CHEMISTRY A LEVEL(FORM SIX) NOTES – INORGANIC CHEMISTRY 1.2-TRANSITION ELEMENT

INORGANIC CHEMISTRY 1.2 – TRANSITION ELEMENT

Is an element which has half-filled ‘d’ orbital. Transition elements are elements which have at least one unpaired electron in the outermost sub-energy level ‘d’. The transition element is an element which has incompletely filled d-orbital. The transition element is known as transition because it has intermediate properties which differ from ‘S’ block element and ‘P’ block element. Transition elements are d-block elements, but not all d-block elements are transition elements. Most transition elements are metal elements called transition metals.

There are three series of transition metals:

1. Transition Metal of the first series.
2. Transition metal of the second series or Lanthanide metals.
3. Transition Metal of the third series or sometimes known as Actinide metals (The strongest metal).

II. TRANSITION METALS OF THE FIRST SERIES

Are those which have half-filled 3d orbital and have at least one unpaired electron in sub-energy level 3d.

Transition metals of the first series are given below:

GENERAL PROPERTIES OF TRANSITION METALS

Transition metals have the following general properties:

• Form colour (colour formation)
• Are paramagnetic substances
• Have variable oxidation states
• Form complex compounds
• Have catalytic action

A. COLOUR FORMATION BY TRANSITION METALS

Transition metals appear coloured at room temperature or form colour when they occur in ionic or combined state. But non-transition metals do not form colour at room temperature. The non-transition metals appear coloured when heated during flame test. The energy of flame excites electrons which jump from low energy level. This causes the atom to be unstable. To maintain stability, the effective nuclear force returns the electron to the ground state. When the electron drops back, it emits radiant energy which has a wavelength detected by the human eye. This radiant energy has a definite colour.

Magnesium does not produce colour during flame test; instead, it remains colourless because the energy of the flame is not enough to excite electrons of magnesium.

Transition metals colour formation can be explained by two theories:

1. 4s – 3d electron transition theory
2. 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 electrons between sub-energy levels 4s and 3d. The energy difference between sub-energy levels 3d and 4s (ΔE) is so small that normal radiant energy absorbed causes electrons to move from 4s level to 3d level. Radiant energy absorbed from the sun makes 4s electrons jump to 3d orbitals and the atoms become unstable. To maintain stability, excited electrons fall back to 4s orbital. During this process, when excited electrons fall to their ground state, heat energy is emitted in the form of radiation with wavelengths detectable by human eyes, producing a specific definite colour.

(ii) d – d ELECTRON TRANSITION THEORY OR CRYSTAL FIELD THEORY

Colour formation by transition metals can be explained by crystal field theory. This theory suggests that for transition metals to express colour, two conditions must be fulfilled:

• There must be at least one unpaired electron in 3d orbital.
• Presence of ligands.

PROCESS OF COLOUR FORMATION

In isolated transition atoms, all five 3d orbitals are degenerate and have equal energy. In the presence of ligands, the d-orbitals split into two groups: double degenerate and triple degenerate, which have different energies. Ligands exert an electric field (repulsive force) on unpaired 3d electrons. This repulsion causes the orbitals to split into two degenerates: double and triple degenerates.

The energy of separation (ΔE) between the two degenerates is small such that normal radiant energy absorbed from the sun is enough to make electrons jump from triple degenerate orbital to double degenerate orbital and atoms become excited. Electrons fall back to their ground states, accompanied by emission of heat energy absorbed from the sun. The energy emitted by the falling electrons comes out in the form of radiation with specific definite colour.

