The Nature of Chemical Bonds and Theories of Chemical Bonding

 

The Nature of Chemical Bonds and Theories of Chemical Bonding

The Nature of Chemical Bonds and Theories of Chemical Bonding

1. Introduction to Matter and Elements:
• Matter consists of distinct elements.
• Under normal conditions, elements exist as independent atoms except for noble gases.
2. Formation of Molecules:
• Atoms group together to form molecules.
• Molecules are collections of atoms with characteristic properties.
• A molecule is held together by a force known as a chemical bond.
3. The Concept of Chemical Bonds:
• Chemical bonds are attractive forces between atoms, ions, etc.
• Chemical bonds are responsible for holding together the constituents of different chemical species.
4. Questions about Chemical Bonding:
• The process of forming chemical compounds from combinations of atoms raises questions.
• Why do atoms combine? Why are specific combinations possible?
• What determines why certain atoms combine while others do not?
• Why do molecules have definite shapes?
5. Theories and Concepts of Chemical Bonding:
• Kössel-Lewis Approach:
• A theory explaining the transfer of electrons between atoms to achieve stable electron configurations.
• Focuses on achieving noble gas electron configurations through electron transfer.
• Valence Shell Electron Pair Repulsion (VSEPR) Theory:
• Describes molecular shapes based on the repulsion between valence electron pairs.
• Explains the three-dimensional arrangement of atoms in molecules.
• Valence Bond (VB) Theory:
• Explains chemical bonding in terms of overlapping atomic orbitals.
• Emphasizes the role of unpaired electrons in the formation of bonds.
• Molecular Orbital (MO) Theory:
• Describes chemical bonding using molecular orbitals formed by the combination of atomic orbitals.
• Electrons are treated as wave-like entities, leading to a more comprehensive understanding of bonding.
6. Relation to Atomic Structure and Periodic Table:
• Development of bonding theories closely linked to advancements in atomic structure understanding.
• Electronic configurations of elements and their placement in the periodic table influence bonding behavior.
7. Stability and Bonding:
• Systems tend to achieve stability.
• Bonding is a natural process that reduces the energy of a system, leading to increased stability.

In summary, the nature of chemical bonds and the theories of chemical bonding have evolved over time to address fundamental questions about atomic combinations, molecule formation, shapes, and stability. The theories, including Kössel-Lewis, VSEPR, Valence Bond, and Molecular Orbital theories, provide different perspectives on the forces that hold atoms together in molecules, considering electron configurations, atomic orbitals, and molecular shapes. These theories have been influenced by our understanding of atomic structure and the periodic table, ultimately explaining why atoms combine, the possible combinations, and the shapes of molecules. Bonding serves as nature's mechanism to achieve stability within systems.

 

Kössel-Lewis Approach to Chemical Bonding

1. Introduction:
• In 1916, independent efforts by Kössel and Lewis brought about a satisfactory explanation for chemical bonding based on electron interactions.
• Their approach was grounded in understanding valence and drew inspiration from the inert properties of noble gases.
2. Lewis's Model of the Atom:
• Lewis conceptualized atoms as having a "Kernel" comprising the nucleus and inner electrons, surrounded by an outer shell.
• The outer shell could accommodate up to eight electrons, distributed at the corners of a cube that enveloped the "Kernel."
• This arrangement formed a stable octet of electrons, promoting stability in the atom.
3. Stable Octet and Chemical Bonds:
• Lewis proposed that atoms attain stability by achieving the octet electron configuration.
• For atoms to achieve this stable state, they form chemical bonds.
• In the example of sodium (Na) and chlorine (Cl), sodium donates an electron to chlorine, resulting in the formation of Na+ and Cl– ions.
• The transfer of an electron from one atom to another leads to the formation of ionic bonds.
4. Covalent Bonds and Octet Rule:
• In molecules like Cl₂, H₂, F₂, etc., atoms form bonds by sharing pairs of electrons.
• This sharing of electrons enables each atom to complete its outer shell and attain the stable octet configuration.
• The concept of the octet rule underlines the tendency of atoms to seek eight electrons in their outer shell, mirroring noble gas electron configurations.
5. Lewis Symbols:
• In molecular formation, only the outer shell electrons (valence electrons) partake in chemical bonding.
• Inner shell electrons are shielded and usually remain uninvolved in bonding.
• G.N. Lewis, an American chemist, introduced a simplified notation to represent valence electrons in atoms, known as Lewis symbols.

Advantages and Insights of Kössel-Lewis Approach:

• Explanation of Valence: Kössel and Lewis provided a logical explanation for valence based on the stable octet configuration.
• Inertness of Noble Gases: The concept of achieving noble gas-like electron configurations explained why noble gases are chemically inert.
• Ionic and Covalent Bonds: The approach distinguished between ionic and covalent bonds, elucidating how electron transfer and sharing contribute to bond formation.
• Octet Rule: The octet rule became a guiding principle for understanding chemical behavior and predicting molecular stability.

In Summary: The Kössel-Lewis approach to chemical bonding, developed by Kössel and Lewis, introduced a model of atoms based on the outer shell's electron arrangement. This approach elucidated the significance of achieving a stable octet configuration for atoms through chemical bonding. The concept of Lewis symbols simplified the representation of valence electrons. The approach's insights into ionic and covalent bonding, the octet rule, and the inertness of noble gases contributed significantly to the understanding of chemical interactions and molecular stability.

 

Significance of Lewis Symbols and Kössel's Contributions to Chemical Bonding

1. Lewis Symbols and Valence Electrons:
• Lewis symbols represent an element's valence electrons as dots around its chemical symbol.
• The number of dots in a Lewis symbol corresponds to the number of valence electrons.
• This representation aids in calculating the common or group valence of the element.
2. Group Valence Calculation:
• Group valence of an element is often either equal to the number of dots in its Lewis symbol or 8 minus the number of dots (valence electrons).
3. Kössel's Observations in Chemical Bonding:
• Kössel's contributions to chemical bonding shed light on critical observations:
• The periodic table separates highly electronegative halogens and highly electropositive alkali metals with noble gases in between.
• Formation of negative ions from halogen atoms and positive ions from alkali metal atoms involves electron loss and gain, respectively.
• Negative and positive ions formed acquire stable noble gas electronic configurations, especially the octet (eight electrons) in the outer shell (ns²np⁶).
• The stability of noble gas configurations implies that ions aim to achieve similar electron arrangements.
4. Stabilization through Electrostatic Attraction:
• The interaction between positive and negative ions formed due to electron transfer results in electrostatic attraction.
• This type of bond was termed the "electrovalent bond."
• Electrovalence corresponds to the number of unit charges on the ion. For instance, calcium carries a positive electrovalence of two, while chlorine bears a negative electrovalence of one.
5. Implications and Applications of Kössel's Postulations:
• Kössel's ideas laid the groundwork for modern concepts related to ion formation through electron transfer and the creation of ionic crystalline compounds.
• These insights contributed significantly to understanding and systematizing ionic compounds.
• While Kössel's concepts were valuable, they also recognized that certain compounds deviated from these ideas.

