What Are Transition Elements? Properties, Periodic Table Location, and Atomisation Enthalpy

What Are Transition Elements?

Transition elements, also known as transition metals, are a group of elements located in the d-block of the periodic table, specifically in Groups 3 to 12. These elements are characterized by the presence of partially filled d-orbitals, either in their neutral atomic state or in at least one of their stable oxidation states.

IUPAC Definition 

A transition element, according to IUPAC, is an element that either:

• Has a partially filled d-subshell in its atom,
  or
  
• Can form stable positive ions (cations) with an incompletely filled d-orbital.

In simpler terms, even if the d-orbitals are full in the neutral atom, the element is still considered a transition metal if it forms ions where the d-orbitals are not completely filled.

This feature allows transition elements to show a wide range of chemical behavior, such as:

• Multiple oxidation states
  
• Ability to form colored compounds
  
• Magnetic properties
  
• Catalytic activity

These special characteristics come from the behavior of electrons in the d-orbitals.

Position of Transition Elements on the Periodic Table

Transition elements occupy the central region of the modern periodic table, within the d-block.

• Groups: 3 to 12
  
• Periods: 4, 5, and 6
  
• Block: d-block (main transition metals)
  

In addition, the f-block elements, which include the lanthanides and actinides, are sometimes referred to as inner transition elements. These elements have partially filled f-orbitals rather than d-orbitals, and are usually considered separately from the main transition series.

Also Check- The Periodic Table of Elements: What It Is and How It Works

Inner Transition Elements

Inner transition elements are a subset of transition elements found in the f-block of the periodic table. They are located below the main table and are divided into two series-

• Lanthanides (atomic numbers 57–71)
  
• Actinides (atomic numbers 89–103)
  

These elements are characterized by the gradual filling of the 4f (lanthanides) and 5f (actinides) orbitals.

Key Characteristics-

• Partially filled f-orbitals in their atoms or common oxidation states
  
• Exhibit variable oxidation states, though less varied than d-block elements
  
• Form colored ions, though colors are generally less intense than those of d-block metals
  
• Show paramagnetism, often due to unpaired f-electrons
  
• Tend to form complexes with ligands, especially oxygen- and nitrogen-donors
  
• Many actinides are radioactive, with elements beyond uranium (U) being synthetic
  

Lanthanides (4f-series)-

• From Lanthanum (La, Z=57) to Lutetium (Lu, Z=71)
  
• Known for lanthanide contraction– gradual decrease in atomic and ionic size across the series
  
• Common oxidation state- +3
  
• Used in- magnets, phosphors, lasers, and hybrid vehicles
  

Actinides (5f-series)-

• From Actinium (Ac, Z=89) to Lawrencium (Lr, Z=103)
  
• Exhibit more variable oxidation states (e.g., U- +3 to +6)
  
• High radioactivity; many are short-lived
  
• Used in- nuclear reactors (e.g., U, Pu), radiopharmaceuticals
  

Comparison with d-Block Transition Elements-

Property	d-Block (Transition)	f-Block (Inner Transition)
Orbitals filled	(n−1)d	(n−2)f
Series	3d, 4d, 5d	4f (lanthanides), 5f (actinides)
Common oxidation states	Variable (e.g., +2 to +7)	+3 dominant in lanthanides; variable in actinides
Color intensity	Often vivid	Less intense
Complex formation	Extensive	Moderate (more covalent in actinides)
Magnetic behavior	From unpaired d-electrons	From unpaired f-electrons

Examples of d-Block Transition Elements

Element	Symbol	Group	Period
Scandium	Sc	3	4
Titanium	Ti	4	4
Iron	Fe	8	4
Copper	Cu	11	4

Zinc (Zn), along with cadmium (Cd) and mercury (Hg), is located in the d-block but is not classified as a transition element. This is because their atoms and common ions (e.g., Zn²⁺, Cd²⁺, Hg²⁺) have completely filled d-subshells, typically with the configuration (n−1)d¹⁰. They do not form ions with partially filled d-orbitals and therefore do not meet the IUPAC definition of a transition element.

