. The different between the two degenerated is so small that just normal radiant energy from the sun is enough to excite electrons and so small that electrons from sub energy level 4s and 3d can be removed by almost the same amount of ionization energy. So that during ionization, transition
metals give off 4s – electrons first followed by 3d Transition metals depends on the number of electrons present in 4s – orbital. Chromium has lowest
oxidation state of +2, looses two electrons in 4s – orbital and copper the lowest oxidation state is +1. The maximum or the highest oxidation state is (achieved) attained after losing all the 3d electrons.
The extent of losing all the 3d – electrons is possible only for the first five elements from Sc to Mn. Elements beyond manganese can ionize to an extent of losing all the electrons because of large nuclear charge in these elements which cause strong effective charge pull, over the 3d electrons. The only oxidation state for elements found beyond manganese is +3 and for copper stable oxidation state is +2 which is attained after loss of the single 4s – electron and another from sub energy level 3d. The nuclear charge is high in zinc and hence the nuclear pull in 3d electron for zinc is maximum. Hence it is difficult to remove electrons from sub energy 3d in zinc. Zn has only one stable oxidation state of +2 which is formed after losing the two electrons. Only one stable oxidation state for zinc is one of the factors which excludes zinc from transition metals.
For these elements with atomic number 21 to 25 which can ionize to an extent of losing all the d – electrons, the tendency is that increase in oxidation states accompany with increase in acidic character. Lower oxidation states have basic character and higher oxidation states are acidic in nature. This is caused by increase in ionization energy needed to form the ions.The increase in acidic character with increasing oxidation states can be justified by considering the various oxides of chromium.
The various oxides of chromium include.
Base oxides amphoteric oxides Acidic oxides
The first two oxides, chromium (iii) oxide [Cr2O] and chromium (ii) oxide [CrO] are basic solutions which have no chemical reaction with other basic solution even alkaline solutions.
react with acidic solution
Chromium (ii) chloride
Chromium (iv) oxide [Cr2O3] is amphoteric. It dissolve in water and form hydroxides which dissolve in both alkaline solutions and acidic solutions.
Chromium (iv) hydroxide react acid solution
Chromium (iv) hydroxide dissolves in alkaline solution and form complex salt
[Cr (0H) 6
NaCrO3 + 3H20
Sodium chromide (iii)
The last two oxides, chromium (v) oxide Cr205 and chromium (vi) oxide (CrO3) are ACIDIC
OXIDES. They dissolve in water and form strong acidic solutions
CrO3 + H2O H2CrO4
Chromic acid is so strong acid that has no chemical reaction with other acidic solution
H2CrO4 + H2SO4 reaction
Chromic acid reacts with basis solutions and form normal salts
Reaction of the various oxides of chromium real justify with increase in Acidic character.
Change in base acidic character in transition metals during ionization explain why Redox reactions take place either in acidic medium or basic medium, lower
oxidation states which are basic in nature, become more stable if the medium in which they are formed is acidic. On the other hand, high oxidation states which are acidic tend to be more stable if they are formed in basic medium. Therefore in redox reactions, when the elements change its oxidation states from high to lower oxidation state the process will take place in acidic medium from low oxidation state to high oxidation state, the process will occur in basic medium.
i. Low oxidation state High oxidation state
Basic in nature Mn2+ Mn04–
ii. High oxidation state lower oxidation state
[Acid] MnO4-1 Mn2+ [base]
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