Key Contributions and Insights:

• Electronegativity and Electropositivity: Kössel highlighted the contrasting properties of halogens and alkali metals, separated by noble gases in the periodic table.
• Ionic Formation: Kössel's observations explained the electron gain and loss during ion formation, leading to stable noble gas-like configurations.
• Stable Outer Shell: Kössel emphasized that negative and positive ions achieve stability by acquiring noble gas outer shell electron configurations.
• Electrovalent Bonds: The electrostatic attraction between oppositely charged ions was termed as electrovalent bonding, with electrovalence representing ion charge.
• Ionic Compounds Understanding: Kössel's ideas significantly enhanced the comprehension and organization of ionic compounds.

In Summary: Kössel's contributions revolutionized the understanding of chemical bonding, emphasizing the connection between electron transfer and ion formation. His insights into electronegativity, electron gain and loss, stable electron configurations, and electrovalent bonding laid the foundation for modern concepts. While his ideas provided a profound understanding of ionic compounds, they also acknowledged that certain compounds didn't conform to these concepts, reflecting the evolving nature of chemical understanding.

 

The Octet Rule in Chemical Bonding

1. Introduction:
• In 1916, Kössel and Lewis introduced the electronic theory of chemical bonding.
• This theory explains how atoms combine through the transfer or sharing of valence electrons.
2. Fundamental Idea - Octet Rule:
• Atoms can attain stability by having a complete outer electron shell.
• The octet rule states that atoms tend to combine in a way that allows them to achieve a stable configuration of eight electrons in their valence shell, resembling the noble gases' electronic configuration.
3. Two Modes of Combination:
• Electron Transfer (Ionic Bonding):
• Involves the transfer of valence electrons from one atom to another.
• One atom gains electrons to fill its valence shell, becoming negatively charged (anion), while the other loses electrons and becomes positively charged (cation).
• The electrostatic attraction between oppositely charged ions leads to the formation of ionic compounds.
• Electron Sharing (Covalent Bonding):
• Atoms share pairs of electrons to complete their valence shells.
• By sharing electrons, each atom achieves a stable configuration similar to noble gases.
• Covalent bonding is common in molecular compounds and forms when atoms have similar electronegativities.
4. Significance of the Octet Rule:
• The octet rule guides the formation of various chemical compounds by dictating how atoms will interact to achieve stability.
• It explains why atoms either gain, lose, or share electrons during bonding.
5. Stability and Noble Gas Configuration:
• Noble gases possess a stable electron configuration with eight electrons in their valence shell (except helium, which has two).
• Other elements aim to emulate this stable state by following the octet rule during chemical bonding.
6. Predictive Power:
• The octet rule aids in predicting the types of bonds that will form between different elements.
• It also provides insights into the properties of resulting compounds.
7. Limitations and Exceptions:
• While the octet rule is a useful guideline, it doesn't explain every type of chemical bonding.
• Some molecules and compounds don't strictly adhere to the octet rule, especially for elements with d or f orbitals that can accommodate more than eight electrons.

Implications of the Octet Rule:

• Ionic and Covalent Bond Types: The octet rule underlies the fundamental distinction between ionic and covalent bonds based on electron transfer and sharing, respectively.
• Stability and Noble Gas Mimicry: The pursuit of the octet configuration drives atoms towards enhanced stability by resembling the noble gas electron arrangements.
• Predictive Tool: The octet rule is an invaluable predictive tool, aiding in explaining the behavior of various elements during bonding.
• Beyond Octet Rule: While essential, the octet rule has limitations, especially for elements with more complex electron configurations.

In Summary: The octet rule, developed by Kössel and Lewis, forms the cornerstone of chemical bonding theories. It elucidates how atoms combine by either transferring or sharing valence electrons to achieve a stable configuration with eight electrons in their valence shell. While it provides a strong basis for understanding most bonding scenarios, exceptions exist due to the intricate electron arrangements of certain elements.

Covalent Bonding and Lewis-Langmuir Theory

1. Introduction:
• Langmuir (1919) built upon Lewis's ideas, refining the concept of chemical bonding.
• Langmuir introduced the term "covalent bond" and expanded on Lewis's octet rule.
2. Evolution from Lewis to Lewis-Langmuir Theory:
• Lewis's idea of atoms bonding through electron sharing was enhanced by Langmuir.
• Langmuir abandoned the concept of a fixed cubic arrangement of the octet and introduced the term "covalent bond."
3. Chlorine Molecule Example - Cl2:
• The Cl atom has the electronic configuration [Ne]3s²3p⁵, lacking one electron for the argon configuration.
• Cl2 formation involves the sharing of a pair of electrons between two chlorine atoms.
• Each chlorine atom contributes one electron to the shared pair.
• The result is that both chlorine atoms achieve the outer shell octet of argon, leading to a stable configuration.
4. Representation of Covalent Bonds - Lewis Dot Structures:
• Lewis dot structures use dots to represent electrons in atoms and molecules.
• These structures provide visual insights into the arrangement of shared electrons.
• Lewis dot structures are applicable to different molecules with identical or different combining atoms.
5. Key Conditions of Covalent Bonding:
• Electron Pair Sharing: Covalent bonds form by sharing an electron pair between atoms.
• Contribution of Electrons: Each atom involved contributes at least one electron to the shared pair.
• Achievement of Stable Configuration: Combining atoms achieve noble gas configurations due to shared electrons.
6. Examples of Covalent Bonds:
• Water (H₂O) and Carbon Tetrachloride (CCl₄): In these molecules, atoms share electron pairs to form covalent bonds, fulfilling the conditions of electron sharing and noble gas configurations.
7. Multiple Bonds:
• Single Covalent Bond: Formed when two atoms share one electron pair.
• Double Bond: Two pairs of electrons are shared between atoms. Example: Carbon dioxide (CO₂).
• Triple Bond: Three pairs of electrons are shared between atoms. Examples: Nitrogen gas (N₂), ethyne (C₂H₂).

Implications of the Lewis-Langmuir Theory:

• Advancement of Covalent Bond Concept: Langmuir's covalent bond concept refined Lewis's ideas, focusing on electron sharing and arrangement.
• Visualization through Lewis Dot Structures: Lewis dot structures provide a visual representation of electron distribution in molecules.
• Universal Application: The Lewis-Langmuir theory's principles apply to various compounds, guiding our understanding of covalent bonding.
• Explanation of Multiple Bonds: The theory explains the formation of single, double, and triple bonds through shared electron pairs.