NOTE:

The intensity of the colour, whether faint or deep, depends on the electric field and splitting power of the ligand. If the ligands exert a weak electric field (weak repulsive forces), the d-orbitals split to a small extent. The energy of separation ΔE becomes very small. Low energy is required to excite electrons and make them jump to the double degenerate orbital. Similarly, a low amount of heat energy is emitted in the form of radiation when the excited electrons fall to their ground state. The intensity of the waves in the radiation will be low, and hence the resulting colour appears light or faint. Ligands which exert weak electric fields and cause small separation of the d-orbitals resulting in light or faint colour are called “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 a strong electric field (strong repulsive force), the d-orbitals split and separate to a large extent. The energy of separation (ΔE) becomes large. Large heat energy is required to excite electrons, and equally large energy is emitted in the form of radiation when electrons fall back to their ground state. The intensity of the waves in the radiation will be high, and hence the resulting colour appears deep. Ligands which exert strong electric fields and cause large separation of the d-orbitals resulting in deep colours are called “ligands of low spin.” Ligands of low spin include nitrogen-containing compounds, e.g., NH₃, CN^–, etc.

1. Ligands of high spin

• Exert weak electric field
• d-orbitals split to a small extent
• ΔE energy of separation is small
• The colour is faint or light

2. Ligands of low spin

• Exert strong electric field
• d-orbitals separate to a large extent
• ΔE energy of separation is large
• The colour is deep

NOTE:

It has been noted that colour formation by transition metals can be explained by two theories:

1. 4s – 3d electron transition theory
2. 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 facts justify the statement:

1. 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.
2. 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, whether they are of high or low spin.
3. 4s – 3d electron transition theory cannot explain why Sc³⁺, Zn²⁺, and Cu⁺ are colourless in aqueous solution, but crystal field theory can account for this observation.
4. 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-orbitals have no electrons. So no splitting of the d-orbitals and no d-electron transition in Sc³⁺. Cu⁺ and Zn²⁺ are colourless because all the five 3d orbitals on these ions are fully occupied by two electrons, i.e., the sub-energy level 3d in Cu⁺ and Zn²⁺ have ten electrons. No unpaired electron in 3d orbital of Zn²⁺ and Cu⁺.

NOTE:

Despite the fact that crystal field theory explains colour formation by transition metals better than 4s – 3d electron transition theory, it has some weaknesses. 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 and make the aqueous solution purple.

[B] COMPLEX COMPOUND

A complex compound is a compound which contains a central atom and several ligands. The complex compound consists of a central atom attached to several atoms or groups of atoms.

Central Atom

Is a metallic ion which accepts or accommodates pairs of electrons during formation of complex compounds. Most central atoms are transition elements.

The characteristic features which result in a transition metal or another metal forming a complex compound include the following:

• They have vacant orbitals or empty orbitals which accommodate pairs 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 atoms or metallic ions have a constant total number of ligands accommodated during complex compound formation. This is according to the number of vacant orbitals and atomic size of metallic ions. If the central atoms accommodate ligands above those required, the compound becomes unstable. The following include the central atoms together with their total number of ligands:

• Ag⁺ – 2
• Hg²⁺ – 4
• Al³⁺ – 4
• Zn³⁺ – 4
• Pb²⁺ – 4
• Cu²⁺ – 4
• Cr³⁺ – 6
• Mn²⁺ – 5/6
• Fe^2+/3+ – 4/6
• Ni²⁺ – 4/6
• CO²⁺ – 4/5
• Pb²⁺ – 4/6

LIGAND

A ligand is a non-metallic ion or molecule which donates a pair of electrons to the central metal atom/ion during complex compound/ion formation. The non-metallic ions or molecules have pairs of electrons in the valence orbital. These pairs of electrons are called lone pairs.

There are two kinds of ligands:

1. Neutral ligand: electrically neutral ligands, generally molecular species having one or more lone pairs of electrons.
   Examples: Amine NH₃, Aqua H₂O, Carbonyl CO, Ethane-1,2-diamine [en] NH₂-CH₂-CH₂-NH₂.
2. Anionic ligands: electrically charged ligands carrying negative charges.
   Examples: Fluoro F^–, Chloro Cl^–, Iodo I^–, Cyano CN^–, Hydroxo OH^–, Nitro NO₂^–, Nitrate NO₃^–, Carbonate CO₃^2-, Sulphate SO₄^2-, Oxalato C₂O₄^2-.

