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.

 

PERIODIC CLASSIFICATION: PART 3

 

p-Block Elements

• The p-Block Elements encompass elements found in Group 13 to 18 on the periodic table.
• When combined with the s-Block Elements, they are referred to as Representative or Main Group Elements.
• The outermost electron arrangement ranges from ns²np¹ to ns²np⁶ within each period.
• At the conclusion of each period lies a noble gas element with a closed valence shell ns²np⁶ configuration.
• Noble gases possess fully populated valence shell orbitals, making their electron arrangements stable and resistant to change.
• Due to this stability, noble gases exhibit minimal chemical reactivity.
• Just before the noble gas series, two significant groups of nonmetals can be found.
• These groups are the halogens (Group 17) and the chalcogens (Group 16).
• Both the halogens and chalcogens have notably negative electron gain enthalpies.
• They readily accept one or two electrons respectively to achieve the stable noble gas configuration.
• Non-metallic tendencies intensify as we progress from left to right across a period.
• Conversely, metallic properties increase as we move down the group.

d-Block Elements

• These elements occupy Group 3 to 12 at the center of the Periodic Table.
• They are identified by their occupation of inner d orbitals by electrons and are termed d-Block Elements.
• These elements typically have the general outer electronic configuration of (n-1)d^1-10ns^0-2, with an exception being Pd which has the electronic configuration 4d¹⁰5s⁰.
• All of these elements are categorized as metals.
• They commonly yield ions with distinct colors, display variable valence (oxidation states), and exhibit paramagnetism.
• These elements often serve as catalysts.
• Notably, Zn, Cd, and Hg possess the electronic configuration (n-1)d¹⁰ns², and consequently, they lack several properties exhibited by typical transition elements.
• Transition metals act as a connection between the more chemically active metals found in the s-block and the comparatively less reactive elements in Groups 13 and 14.

f-Block Elements

• The f-Block Elements, also known as Inner-Transition Elements, encompass two rows at the bottom of the Periodic Table.
• These rows consist of Lanthanoids (Ce(Z = 58) – Lu(Z = 71)) and Actinoids (Th(Z = 90) – Lr (Z = 103)).
• Their outer electronic configuration is defined as (n-2)f^1-14 (n-1)d^0–1 ns².
• The last added electron for each element occupies an f-orbital.
• The Inner-Transition Elements are collectively named due to this arrangement.
• All elements in this category are metals.
• Elements within each series share similar properties.
• Among the Actinoids, chemistry is more intricate due to the multitude of possible oxidation states.
• Actinoid elements are radioactive in nature.
• Many actinoids are produced only in extremely small quantities through nuclear reactions, limiting their chemical study.
• Elements beyond uranium are termed Transuranium Elements.

Metals, Non-metals, and Metalloids:

• Apart from the categorization into s-, p-, d-, and f-blocks, another classification of elements is based on their properties.
• This classification divides elements into Metals and Non-Metals.
• Metals constitute over 78% of known elements and are positioned on the left side of the Periodic Table.
• Typically, metals are solids at room temperature, except for mercury. Gallium and caesium also have notably low melting points (303K and 302K, respectively).
• Metals generally possess high melting and boiling points.
• They excel in conducting heat and electricity.
• Metals display malleability (capable of being flattened into thin sheets) and ductility (can be drawn into wires).
• On the other hand, non-metals are found on the upper right-hand side of the Periodic Table.
• Within a horizontal row, elements transition from metallic on the left to non-metallic on the right.
• Non-metals are primarily solids or gases at room temperature, with exceptions like boron and carbon.
• They have low melting and boiling points.
• Non-metals are poor conductors of heat and electricity.
• Most non-metallic solids are brittle and lack malleability and ductility.
• The metallic nature increases as we descend a group, while non-metallic characteristics intensify from left to right across the Periodic Table.
• The shift from metallic to non-metallic traits is gradual, as denoted by the thick zig-zag line.
• Elements such as silicon, germanium, arsenic, antimony, and tellurium situated along this line and forming a diagonal across the Periodic Table exhibit features that resemble both metals and non-metals.
• These elements are referred to as Semi-metals or Metalloids.

 

Relationship between Ionization Enthalpy and Atomic Radius:

• Ionization enthalpy and atomic radius are interconnected properties.
• Understanding their trends involves considering electron attraction to the nucleus and electron-electron repulsion.
• Two main factors influencing these trends are: (i) electron attraction toward the nucleus, and (ii) electron-electron repulsion.

Effective Nuclear Charge and Shielding:

• Valence electrons experience an effective nuclear charge due to shielding by inner core electrons.
• Shielding or screening reduces the net positive charge experienced by valence electrons.
• Shielding is more effective when inner shell orbitals are fully filled.
• Alkali metals exhibit effective shielding with a single outermost ns-electron following a noble gas electronic configuration.