In Summary: Langmuir's extension of Lewis's ideas gave rise to the Lewis-Langmuir theory of covalent bonding. This theory emphasizes electron sharing between atoms to achieve stable electron configurations, visualized through Lewis dot structures. Examples like the chlorine molecule showcase how covalent bonds lead to noble gas-like outer shell configurations. Additionally, the theory accounts for multiple bonds, elucidating the formation of double and triple bonds through shared electron pairs.

 

 

Oxidation States in d-Block Elements

 

Point-wise Explanation of Oxidation States in d-Block Elements:

1. Oxidation State Variation in d-Block Elements:
• The d-block elements, also known as transition metals, exhibit various oxidation states.
• The central idea is that elements near the middle of the d-block tend to display the greatest number of oxidation states.
2. Manganese's Example:
• Manganese (Mn) serves as an example; it displays oxidation states from +2 to +7.
• This wide range of oxidation states occurs due to the presence of a suitable number of d electrons for both gaining and losing electrons.
3. Factors Influencing Oxidation States:
• Oxidation states at the extremes of the d-block series are limited by specific electron configurations and orbital availability.
4. Early Series Elements (Scandium and Titanium):
• Scandium (Sc) and Titanium (Ti) are early elements in the d-block series.
• Scandium (II) oxidation state is uncommon due to having too few electrons to effectively share or lose.
• Titanium (IV) is more stable than Ti(III) or Ti(II) because it loses its 4s electrons before 3d electrons, resulting in a more stable configuration.
5. Late Series Elements (Copper and Zinc):
• Copper (Cu) and Zinc (Zn) are later elements in the d-block series.
• Copper has a limited range of oxidation states (I and II) due to having too many d electrons, which results in fewer available orbitals for sharing electrons.
• Zinc's only stable oxidation state is +2, as it loses only its 4s electrons, and its d orbitals remain uninvolved.
6. Maximum Stable Oxidation States:
• The most stable oxidation states typically correspond to the sum of s and d electrons up to manganese.
7. Examples of Maximum Stable Oxidation States:
• Titanium (IV) oxide: Ti(IV)O2
• Vanadium (V) oxide cation: V(V)O2+
• Chromium (VI) oxide anion: Cr(VI)O4(2–)
• Manganese (VII) oxide anion: Mn(VII)O(4–)
8. Stability Trend Beyond Manganese:
• After manganese, there's an abrupt decrease in the stability of higher oxidation states.
• This phenomenon arises due to increased repulsion between multiple electrons in the d orbitals, making higher oxidation states less favorable.
9. Common Oxidation States for Later Elements:
• Iron (Fe) exists in oxidation states II and III.
• Cobalt (Co) can be found in oxidation states II and III.
• Nickel (Ni) is commonly seen in the +2 oxidation state.
• Copper (Cu) can be found in oxidation states I and II.
• Zinc (Zn) primarily occurs in the +2 oxidation state due to its full d orbitals.

Phenomenon Explanation:

The phenomenon of varying oxidation states in d-block elements can be attributed to electron configuration and orbital availability. Elements in the middle of the d-block have an optimal balance of d electrons, allowing for a wider range of oxidation states. As you move toward the extremes of the series, the electron configuration leads to limitations in the number of stable oxidation states. Elements with too few or too many d electrons find it challenging to achieve various oxidation states due to electron repulsion and the availability of orbitals for sharing electrons.

This pattern helps to understand why certain oxidation states are common for specific elements and provides insights into the chemical behavior of transition metals.

 

Point-wise Explanation of Oxidation State Variability in Transition Elements:

1. Characteristic of Transition Elements:
• Transition elements exhibit a unique feature known as the variability of oxidation states.
• This variability is a consequence of the incomplete filling of d orbitals in their electron configurations.
2. Incomplete Filling of d Orbitals:
• Transition elements have partially filled d orbitals in their electron configurations.
• The arrangement of electrons in these orbitals allows for a range of possible oxidation states.
3. Differing Oxidation States by Unity:
• The variability of oxidation states in transition elements is distinct in that the states differ from each other by a unit of one.
• For example, a transition element like vanadium (V) can exhibit oxidation states such as +2, +3, +4, and +5.
4. Example of Vanadium (V):
• Vanadium serves as an illustrative example.
• Its oxidation states include V(II), V(III), V(IV), and V(V), differing from each other by a unit of one.
5. Contrast with Non-Transition Elements:
• The variability of oxidation states in transition elements contrasts with that of non-transition elements.
• In non-transition elements, oxidation states often differ by a unit of two.

Phenomenon Explanation:

The variability of oxidation states in transition elements arises from the specific arrangement of electrons in their electron configurations. Transition elements have partially filled d orbitals, which provide a flexible environment for electron interactions and transfers. This results in a range of possible oxidation states, each differing from the others by a single unit. For instance, in the case of vanadium (V), the incomplete filling of its d orbitals allows it to exhibit oxidation states such as +2, +3, +4, and +5, which differ from each other by a unit of one.

On the other hand, non-transition elements typically have completely filled s and p orbitals in their valence shells. This arrangement leads to a different pattern of oxidation state variability. In non-transition elements, oxidation states often differ by a unit of two due to the filling of these orbitals and the way electrons are gained or lost during chemical reactions.

In summary, the unique electron configuration of transition elements, particularly the incomplete filling of d orbitals, enables them to display a wider range of oxidation states, with the states differing by a single unit. This distinct behavior contributes to the versatile chemistry exhibited by transition metals.

 

Point-wise Explanation of Oxidation State Variability in d-Block Elements within Groups:

1. Distinct Oxidation State Behavior in Groups 4 to 10:
• Within the d-block elements (groups 4 to 10), a unique trend is observed regarding oxidation state variability.
2. Contrasting Oxidation State Preference in p-Block and d-Block:
• In the p-block elements, heavier members tend to favor lower oxidation states due to the inert pair effect.
• However, among the d-block elements, the situation is opposite, with heavier members favoring higher oxidation states.
3. Example in Group 6:
• Group 6 d-block elements illustrate this trend.
• Molybdenum (Mo) and tungsten (W) in their +6 oxidation states (Mo(VI) and W(VI)) are more stable than chromium (Cr) in its +6 oxidation state (Cr(VI)).
4. Consequences of Stability Difference:
• Due to the stability difference, dichromate (Cr(VI)) in acidic medium acts as a potent oxidizing agent, while molybdenum trioxide (MoO₃) and tungsten trioxide (WO₃) do not exhibit the same oxidizing capability.
5. Influence of Ligand Properties on Oxidation States:
• Oxidation states can also be influenced by the nature of ligands in complex compounds.
6. Role of Ligand p-Acceptor Character:
• Low oxidation states are favored when complex compounds have ligands with p-acceptor character in addition to s-bonding.
• These ligands can participate in electron acceptance through their p-orbitals.
7. Example with Nickel Carbonyl (Ni(CO)₄) and Iron Pentacarbonyl (Fe(CO)₅):
• In the complex compounds Ni(CO)₄ and Fe(CO)₅, nickel and iron are found in a zero oxidation state.
• This is due to the combined effect of ligand's electron donation through s-bonding and electron acceptance through p-acceptor character.