Also Check – What Are Chemical Elements? A Complete Guide for Students

Electronic Configuration of Transition Elements

Transition elements have unique electron arrangements that involve filling the (n−1)d subshell and the ns subshell. This electron configuration influences many of their chemical and physical properties, such as variable oxidation states, colored compounds, and catalytic activity.

General Electron Configuration

The general electron configuration for transition elements is written as:

(n−1)d¹–¹⁰ ns¹–²

This means:

• Electrons are added to the d-orbital of the second outermost shell (n−1).
  
• The outermost s-orbital (ns) is also being filled with one or two electrons.
  

For example, if n = 4 (as in the 4th period), the configuration involves filling the 3d and 4s orbitals.

List of Transition Elements and Their Electron Configurations

Below is a table of selected transition elements along with their atomic number, element name, and electronic configuration:

Element Name	Symbol	Atomic Number	Electron Configuration
Scandium	Sc	21	[Ar] 3d¹ 4s²
Titanium	Ti	22	[Ar] 3d² 4s²
Vanadium	V	23	[Ar] 3d³ 4s²
Chromium	Cr	24	[Ar] 3d⁵ 4s¹ (exception)
Manganese	Mn	25	[Ar] 3d⁵ 4s²
Iron	Fe	26	[Ar] 3d⁶ 4s²
Cobalt	Co	27	[Ar] 3d⁷ 4s²
Nickel	Ni	28	[Ar] 3d⁸ 4s²
Copper	Cu	29	[Ar] 3d¹⁰ 4s¹ (exception)
Zinc	Zn	30	[Ar] 3d¹⁰ 4s² (not a true transition element)
Yttrium	Y	39	[Kr] 4d¹ 5s²
Zirconium	Zr	40	[Kr] 4d² 5s²
Niobium	Nb	41	[Kr] 4d⁴ 5s¹ (exception)
Molybdenum	Mo	42	[Kr] 4d⁵ 5s¹ (exception)
Technetium	Tc	43	[Kr] 4d⁵ 5s²
Ruthenium	Ru	44	[Kr] 4d⁷ 5s¹
Rhodium	Rh	45	[Kr] 4d⁸ 5s¹
Palladium	Pd	46	[Kr] 4d¹⁰ (fully filled d-shell, no 5s electron)
Silver	Ag	47	[Kr] 4d¹⁰ 5s¹ (exception)
Cadmium	Cd	48	[Kr] 4d¹⁰ 5s² (not a true transition element)

Key Observations and Exceptions

• Most transition elements follow the general pattern of filling the (n−1)d orbitals before or along with the ns orbitals.
  
• However, some elements such as chromium, copper, molybdenum, niobium, and silver do not follow the standard Aufbau principle strictly.
  

Examples of Exceptions:

• Chromium (Cr):
  Expected: [Ar] 3d⁴ 4s²
  Actual: [Ar] 3d⁵ 4s¹
  Reason: A half-filled d-subshell (d⁵) provides extra stability.
  
• Copper (Cu):
  Expected: [Ar] 3d⁹ 4s²
  Actual: [Ar] 3d¹⁰ 4s¹
  Reason: A fully filled d-subshell (d¹⁰) is more stable.
  
• Palladium (Pd):
  Has the unusual configuration: [Kr] 4d¹⁰ (no 5s electrons). This is rare and specific to its electron interactions.
  

These configurations are well-established experimentally and occur because the energy difference between the ns and (n−1)d orbitals is very small, allowing electrons to shift to achieve a more stable arrangement (half-filled or fully filled d-subshell).

Important Note on Zinc, Cadmium, and Mercury

Although Zn, Cd, and Hg are in the d-block, they are not considered true transition elements because:

• They do not have partially filled d-orbitals in either their neutral atoms or common ions.
  
• Their d-subshells are completely filled: d¹⁰ configuration.
  