Anionic ligands generally form anionic complex ions. For example, [Ag(CN)]^–, [Pb(CN)]^2-, [Al(CN)]^2-, [Ag(CN)]^–.

Ligands can also be classified based on the mode of attachment:

1. Monodentate (Unidentate) ligands: attach to the central metal atom/ion through one point.
2. Bidentate ligands: attach to the metal ion through two points.
3. Tridentate ligands: attach to the metal ion through three points.
4. Tetradentate ligands: attach through four points.
5. Pentadentate ligands: attach through five points.

FORMATION OF COMPLEX COMPOUND

The central atom provides vacant orbitals and ligands provide pairs of electrons. For a central atom to form a complex compound, it should first lose electrons equivalent to its valency to become metallic ions, then attract ligands. For transition elements which form complex compounds, use the following vacant orbitals. For central atoms which accommodate four ligands, use ns and np vacant orbitals which result in sp³ hybridization.

The central atom forms complex compounds if the ligand occurs in excess or high concentration. For example, Al³⁺ in a dilute or low amount of NaOH forms a simple compound Al(OH)₃ + Na⁺.

But when the ligand occurs in excess or high concentration, Al³⁺ forms complex compounds.

TYPES OF COMPLEXES

1. Cationic complexes: Complex compounds which are electrically positively charged. The total charge of central atom or metallic ions and all ligands results in the compound being positively charged. Example: [CrCl(H₂O)₃(NH₃)₂]⁺.
2. Anionic complexes: Complex compounds which are negatively charged. The total charge of metallic ions and all ligands results in the compound being negatively charged. Example: [Fe(CN)₆]^3-.
3. Neutral complexes: Complex compounds which are electrically neutral and have no net charge. The total charge of metallic ions and all ligands form a neutral compound. Example: [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).

1. Cations are always named before anions. The oxidation states of the central metal atom or ion are shown in Roman numerals in brackets immediately after its name.
   Example: [Fe(H₂O)₆]Cl₃ Hexaaquairon (III) chloride.
2. Within the complex, ligands are named first followed by the central metal ion. Ligands are named in alphabetical order.
   Example: [Co(NH₃)₄Cl₂] Tetraamminedichlorocobalt (III) ion.
3. The number of particular ligands present must be specified by using the following prefixes:
   di – 2 ligands
   tri – 3 ligands
   tetra – 4 ligands, etc.
4. Names of negative ligands end with the suffix -o. E.g., CN^– is cyano, Cl^– is chloro, etc. Neutral ligands usually retain their normal names except for special cases: H₂O (aqua), NH₃ (ammine), etc.
5. Anionic complexes end with the suffix -ate, often appended to the Latin or English name of the metal.
   Example: [Co(CN)₆]^3- Hexacyanocobaltate (III) ion.
   [Fe(CN)₆]^3- Hexacyanoferrate (III) ion.
6. Cationic complexes use the English name unchanged for the central metal ion.
   Example: [Cr(NH₃)₄Cl₂]⁺ Tetraamminedichlorochromium (III) ion.

Questions

Name the following complexes:

1. K₄[Fe(CN)₆]
   Potassium hexacyanoferric (II)
2. [Co(NH₃)₄(NO₂)]SO₄
   Tetraaminebromonitrocobalt (III) sulphate
3. (Cr(NH₃)₆)(Cr(C₂O₄)₃)^3-
   Hexaminechloro(III) trioxalatochromium (III)
4. [Al(H₂O)(OH)₅]^2-
   Aquopentahydroxoaluminium (III)
5. [Pt(en)₂Cl₂]^2-
   Dichlorodimethylamine platinum
6. K[CoCl₄(H₂O)NH₃]5H₂O
   Pentahydrate potassium amine aquatetrachlorochromate (III)

Questions

1. Complex compound [Co(NH₃)₅(Br)]SO₄
   (i) Name the complex compound above.
   (ii) What is the coordination number of the central atom?
   (iii) If all ligands are replaced by chloride ligands, what is the charge of the complex compound?
   (iv) Write isomers of the compound.
2. Using hybridization principles, prove the following:
   Fe(CN)₆^-4 is d² sp³ hybridized.
   CoF₆^-3 is d² sp³ hybridized.
   Ni(CN)₆^-3 is dsp³ hybridized.