Phenomenon Explanation:

The trend of oxidation state variability in the d-block elements within groups is a consequence of the interplay between the electronic structures of the elements and the properties of the ligands they interact with.

In the p-block, the inert pair effect causes heavier elements to prefer lower oxidation states, but in the d-block, the opposite tendency is observed. This can be attributed to the relative energies of different electron orbitals in the d-block elements and their ability to gain or lose electrons.

Furthermore, the role of ligand properties is crucial. Ligands with p-acceptor character can accept electrons from the metal center, stabilizing lower oxidation states. In complex compounds like Ni(CO)₄ and Fe(CO)₅, the combined effects of s-bonding and p-acceptor interactions lead to the presence of the metal in a zero oxidation state.

This phenomenon illustrates the complex nature of transition metal chemistry, where both the electronic structure of the elements and the interactions with ligands contribute to the observed oxidation state patterns.

 

d-BLOCK AND f-BLOCK ELEMENTS:

 

d-BLOCK AND f-BLOCK ELEMENTS:

 

CLASS XII

 

Transition Metals:

• Located in the d-block of the periodic table (groups 3-12).
• D-orbitals progressively filled in each of the four long periods.
• Chemical properties are transitional between s and p-block elements.
• Transition metals are defined by IUPAC as metals with incomplete d subshells in neutral atoms or ions.
• Notable series: 3d (Sc to Zn), 4d (Y to Cd), 5d (Hf to Hg), and 6d (Ac to Cn).
• Zinc, cadmium, and mercury (group 12) are not considered transition metals due to full d10 configuration.

Inner Transition Metals:

• Located in the f-block of the periodic table.
• 4f series: Lanthanoids (Ce to Lu).
• 5f series: Actinoids (Th to Lr).

Characteristics:

• Presence of partly filled d or f orbitals sets them apart from non-transition elements.
• Compounds of transition elements studied separately due to unique characteristics.
• Valence theory applicable to both transition and non-transition elements.

Metals and Their Significance:

• Precious metals (e.g., silver, gold, platinum) and industrially important metals (e.g., iron, copper, titanium) belong to transition metals series.

Focus of the Unit:

• Electronic configuration, occurrence, and general characteristics of transition elements.
• Emphasis on trends in properties of the first row (3d) transition metals.
• Preparation and properties of significant compounds.
• Discussion of electronic configurations, oxidation states, and chemical reactivity of inner transition metals.

 CLICK TO REGISTER

CLICK FOR TUTORVERSE WEBSITE

CLICK TO CONTACT

Top of Form

 

1. Position on the Periodic Table:
• The d-block is located in the middle of the periodic table, between the s-block and p-block elements.
2. Orbitals Involved:
• The d-block elements have electrons filling the d-orbitals.
• These d-orbitals belong to the penultimate (second-to-last) energy level of the atoms.
3. Rows of Transition Metals:
• The d-block is divided into four rows of transition metals: 3d, 4d, 5d, and 6d.
• Each row corresponds to a specific energy level (n = 3, 4, 5, and 6).
4. Electronic Configuration Pattern:
• The outermost electronic configuration of these elements follows a consistent pattern.
• It can be represented as (n-1)d^1–10ns^1–2, where 'n' represents the energy level of the element.
5. Exceptions to the Configuration:
• Palladium (Pd) is an exception to this general configuration pattern.
• Its electronic configuration is 4d¹⁰ 5s⁰, indicating that it has completely filled its 4d orbitals but lacks electrons in its 5s orbital.

 

1. Inner d Orbitals and Outer ns Orbitals:
• In the (n–1)d^1–10ns^1–2 electronic configuration pattern:
• (n–1) represents the inner d orbitals, which can accommodate one to ten electrons.
• ns represents the outermost orbital, which can hold one or two electrons.
2. Exceptions to the Generalization:
• The (n–1)d and ns orbitals have similar energy levels, leading to exceptions in the electronic configurations due to small energy differences.
• Half-filled and completely filled sets of orbitals are relatively more stable.
3. Cr and Cu Configurations in the 3d Series:
• Chromium (Cr) and copper (Cu) demonstrate the consequences of small energy differences:
• Cr's configuration is 3d⁵ 4s¹, rather than the expected 3d⁴ 4s².
• Cu's configuration is 3d¹⁰ 4s¹, instead of 3d⁹ 4s².
• The energy gap between the 3d and 4s orbitals prevents electrons from entering the 3d orbitals, creating a more stable configuration.
4. Electronic Configurations of Zn, Cd, Hg, and Cn:
• Zn, Cd, Hg, and Cn follow the general formula (n-1)d¹⁰ns² for their outer orbitals.
• These elements have completely filled orbitals in both their ground state and common oxidation states.
• Consequently, they are not considered as transition elements due to their stable configurations.

 

1. Enhanced Protrusion of d Orbitals:
• Transition elements have d orbitals that extend further from the nucleus compared to s and p orbitals.
• This positioning makes the d orbitals more exposed to the surrounding environment.
• The influence of the surroundings on d orbitals is substantial, affecting both the atoms or molecules containing these elements and the elements themselves.
2. Shared Properties of dn Configurations:
• Transition metal ions with dⁿ configurations (n = 1 – 9) often exhibit similar magnetic and electronic characteristics.
• Similarities arise due to the presence of partly filled d orbitals, contributing to their distinctive behavior.
3. Distinctive Characteristics of Partly Filled d Orbitals:
• Elements with partly filled d orbitals showcase unique properties:
• Display a range of oxidation states due to the versatility of d orbital electrons.
• Formation of colored ions due to electronic transitions within d orbitals.
• Formation of complex compounds with various ligands due to d orbitals' availability for bonding.
• Enter into catalytic reactions, highlighting their catalytic property.
• Exhibit paramagnetic behavior, where unpaired d electrons lead to magnetic attraction.
4. Group Similarities and Horizontal Row Trends:
• Within a horizontal row of transition elements, there are strong resemblances in properties.
• Horizontal rows, particularly the 3d row, demonstrate general trends in properties.
• These trends provide insights into how properties evolve as atomic number increases.
5. Discussion Sequence:
• The unit's content follows a specific order:
• Detailed exploration of the general characteristics of transition elements and their trends in horizontal rows, particularly focusing on the 3d row.
• Examination of group similarities among these elements.

Electronic Configurations of Zn, Cd, Hg, and Cn:

1. The outer orbital electronic configurations of Zn, Cd, Hg, and Cn follow the general formula      (n-1)d ¹⁰ ns ².
2. In this configuration, the (n-1)d orbitals contain 10 electrons, while the ns orbital holds 2 electrons.
3. These elements exhibit complete filling of their outer orbitals in both the ground state and common oxidation states.