Atomic and Ionic Radii of Transition Elements

The atomic and ionic radii of transition elements show distinct trends across periods and down groups in the periodic table.

Across a Period (Left to Right):

• From Group 3 to Group 6, the atomic and ionic radii decrease.
  
• This happens because the nuclear charge increases, pulling the electrons closer to the nucleus.
  
• However, the shielding effect of d-electrons is relatively poor, so the size decreases more noticeably.
  

Between Groups 7 to 10:

• The radii remain fairly constant.
  
• This is due to a balance between increasing nuclear charge and increased electron–electron repulsion within the d-orbitals.
  

Groups 11 and 12:

• Elements like Cu, Zn and Cd show a slight increase in radius.
  
• This is because the repulsion between fully filled d-electrons slightly cancels out the nuclear attraction.

Down a Group:

• As we move down a group, the atomic and ionic radii increase.
  
• This is due to the addition of new electron shells (energy levels).
  
• Even though the nuclear charge increases, the outer electrons are farther from the nucleus, leading to a larger atomic size.

Ionization Enthalpy of Transition Elements

Definition:

Ionization enthalpy (or ionization energy) is the energy required to remove the outermost electron from a neutral atom in the gaseous state.

Transition Metal Trends:

• The ionization enthalpy of transition elements is generally higher than that of s-block elements.
  
• This is because transition metals have:
  
  • Higher nuclear charge
    
  • Smaller atomic radii
    
  • Stronger attraction between nucleus and valence electrons
    

Across a Period:

• As we move from left to right across a transition series:
  
  • Ionization enthalpy generally increases.
    
  • This trend is not as regular as in the s- and p-blocks because of irregularities in electron configurations and subshell repulsions.
    

Relation to Atomic Radius:

• Atoms with smaller radii (more nuclear attraction) tend to have higher ionization energies.
  
• This explains why early transition metals (like Sc, Ti) have lower ionization enthalpies than later ones (like Fe, Cu).

Why These Properties Matter:

• Atomic radius affects bonding length and metallic properties.
  
• Ionization enthalpy influences chemical reactivity, oxidation states, and compound stability.

What Are Transition Metals Properties?

Transition metals, located in the d-block of the periodic table (Groups 3–12), exhibit a unique set of chemical and physical properties due to the presence of partially filled d-orbitals. These properties distinguish them from s- and p-block elements.

1. Variable Oxidation States

Transition metals can exhibit more than one oxidation state because of the relatively small energy difference between their (n−1)d and ns orbitals. This allows them to lose different numbers of electrons under different chemical conditions.

• Example: Iron (Fe) exists as both Fe²⁺ and Fe³⁺.
  
• Example: Manganese (Mn) shows oxidation states ranging from +2 to +7.
  

This versatility enables their widespread use in redox reactions and catalysis.

2. Formation of Coloured Compounds

Many transition metal compounds are distinctively colored, which arises from electronic transitions between d-orbitals of different energy levels in the presence of ligands (coordination environments). These are called d–d transitions.

• The specific color depends on:
  
  • The metal ion
    
  • Its oxidation state
    
  • The nature of the surrounding ligands
    
• Example: [Cu(H2O)6]2+[Cu(H₂O)_6]^{2+}[Cu(H2​O)6​]2+ appears blue due to absorption in the red-orange region.
  

This property is utilized in analytical chemistry, pigments, and biological systems (e.g., hemoglobin coloration due to Fe²⁺).

3. Catalytic Properties

Transition metals and their compounds frequently act as catalysts due to their ability to:

• Change oxidation states reversibly
  
• Form temporary complexes with reactants
  
• Provide alternative reaction pathways with lower activation energy
  
• Example: Iron (Fe) is used in the Haber process for ammonia synthesis.
  
• Example: Vanadium(V) oxide (V₂O₅) is used in the Contact process for sulfuric acid production.
  

These properties are especially important in industrial and environmental chemistry.

4. Magnetic Properties

Transition elements often exhibit magnetism due to unpaired d-electrons.