Solution

1. [Co(NH₃)₅Br]SO₄
   (i) Diaminebromocobalt (III) sulphate
   (ii) Coordination number = 3
   (iii) [Co(NH₃)Br]²⁺ [CoCl₃] the complex compound is neutral.
   (iv) Isomers of complex compound:
   [Co(NH₃)₂Br]²⁺, [CoNH₃Br₂]⁺. Isomers of ligand obtained through changing the number of each ligand.
2. [Fe(CN)₆]^-4 d² sp³ hybridization.

Question:

Provide the IUPAC name of the following compounds:

1. [Pt(NH₃)₃Cl₃]Cl
2. Na[Al(CN)₂]
3. K₆(CO)(en)₁₀4H₂O
4. [Ni(H₂O)₄][Ni(CN)₄]
5. [Mn(NH₃)₆]Cl₃
6. Na₃(Cr)CN₆Cl₃
7. Na₂[Zn(CN)₄]
8. [Cr(H₂O)₅Cl]Cl₂.H₂O
9. CrCl₂[H₂O]₄Cl.2H₂O

Question

Consider the complex compound K[Cr(NH₃)₂Cl₂CO₄]

1. What is the name of the central atom?
2. List neutral ligands in the compound.
3. Name the charged ligands in the compound.
4. What is the name of the complex compound?
5. What is the coordination number of the central atom?

Solution

1. Cr = Chromium
2. NH₃ = Amine
3. Cl^– = Chloro and C₂O₄^2- = Oxalato
4. Potassium diamminedichloroooxalatochromium (III)
5. Coordination number = 6

Question

Write the formula of the following complex compounds:

1. Dichlorotetraammine cobalt (III) chloride, Tetraammine aquacopper sulphate
2. Sulphate
3. Chloropentaaquacobalt (III) sulphate
4. Diaquatetraammine manganese (II) bromide
5. Calcium di-iodo oxalatoferrate (III)
6. Trichlorotriammine chromium (III)
7. Trichlorotriammine platinum (IV) chloride
8. Potassium disulphato tetra aquachromate (III)
9. Tetrammine copper (II) sulphate monohydrate
10. Potassium hepta oxodichromate (VI)
11. Trichlorotriammine platinum (IV) chloride

Solution

1. [CoCl₂(NH₃)₄]Cl₃
2. [Cu(NH₃)₄H₂O]SO₄

Reaction of the following salts with dilute NaOH

NOTE: Precipitates of Zn, Pb, and Al are able to dissolve in excess alkali because they form soluble complexes.

Reaction of the following salts with dilute NH₃ solution

NOTE:

All complex compounds are soluble. Pb is insoluble in excess NH₃ (It forms ppt).

Complete the following equations:

1. Cu(OH)₂(s) + NH₃(aq) → [Cu(NH₃)₄](NO₃)₂(aq)
2. AgCl(s) + NH₃ → No reaction
3. FeSO₄ + KCN → [Fe(CN)₆]SO₄

GEOMETRIC SHAPE OF COMPLEX COMPOUND

The number of ligands directly coordinated to the metal ion is the coordination number of that atom in the complex.

(a) Coordination Number 6:

Complexes with this coordination number have octahedral structure.
Example: [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) Coordination Number 5:

Complexes with this coordination number have trigonal bipyramidal structure.
Example: [Ni(CN)₅]^3-.

(c) Coordination Number 4:

The most common complexes with this coordination number have square planar structure.
Example: [Cu(NH₃)₄]²⁺.

The shape of square planar complexes is flat.