Transition Element Status:

1. Due to their completely filled outer orbitals, Zn, Cd, Hg, and Cn are not classified as transition elements.
2. Transition elements typically possess incompletely filled d orbitals, which contributes to their distinctive properties.

Distinguishing d Orbitals in Transition Elements:

1. In transition elements, d orbitals extend more towards the periphery of an atom compared to s and p orbitals.
2. The positioning of d orbitals makes them more susceptible to external influences and allows them to impact surrounding atoms and molecules.

Shared Properties of dn Configurations:

1. Ions with a given dn configuration (n = 1 – 9) exhibit similar electronic and magnetic properties.
2. This similarity stems from the presence of partly filled d orbitals, which contribute to consistent behavior.

Distinctive Traits of Partly Filled d Orbitals:

1. Elements with partially filled d orbitals demonstrate distinct characteristics:
• They can display a variety of oxidation states due to the flexibility of d orbital electrons.
• Formation of colored ions arises from electronic transitions within the d orbitals.
• Complex formation with diverse ligands is possible due to availability of d orbitals for bonding.
• They often engage in catalytic reactions and show paramagnetic behavior due to unpaired d electrons.

Organized Sequence of Discussion:

1. The unit's content follows a specific order:
• Explanation of the electronic configurations of Zn, Cd, Hg, and Cn, which are distinct from typical transition elements.
• Clarification on the non-transition element status of these elements.
• Exploration of the protruding d orbitals in transition elements and their influence on surroundings.
• Discussion on the shared properties of ions with dn configurations.
• Elaboration on the special traits exhibited by elements with partly filled d orbitals.

 CLICK TO REGISTER

CLICK FOR TUTORVERSE WEBSITE

CLICK TO CONTACT

 

Catalytic Property and Paramagnetic Behavior:

1. Transition metals and their compounds showcase the ability to catalyze chemical reactions.
2. They also exhibit paramagnetic behavior due to the presence of unpaired electrons.
3. These properties will be discussed further in detail later within this unit.

Comparative Characteristics of Transition Elements:

1. Transition elements within a horizontal row share greater similarities in properties compared to non-transition elements.
2. The collective behavior of transition elements within the same row indicates recurring trends.

Focus on Horizontal Row and 3d Series:

1. The study begins by examining general characteristics and trends of transition elements within horizontal rows.
2. Particular emphasis is given to the 3d row, investigating how properties change as the atomic number increases.

Exploration of Group Similarities:

1. Group similarities among transition elements will also be explored.
2. While horizontal row behaviors are prominent, there exist noteworthy commonalities within specific groups.

Sequence of Study:

1. The approach follows a specific sequence:
• In-depth analysis of the catalytic property and paramagnetic behavior of transition metals and their compounds.
• Emphasis on the comparative properties of transition elements within horizontal rows versus non-transition elements.
• Concentration on the general characteristics and trends of transition elements within the 3d series of the periodic table.
• Discussion of common characteristics shared among transition elements within certain groups.

PHYSICAL PROPERTIES OF TRANSITION METALS:

1. Typical Metallic Properties:
• Transition elements exhibit characteristic metallic properties such as high tensile strength, ductility, malleability, high thermal and electrical conductivity, and a metallic luster.
• However, Zn, Cd, Hg, and Mn deviate from these properties.
2. Crystal Structures:
• Most transition metals have one or more common metallic crystal structures at normal temperatures, except for Zn, Cd, and Hg.
3. Hardness and Volatility:
• Transition metals (excluding Zn, Cd, and Hg) are known for their hardness and low volatility.
4. Melting and Boiling Points:
• Transition metals generally have high melting and boiling points.
• The graph (Fig. 4.1) shows the melting points of transition metals from 3d, 4d, and 5d series.
• High melting points are attributed to increased involvement of both (n-1)d and ns electrons in interatomic metallic bonding.
• In each series, melting points rise to a maximum at d5 electron configuration, except for anomalies seen in Mn and Tc.
• Melting points tend to decrease as the atomic number within a series increases.
5. Enthalpy of Atomization:
• Transition metals have high enthalpies of atomization, as depicted in Fig. 4.2.
• Peaks at the midpoint of each series indicate that having one unpaired electron per d orbital greatly favors strong interatomic interactions.
• Generally, a higher number of valence electrons lead to stronger resulting bonding.
6. Effect on Electrode Potential:
• Enthalpy of atomization significantly influences the standard electrode potential of a metal.
• Metals with very high enthalpy of atomization (high boiling points) tend to exhibit noble behavior in reactions.
• (Note: More information on electrode potentials is provided later.)

 

1. 
   Enthalpy of Atomization Definition:
• Enthalpy of atomization refers to the energy required to completely separate one mole of a solid metal into its individual atoms in the gaseous state at a given temperature.
2. Transition Metals and Enthalpies:
• Transition metals, located in the d-block of the periodic table, exhibit high enthalpies of atomization.
• This is due to the strong metallic bonding resulting from the presence of unpaired electrons in the d orbitals of these metals.
3. Unpaired Electrons and Interatomic Interaction:
• Transition metals often have unpaired electrons in their d orbitals, leading to strong interatomic interactions.
• The presence of unpaired electrons allows for effective sharing of electrons among neighboring atoms, promoting stability in the metallic lattice.
4. Formation of Metallic Bonds:
• The enthalpy of atomization is influenced by the strength of metallic bonds formed between atoms in the solid metal.
• Transition metals' unpaired d electrons contribute significantly to these strong metallic bonds.
5. Peak Enthalpies and Electron Configurations:
• Peaks in the graph of enthalpies of atomization (as shown in Fig. 4.2) often occur around elements with configurations like d5 in the middle of each series (3d, 4d, 5d).
• Having one unpaired electron per d orbital enhances interatomic interactions, resulting in higher enthalpies of atomization.
6. Valence Electron Count and Bonding Strength:
• The number of valence electrons plays a crucial role in determining the strength of the resultant bonding.
• Transition metals generally have a higher number of valence electrons compared to main-group elements, leading to stronger bonding and higher enthalpies of atomization.
7. Relationship to Standard Electrode Potential:
• Enthalpy of atomization significantly affects the standard electrode potential of a metal.
• Metals with high enthalpies of atomization (and hence strong metallic bonds) tend to exhibit noble behavior in reactions and have higher standard electrode potentials.
8. Variation Across Periods and Series:
• Within a series of transition elements (3d, 4d, 5d), enthalpies of atomization show trends based on electron configurations and atomic sizes.
• Anomalies might occur, such as lower values for certain elements (e.g., Mn and Tc), due to particular electronic configurations.