• Paramagnetism: Caused by unpaired electrons; these metals are weakly attracted to a magnetic field (e.g., Mn²⁺, Fe³⁺).
  
• Ferromagnetism: Found in elements like Fe, Co, and Ni, where unpaired electrons align in parallel, resulting in strong magnetic behavior.
  

This makes them essential in electronics, motors, and data storage technologies.

5. Complex Formation (Coordination Compounds)

Transition metals form a wide variety of stable complex ions by accepting electron pairs from ligands into their empty d-orbitals.

• Common ligands include: H₂O, NH₃, Cl⁻, CN⁻
  
• These compounds can have different geometries: octahedral, tetrahedral, square planar
  
• Example: [Fe(CN)6]3−[Fe(CN)_6]^{3−}[Fe(CN)6​]3−, [Ni(NH3)6]2+[Ni(NH₃)_6]^{2+}[Ni(NH3​)6​]2+
  

Complex formation is fundamental in biochemistry, coordination chemistry, and metal extraction processes.

6. High Melting and Boiling Points

Transition metals exhibit high thermal stability due to:

• Strong metallic bonding, which arises from the delocalization of electrons from both ns and (n−1)d orbitals
  
• The presence of multiple unpaired electrons, allowing for stronger bonding interactions between atoms
  

This makes them suitable for structural materials and high-temperature applications.

7. High Density and Hardness

Transition metals are generally:

• Hard
  
• Dense
  
• Mechanically strong
  

This is because their atoms are closely packed, and metallic bonding is stronger due to partially filled d-orbitals. These features are useful in construction materials, cutting tools, and alloy manufacture.

8. Good Electrical Conductivity

Transition metals conduct electricity efficiently due to the mobility of delocalized electrons in their metallic structure.

The overlapping of orbitals in the metallic lattice allows free movement of electrons, making them ideal for electrical wiring and components (e.g., copper wires, silver contacts).

Why Transition Elements Have High Enthalpy of Atomisation

Definition:

Enthalpy of atomisation is the amount of energy required to break all bonds in one mole of a metallic element to form separate gaseous atoms.

Explanation:

Transition elements have high enthalpies of atomisation because:

1. Strong Metallic Bonding:
   
   • Metallic bonds involve delocalized electrons from both s- and d-orbitals.
     
   • The more delocalized electrons involved, the stronger the metallic bond.
     
2. Presence of Unpaired d-Electrons:
   
   • Unpaired electrons contribute to interatomic attractions in the metallic lattice.
     
   • More unpaired d-electrons → more bonding interactions → higher energy needed to break them.
     
3. Close Atomic Packing:
   
   • Transition metals typically form close-packed structures, increasing bond density.

The enthalpy of atomisation peaks near the center of the transition series, where the number of unpaired d-electrons is highest (e.g., Cr, Fe).

Element	Enthalpy of Atomisation (kJ/mol)
Iron (Fe)	~416
Chromium (Cr)	~397
Manganese (Mn)	~281

This property is important in understanding melting points, bonding strength, and metallic behavior of the d-block elements.

FAQ Questions:

Q1: What are transition elements in simple terms?
A1: Transition elements are d-block metals with partially filled d-orbitals, known for forming colored compounds, variable oxidation states, and acting as catalysts.

Q2: Where are transition metals located on the periodic table?
A2: They are found in Groups 3 to 12, in periods 4, 5, and 6, in the central d-block of the periodic table.

Q3: Why do transition metals have variable oxidation states?
A3: Transition metals have similar energies in their ns and (n−1)d orbitals, allowing them to lose different numbers of electrons easily.

Q4: What makes transition metal compounds colored?
A4: Their ions absorb specific wavelengths of light due to d-d electron transitions, resulting in vivid colors.

Q5: Why do transition elements have high enthalpy of atomisation?
A5: Because of strong metallic bonding involving unpaired d-electrons, more energy is needed to separate atoms into the gaseous phase.