Question

Write the geometrical shape of the following complexes:

1. [Zn(OH)₄]^2-
2. [Co(CN)₃(OH)₂]^2-
3. [Fe(H₂O)₂(CN)₄]^–

(C) CATALYTIC ACTION

With the exception of zinc, transition metals have catalytic action. They are used as catalysts in most chemical reactions. The catalytic power of catalysts is explained by their ability to exist in stable oxidation states. In their catalytic ion, transition metals act as electron carriers, transferring electrons from electron donors to electron acceptors.

Example: Iron and nickel are catalysts during hydrogenation of alkenes. Vanadium pentoxide (V₂O₅) is used as a catalyst during the contact process. Manganese (IV) oxide is used as a catalyst during preparation of oxygen. Iron is used as a catalyst during the Haber process.

(D) PARAMAGNETIC PROPERTY

Paramagnetic substances can be magnetized and attracted by a magnetic bar. For a substance to be paramagnetic, it must possess unpaired electrons in its electronic structure. Transition metals, with the exception of zinc, are paramagnetic because they have unpaired electrons in 3d orbitals. Unpaired electrons form weak magnetic fields and magnetic forces arise due to weak electric currents formed in the orbitals as a result of electron spin. When a magnetic bar is brought near a paramagnetic object, there is a tendency of domains or unpaired electrons to shift from the side of weak magnetic field (from the object) to the side of strong magnetic field (magnet), causing the whole object to be attracted towards the magnetic field.

When electrons in the orbitals are paired, magnetic fields cancel each other out as the magnetic field of two paired electrons move in exactly opposite directions (vector). Such substances cannot be magnetized and are not attracted by a magnetic bar. These substances are called diamagnetic substances.

Transition metals are paramagnetic in pure metal or when electrons are present in the d-electronic structure of the transition metallic ion. For example, ferrous sulfate [FeSO₄] is paramagnetic because Fe²⁺ has unpaired electrons in 3d orbitals.

This is because ligands of high spin, when approaching the d-orbitals, cause small separation of the d-orbitals. The energy of separation (ΔE) between double degenerate and triple degenerate orbitals is small. Electrons will spread according to Hund’s Rule, where each electron occupies its orbital singly before pairing. In that case, the central transition metallic ion remains with unpaired electrons in the electronic structure. Hence, the complex compound will be paramagnetic. Consider Fe²⁺ ion complexed with ligands of high spin like water [H₂O].

This complex [Fe(H₂O)₄]²⁺ is a 4sp³ hybrid complex and is paramagnetic since Fe²⁺ ion in the complex contains unpaired electrons in the d-electrons. If Fe²⁺ complexes with ligands of low spin, the paramagnetic property will be destroyed and 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 (ΔE) between the two degenerates becomes large. If the energy of separation is large, electrons fill the orbitals according to Aufbau’s Principle, pairing themselves in the triple degenerate orbitals which have lower energy. Consider Fe²⁺ complexes with ammonia, a ligand of low spin. This complex is 3d⁴ sp³ hybrid and is diamagnetic since there are no unpaired electrons in the electronic structure of Fe²⁺. All electrons in Fe²⁺ are paired in triple degenerate orbitals.

QUESTIONS

NECTA 2001 P2 Question 6c

Use the configuration of 3d orbital electrons on cobalt (III) ion to explain why [CoF₆]^3- is paramagnetic while [CoCN₆]^3- is not paramagnetic? [10%]

Solution

Consider the electronic structure of CO and CO³⁺ ions given below:

Then, complex [CoF₆]^3- is paramagnetic because fluoride ion is a ligand of high spin, hence causes small magnetic field splitting of 3d orbitals of cobalt (III) ion. The energy separation is small and electrons fill according to Hund’s Rule, thus cobalt (III) ion contains unpaired electrons in the 3d orbitals. Hence, it is a paramagnetic substance.

This is a 4sp³ hybrid complex and it is paramagnetic due to unpaired electrons in 3d orbitals.