In summary, the enthalpies of atomization of d-block metals are influenced by the presence of unpaired electrons, strong interatomic interactions, valence electron count, and electron configurations. These factors contribute to the high stability and metallic bonding observed in transition metals, leading to their characteristic high enthalpies of atomization.

Enthalpies of Atomization of d-Block Metals and Periodic Trend:

1. Definition of Enthalpy of Atomization:
• Enthalpy of atomization is the energy required to break one mole of a solid metal into individual gas-phase atoms under standard conditions.
2. Strong Metallic Bonding:
• Transition metals in the d-block tend to have high enthalpies of atomization due to their strong metallic bonding.
• Metallic bonding arises from the sharing of delocalized electrons in the metal lattice.
3. Unpaired Electrons and Bonding Strength:
• Transition metals often have unpaired electrons in their d orbitals, leading to stronger interatomic interactions.
• These unpaired electrons contribute to the formation of stronger metallic bonds, resulting in higher enthalpies of atomization.
4. Trends Across Periods (Rows):
• Moving across a period, from left to right, enthalpies of atomization generally increase for transition metals.
• This trend is due to the increasing effective nuclear charge, which enhances attraction between the nucleus and valence electrons, leading to stronger bonding.
5. Trends Down Groups (Columns):
• Going down a group, from top to bottom, enthalpies of atomization usually increase for transition metals.
• This trend is attributed to the larger atomic size and increased electron shielding, which decrease the effective nuclear charge felt by valence electrons. As a result, these electrons are less tightly held, allowing for stronger metallic bonding.
6. Exceptions and Anomalies:
• There are exceptions to these trends due to variations in electron configurations and electron filling patterns.
• Anomalies, such as lower enthalpies for certain elements like Mn and Tc, can be attributed to their specific electronic structures.
7. Effect on Reactivity and Stability:
• Transition metals with higher enthalpies of atomization tend to be more stable and less reactive.
• These metals are less likely to lose their atoms or electrons in reactions due to the strong bonding and the energy required to break these bonds.
8. Correlation with Standard Electrode Potential:
• Enthalpy of atomization affects the standard electrode potential of a metal, influencing its reactivity in redox reactions.
• Metals with higher enthalpies of atomization often exhibit nobler behavior and more positive electrode potentials.
9. Applications and Properties:
• The high enthalpies of atomization contribute to the characteristic properties of transition metals, such as their high melting and boiling points and strong mechanical properties.
10. Overall Periodic Trend:
• The periodic trend of increasing enthalpies of atomization across periods and down groups reflects the gradual strengthening of metallic bonds as atomic size, effective nuclear charge, and electron shielding change in these directions.

 CLICK TO REGISTER

CLICK FOR TUTORVERSE WEBSITE

CLICK TO CONTACT

In summary, the enthalpies of atomization of d-block metals follow a periodic trend with increasing values across periods and down groups. This trend is a result of the interplay between electron configurations, atomic sizes, and effective nuclear charges, all of which influence the strength of metallic bonding and the energy required to break these bonds.

Variation in Atomic and Ionic Sizes of Transition Metals:

Variation in Atomic and Ionic Sizes of Transition Metals:

1. Progressive Decrease in Radius within a Series:

• Ions of the same charge within a series exhibit a consistent trend of decreasing radius as atomic number increases.
• This decrease is due to the addition of electrons to d orbitals as the nuclear charge increases.
• D electrons provide less effective shielding, causing a stronger electrostatic attraction between the nucleus and outermost electrons.
• This effect leads to a decrease in ionic radius within a series.

2. Similar Variation in Atomic Radii:

• Similar trends are observed in atomic radii within a series, though the variation is relatively small.
• The same phenomenon that causes ionic radii decrease also influences atomic radii.

3. Comparison Across Different Series:

• Comparison of atomic sizes between different series reveals interesting patterns.
• The second (4d) series shows an increase in atomic size compared to the first (3d) series.
• However, the third (5d) series appears to have virtually the same atomic radii as the corresponding elements in the second series.

4. Lanthanoid Contraction Explanation:

• The similarity in atomic radii between the second and third series is due to the "lanthanoid contraction."
• Before the 5d series begins, the 4f orbitals must be filled.
• Filling 4f before 5d leads to a regular decrease in atomic radii, compensating for the expected increase.
• Lanthanoid contraction arises from imperfect shielding of one electron by another within the same set of orbitals.

5. Factors Behind Lanthanoid Contraction:

• Similar to other transition series, imperfect shielding of electrons in orbitals contributes to lanthanoid contraction.
• Shielding of 4f electrons by each other is less effective than for d electrons.
• As the series progresses and nuclear charge increases, the entire 4fn orbitals experience a consistent decrease in size.

6. Consequences of Lanthanoid Contraction:

• The decrease in size of 4fn orbitals, coupled with an increase in atomic mass, results in a general increase in element density.
• Notable density increase can be observed from titanium (Z = 22) to copper (Z = 29).

7. Unusual Similarity in Properties:

• The second and third d series exhibit remarkably similar atomic radii and, consequently, physical and chemical properties.
• This similarity is more significant than what would be expected based on the usual family relationships in the periodic table.

By breaking down the provided text into bullet points, we've highlighted the main concepts and explanations regarding the variation in atomic and ionic sizes of transition metals and the phenomenon of lanthanoid contraction.

 

Ionisation Enthalpies:

1. Increase in Ionisation Enthalpy along Transition Series:

• Transition elements exhibit an increase in ionisation enthalpy from left to right within a series.
• This increase is due to the rising nuclear charge as inner d orbitals get filled.

2. Gradual Increase in Successive Ionisation Enthalpies:

• Compared to non-transition elements, the transition elements' successive ionisation enthalpies show a gentler increase.
• The first three ionisation enthalpies for the first series of transition elements are listed in Table 4.2.
• The increase in the first ionisation enthalpy is moderate, while the second and third ionisation enthalpies increase significantly for successive elements.

3. Lesser Variation in Transition Elements vs. Non-Transition Elements:

• The variation in ionisation enthalpy along a transition series is less pronounced than in a non-transition element period.
• This makes the transition elements' ionisation enthalpy changes smoother and more gradual.

4. Influence of Electron Configuration on Ionisation Enthalpy:

• In the 3d series, as we move along the period from scandium to zinc, the nuclear charge increases.
• Electrons are added to the inner subshell's 3d orbitals, which partially shield the outer 4s electrons from the growing nuclear charge.
• This effective shielding by 3d electrons results in a less rapid decrease in atomic radii and only a slight increase in ionisation energies.

5. Exceptions to the Increasing Ionisation Enthalpy Trend:

• The formation of Mn²⁺ and Fe³⁺ ions shows a deviation from the usual trend.
• Mn²⁺ has a d⁵ configuration and Fe³⁺ has a d⁵ configuration, both causing lower ionisation enthalpies due to exchange energy effects.
• Similar deviations occur at corresponding elements in later transition series.