But the complex [CoCN₆]^3- is not paramagnetic due to absence of unpaired electrons in 3d orbitals in cobalt (III) ion in the complex caused by complexing with ligands of low spin, the cyano ligand. This ligand exerts a strong magnetic field on the unpaired 3d orbitals, causing large separation which makes excitation difficult. Electrons fill according to Aufbau’s Principle.

Paramagnetic property of transition metals can be destroyed in two ways:

1. Temperature/oxidation process
2. Complex compound formation

a. TEMPERATURE/OXIDATION

Raising temperature destroys the magnetic property of the elements. It causes excitation of electrons and ionization of atoms by losing electrons. Ionization to an extent of losing all unpaired electrons destroys the magnetic property of the substance.

b. COMPLEX COMPOUND FORMATION

Formation of complex compounds by coordination with ligands may also destroy the magnetic property of transition metals. However, this depends on the oxidation state of the metals and the nature of the ligands involved in complex compound formation, whether they are ligands of high or low spin.

When the complex compound is formed by ligands of high spin such as CO₄^2-, OH^–, F^–, Br^–, I^–, Cl^–, the energy separation between triple and double degenerates is small. Electrons fill according to Hund’s Rule, resulting in unpaired electrons in the double and triple degenerate orbitals. The complex remains paramagnetic. Example: CoF₆^3- is paramagnetic. The F^– fills in 4s, 4p, and 4d vacant orbitals; the 3d remains with unpaired electrons resulting in paramagnetism.

Question

a. Predict the coordination number of Ni²⁺ and state whether the complex will be paramagnetic or diamagnetic if Ni²⁺ complexes with:

1. Bromine ions, Br^– (ligands of high spin)
2. Ammonia molecules NH₃ (ligands of low spin)

b. Fe³⁺ complexing with NH₃ of low spin.

C) VARIABILITY IN OXIDATION STATES

With the exception of zinc, transition metals form more than one stable oxidation state. They have variable oxidation states. Variability in oxidation state in transition metals is explained by the ability to ionize by losing electrons from both sub-energy levels 4s and 3d. The energy gap between 4s and 3d is very small. The difference between the two degenerate levels is so small that normal radiant energy from the sun is enough to excite electrons, and electrons from sub-energy levels 4s and 3d can be removed by almost the same amount of ionization energy. During ionization, transition metals give off 4s electrons first followed by 3d electrons. Transition metals depend on the number of electrons present in 4s orbital. Chromium has the lowest oxidation state of +2, losing two electrons in 4s orbital, and copper has the lowest oxidation state of +1. The maximum or highest oxidation state is 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 causes strong effective charge pull over the 3d electrons. The only oxidation state for elements found beyond manganese is +3, and for copper, the stable oxidation state is +2, which is attained after loss of a single 4s electron and another from sub-energy level 3d. The nuclear charge is high in zinc and hence the nuclear pull on 3d electrons 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 two electrons. Only one stable oxidation state for zinc is one of the factors which excludes zinc from transition metals.

NOTE:

For elements with atomic numbers 21 to 25 which can ionize to an extent of losing all the d electrons, the tendency is that increase in oxidation states accompanies 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 [Cr₂O₃] and chromium (II) oxide [CrO] are basic oxides which have no chemical reaction with other basic solutions, even alkaline solutions.

They react with acidic solutions.

Chromium (IV) oxide [CrO₃] is amphoteric. It dissolves in water and forms hydroxides which dissolve in both alkaline and acidic solutions.

Sodium chromite (III) is formed:

The last two oxides, chromium (V) oxide Cr₂O₅ and chromium (VI) oxide (CrO₃) are acidic oxides. They dissolve in water and form strong acidic solutions.

Chromic acid is so strong that it has no chemical reaction with other acidic solutions.

Chromic acid reacts with basic solutions and forms normal salts.

Reaction of the various oxides of chromium justifies the increase in acidic character.

NOTE

Change in base-acidic character in transition metals during ionization explains 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 their 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.

SUMMARY

• Low oxidation state is basic in nature (e.g., Mn²⁺).
• High oxidation state is acidic in nature (e.g., MnO₄^–).