6. Exchange Energy and Ionisation Enthalpy:

• Ionisation enthalpy is influenced by three terms: electron-nucleus attraction, electron-electron repulsion, and exchange energy.
• Exchange energy stabilizes electron configuration by favoring maximum parallel spins in degenerate orbitals.
• The loss of exchange energy enhances stability and makes ionisation more challenging.

7. Influence of Electronic Configuration on Ionisation Enthalpy:

• The ionisation enthalpy of Mn⁺ is lower than Cr⁺ due to their different electron configurations (Mn⁺: 3d⁵4s¹, Cr⁺: d⁵).
• Similarly, Fe²⁺ has a lower ionisation enthalpy than Mn²⁺ due to their configurations (Fe²⁺: d⁶, Mn²⁺: 3d⁵).

8. Formation of M2+ Ions and Ionisation Enthalpy:

• To form M²⁺ ions from gaseous atoms, the sum of the first and second ionisation enthalpies, along with atomisation enthalpy, is needed.
• Second ionisation enthalpy dominates this process and has particularly high values for Cr and Cu, with d⁵ and d¹⁰ configurations respectively.
• Zn's second ionisation enthalpy is lower due to the removal of one 4s electron, resulting in a stable d¹⁰ configuration.

9. Third Ionisation Enthalpies:

• The third ionisation enthalpies reveal the difficulty of removing electrons from d⁵ (Mn²⁺) and d¹⁰ (Zn²⁺) ions.
• These enthalpies are generally high and indicate why oxidation states greater than two are challenging for elements like copper, nickel, and zinc.

10. Complexity of Oxidation State Stabilities:

• Ionisation enthalpies offer some insights into relative oxidation state stabilities, but this matter is intricate and not easily generalized.

By presenting the information in pointwise format, we've highlighted the key concepts and explanations surrounding ionisation enthalpies of transition elements.

Oxidation States

Transition elements, also known as transition metals, are a group of elements located in the middle of the periodic table. One of the distinctive features of transition elements is their ability to exhibit a wide variety of oxidation states or valence states. This characteristic arises from the presence of partly filled d orbitals in their electronic configurations. Here's a detailed breakdown of the points related to the variety of oxidation states in transition elements:

1. Partially Filled d Orbitals: Transition elements have partially filled d orbitals in their valence electron configuration. These d orbitals can accommodate a varying number of electrons, allowing the elements to exhibit multiple oxidation states. The number of unpaired electrons in these d orbitals determines the possible oxidation states an element can achieve.
2. Oxidation States: Oxidation state refers to the charge an atom would have if all its bonds were purely ionic. Transition elements can have multiple oxidation states, often differing by a unit of one, due to the availability of multiple d orbitals to lose or gain electrons from.
3. Inner and Outer Orbitals: In transition elements, the outer s orbitals and the inner d orbitals are close in energy. This proximity allows for easy electron exchange between these orbitals, leading to the formation of various oxidation states.
4. Variation Across the Periodic Table: The variety of oxidation states varies across the transition metal series. Elements in the early transition metals (Scandium to Chromium) generally have fewer oxidation states compared to those in the middle and late transition metals (Iron to Copper and beyond). This is due to the availability of different electron shells and d orbitals.
5. Common Oxidation States: While transition elements can potentially exhibit a wide range of oxidation states, there are often certain common oxidation states that they tend to favor. For example, +2 and +3 are common oxidation states for many transition elements. Additionally, some elements like Chromium and Manganese are known to exhibit multiple oxidation states, such as +2, +3, +4, +6 for Chromium and +2, +4, +7 for Manganese.
6. Factors Influencing Oxidation States: The oxidation state that a transition element can achieve is influenced by factors such as the atomic size, nuclear charge, shielding effect, and electron configuration. These factors collectively affect the ease with which electrons can be gained or lost from different orbitals.
7. Formation of Complex Ions: Transition elements are highly capable of forming complex ions by coordinating with ligands (molecules or ions with lone pairs of electrons). This results in the stabilization of certain oxidation states that might not be attainable in simple ionic compounds.
8. Catalytic Activity: The ability of transition metals to exhibit various oxidation states is closely related to their catalytic activity. Many transition metals serve as catalysts in chemical reactions by changing their oxidation state during the reaction.
9. Color and Magnetic Properties: The variety of oxidation states contributes to the vibrant colors and magnetic properties exhibited by transition metal compounds. The absorption of specific wavelengths of light is due to electronic transitions between different oxidation states, leading to colored compounds.
10. Importance in Biological Systems: Transition metals with multiple oxidation states play crucial roles in biological systems. For instance, iron can exist in both +2 and +3 oxidation states and is essential for oxygen transport in hemoglobin and electron transfer in enzymes.

In conclusion, the variety of oxidation states exhibited by transition elements is a result of their unique electronic configurations, which provide multiple accessible d orbitals for electron exchange. This versatility in oxidation states contributes to their diverse chemical, physical, and biological properties.

 

 

Variation in Atomic and Ionic Sizes of Transition Metals:   In general, ions of the same charge in a given series show progressive decrease in radius with increasing atomic number. This is because the new electron enters a d orbital each time the nuclear charge increases by unity. It may be recalled that the shielding effect of a d electron is not that effective, hence the net electrostatic attraction between the nuclear charge and the outermost electron increases and the ionic radius decreases. The same trend is observed in the atomic radii of a given series. However, the variation within a series is quite small. An interesting point emerges when atomic sizes of one series are compared with those of the corresponding elements in the other series. The curves in Fig. 4.3 show an increase from the first (3d) to the second (4d) series of the elements but the radii of the third (5d) series are virtually the same as those of the corresponding members of the second series. This phenomenon is associated with the intervention of the 4f orbitals which must be filled before the 5d series of elements begin. The filling of 4f before 5d orbital results in a regular decrease in atomic radii called Lanthanoid contraction which essentially compensates for the expected increase in atomic size with increasing atomic number. The net result of the lanthanoid contraction is that the second and the third d series exhibit similar radii (e.g., Zr 160 pm, Hf 159 pm) and have very similar physical and chemical properties much more than that expected on the basis of usual family relationship. The factor responsible for the lanthanoid contraction is somewhat similar to that observed in an ordinary transition series and is attributed to similar cause, i.e., the imperfect shielding of one electron by another in the same set of orbitals. However, the shielding of one 4f electron by another is less than that of one d electron by another, and as the nuclear charge increases along the series, there is fairly regular decrease in the size of the entire 4fn orbitals. The decrease in metallic radius coupled with increase in atomic mass results in a general increase in the density of these elements. Thus, from titanium (Z = 22) to copper (Z = 29) the significant increase in the density may be noted.

 CLICK TO REGISTER

CLICK FOR TUTORVERSE WEBSITE

CLICK TO CONTACT

 

 

THE TRANSITION ELEMENTS (d-BLOCK):   The d–block occupies the large middle section of the periodic table flanked between s– and p– blocks in the periodic table. The d–orbitals of the penultimate energy level of atoms receive electrons giving rise to four rows of the transition metals, i.e., 3d, 4d, 5d and 6d. In general the electronic configuration of outer orbitals of these elementsis (n-1)d1– 10ns1–2 except for Pd where its electronic configuration is 4d105s0. The (n–1) stands for the inner d orbitals which may have one to ten electrons and the outermost ns orbital may have one or two electrons. However, this generalisation has several exceptions because of very little energy difference between (n-1)d and ns orbitals. Furthermore, half and completely filled sets of orbitals are relatively more stable. A consequence of this factor is reflected in the electronic configurations of Cr and Cu in the 3d series. For example, consider the case of Cr, which has 3d5 4s1 configuration instead of 3d44s2; the energy gap between the two sets (3d and 4s) of orbitals is small enough to prevent electron entering the 3d orbitals. Similarly in case of Cu, the configuration is 3d104s1 and not 3d94s2.  The electronic configurations of outer orbitals of Zn, Cd, Hg and Cn are represented by the general formula (n-1)d10ns2. The orbitals in these elements are completely filled in the ground state as well as in their common oxidation states. Therefore, they are not regarded as transition elements. The d orbitals of the transition elements protrude to the periphery of an atom more than the other orbitals (i.e., s and p), hence, they are more influenced by the surroundings as well as affect the atoms or molecules surrounding them. In some respects, ions of a given dn configuration (n = 1 – 9) have similar magnetic and electronic properties. With partly filled d orbitals these elements exhibit certain characteristic properties such as display of a variety of oxidation states, formation of coloured ions and entering into complex formation with a variety of ligands. The transition metals and their compounds also exhibit catalytic property and paramagnetic behaviour. All these characteristics have been discussed in detail later in this Unit. There are greater similarities in the properties of the transition elements of a horizontal row in contrast to the non-transition elements. However, some group similarities also exist. We shall first study the general characteristics and their trends in the horizontal rows (particularly 3d row) and then consider some group similarities. The electronic configurations of outer orbitals of Zn, Cd, Hg and Cn are represented by the general formula (n-1)d 10 ns 2 . The orbitals in these elements are completely filled in the ground state as well as in their common oxidation states. Therefore, they are not regarded as transition elements. The d orbitals of the transition elements protrude to the periphery of an atom more than the other orbitals (i.e., s and p), hence, they are more influenced by the surroundings as well as affect the atoms or molecules surrounding them. In some respects, ions of a given d n configuration (n = 1 – 9) have similar magnetic and electronic properties. With partly filled d orbitals these elements exhibit certain characteristic properties such as display of a variety of oxidation states, formation of coloured ions and entering into complex formation with a variety of ligands. The transition metals and their compounds also exhibit catalytic property and paramagnetic behaviour. All these characteristics have been discussed in detail later in this Unit. There are greater similarities in the properties of the transition elements of a horizontal row in contrast to the non-transition elements. However, some group similarities also exist. We shall first study the general characteristics and their trends in the horizontal rows (particularly 3d row) and then consider some group similarities.

 

Ionisation Enthalpies There is an increase in ionisation enthalpy along each series of the transition elements from left to right due to an increase in nuclear charge which accompanies the filling of the inner d orbitals. Table 4.2 gives the values of the first three ionisation enthalpies of the first series of transition elements. These values show that the successive enthalpies of these elements do not increase as steeply as in the case of non-transition elements. The variation in ionisation enthalpy along a series of transition elements is much less in comparison to the variation along a period of non-transition elements. The first ionisation enthalpy, in general, increases, but the magnitude of the increase in the second and third ionisation enthalpies for the successive elements, is much higher along a series. The irregular trend in the first ionisation enthalpy of the metals of 3d series, though of little chemical significance, can be accounted for by considering that the removal of one electron alters the relative energies of 4s and 3d orbitals. You have learnt that when d-block elements form ions, ns electrons are lost before (n – 1) d electrons. As we move along the period in 3d series, we see that nuclear charge increases from scandium to zinc but electrons are added to the orbital of inner subshell, i.e., 3d orbitals. These 3d electrons shield the 4s electrons from the increasing nuclear charge somewhat more effectively than the outer shell electrons can shield one another. Therefore, the atomic radii decrease less rapidly. Thus, ionization energies increase only slightly along the 3d series. The doubly or more highly charged ions have dn configurations with no 4s electrons. A general trend of increasing values of second ionisation enthalpy is expected as the effective nuclear charge increases because one d electron does not shield another electron from the influence of nuclear charge because d-orbitals differ in direction. However, the trend of steady increase in second and third ionisation enthalpy breaks for the formation of Mn2+ and Fe3+ respectively. In both the cases, ions have d 5 configuration. Similar breaks occur at corresponding elements in the later transition series. The interpretation of variation in ionisation enthalpy for an electronic configuration dn is as follows: The three terms responsible for the value of ionisation enthalpy are attraction of each electron towards nucleus, repulsion between the electrons and the exchange energy. Exchange energy is responsible for the stabilisation of energy state. Exchange energy is approximately proportional to the total number of possible pairs of parallel spins in the degenerate orbitals. When several electrons occupy a set of degenerate orbitals, the lowest energy state corresponds to the maximum possible extent of single occupation of orbital and parallel spins (Hunds rule). The loss of exchange energy increases the stability. As the stability increases, the ionisation becomes more difficult. There is no loss of exchange energy at d 6 configuration. Mn+ has 3d54s1 configuration and configuration of Cr+ is d5 , therefore, ionisation enthalpy of Mn+ is lower than Cr+ . In the same way, Fe2+ has d 6 configuration and Mn2+ has 3d5 configuration. Hence, ionisation enthalpy of Fe2+ is lower than the Mn2+ . In other words, we can say that the third ionisation enthalpy of Fe is lower than that of Mn. The lowest common oxidation state of these metals is +2. To form the M2+ ions from the gaseous atoms, the sum of the first and second ionisation enthalpy is required in addition to the enthalpy of atomisation. The dominant term is the second ionisation enthalpy which shows unusually high values for Cr and Cu where M+ ions have the d5 and d10 configurations respectively. The value for Zn is correspondingly low as the ionisation causes the removal of one 4s electron which results in the formation of stable d10 configuration. The trend in the third ionisation enthalpies is not complicated by the 4s orbital factor and shows the greater difficulty of removing an electron from the d5 (Mn2+) and d10 (Zn2+) ions. In general, the third ionisation enthalpies are quite high. Also the high values for third ionisation enthalpies of copper, nickel and zinc indicate why it is difficult to obtain oxidation state greater than two for these elements. Although ionisation enthalpies give some guidance concerning the relative stabilities of oxidation states, this problem is very complex and not amenable to ready generalisation.