Periodic Trend - Across a Period:

• Moving from lithium to fluorine across the second period, successive electrons enter the same principal quantum level.
• Shielding doesn't increase significantly to offset the stronger attraction between electrons and nucleus.
• Increasing nuclear charge dominates over shielding, leading to tighter hold on outermost electrons.
• Ionization enthalpy increases across a period due to stronger electron-nucleus attraction.

Group Trend - Down a Group:

• Moving down a group, outermost electrons are farther from the nucleus.
• Increased shielding by inner electrons outweighs the rising nuclear charge.
• Outermost electron removal requires less energy down a group.

Ionization Enthalpy Anomalies:

• First ionization enthalpy of boron (Z = 5) is slightly less than beryllium (Z = 4) due to electron configuration differences.
• Beryllium's ionization removes an s-electron, while boron's removes a p-electron.
• Penetration of 2s-electron is higher than 2p-electron, leading to greater shielding for boron's p-electron.
• Oxygen's smaller first ionization enthalpy compared to nitrogen's arises from electron configuration differences.
• In oxygen, electron-electron repulsion increases due to pairing in 2p-orbitals, making removal of the fourth 2p-electron easier.

Electron Gain Enthalpy:

• Electron Gain Enthalpy (∆_egH) measures the enthalpy change when a neutral gaseous atom gains an electron to form a negative ion (anion).
• This process is represented by the equation: X(g) + e– → X–(g).

Exothermic and Endothermic Processes:

• The addition of an electron to an atom can be either exothermic (energy released) or endothermic (energy absorbed), depending on the element.
• Elements like halogens (group 17) release energy when gaining an electron due to reaching stable noble gas configurations.
• Noble gases, however, have positive electron gain enthalpies as the added electron enters a higher principal quantum level, resulting in an unstable electronic configuration.

Trends Across the Periodic Table:

• Electron gain enthalpy is more negative towards the upper right of the periodic table before the noble gases.
• Generally, electron gain enthalpy becomes more negative as atomic number increases across a period.
• Increasing effective nuclear charge across a period makes it easier to add an electron to smaller atoms due to stronger attraction to the nucleus.

Trends Down a Group:

• Electron gain enthalpy becomes less negative as you move down a group.
• The larger atomic size results in the added electron being farther from the nucleus.

Anomalies for Oxygen and Fluorine:

• Electron gain enthalpy of oxygen (O) and fluorine (F) is less negative than that of the following element in the period.
• Adding an electron to O or F places it in the smaller n = 2 quantum level, causing significant repulsion from other electrons present.
• For Sulfur (S) or Chlorine (Cl) in the n = 3 quantum level, the added electron occupies more space, leading to less electron-electron repulsion.

 

 

Electronegativity:

• Electronegativity is a qualitative measure of an atom's ability in a chemical compound to attract shared electrons towards itself.
• It's not a directly measurable quantity but can be estimated using various numerical scales.
• Notable scales include Pauling scale, Mulliken-Jaffe scale, and Allred-Rochow scale, with Pauling scale being the most widely used.
• Linus Pauling assigned an arbitrary value of 4.0 to fluorine as a reference point for electronegativity.

Variation in Electronegativity:

• Electronegativity varies based on the element it's bound to.
• It offers insights into the nature of the bonding force between atoms.

Trends in Electronegativity:

• Across a period (left to right), electronegativity generally increases (e.g., lithium to fluorine).
• Down a group (top to bottom), electronegativity generally decreases (e.g., fluorine to astatine).
• This trend is related to atomic radii: electronegativity increases across periods as atomic radii decrease, and it decreases down groups as atomic radii increase.

Relationship with Non-Metallic Properties:

• Non-metallic elements tend to gain electrons, making their electronegativity high.
• Electronegativity correlates with non-metallic properties of elements.
• Electronegativity is inversely related to metallic properties: higher electronegativity corresponds to lower metallic properties.
• Across a period, as electronegativity increases, non-metallic properties increase (metallic properties decrease).
• Down a group, as electronegativity decreases, non-metallic properties decrease (metallic properties increase).

Periodic Trends in Chemical Properties

• Valence and Oxidation States:
• Valence is a key characteristic property of elements and is linked to their electronic configurations.
• For representative elements, valence is often (though not always) equal to the number of electrons in their outermost orbitals, or eight minus the number of outermost electrons.
• The term "oxidation state" is commonly used interchangeably with valence.
• Oxidation States in Compounds:
• Consider compounds OF₂ and Na₂O with the elements F, O, and Na.
• Electronegativity order: F > O > Na.
• In OF₂, each fluorine (F) atom shares one electron with oxygen (O), resulting in F having an oxidation state of -1 due to its high electronegativity.
• Oxygen in OF₂ shares two electrons with fluorine atoms, leading to an oxidation state of +2.
• In Na2O, oxygen accepts two electrons (oxidation state -2) from two sodium (Na) atoms.
• Sodium loses an electron to oxygen, resulting in an oxidation state of +1.
• Oxidation state of an element in a compound is defined as the charge it acquires based on electronegativity considerations from other atoms in the molecule.
• Periodic Trends in Valence:
• Various periodic trends are observed in the valence of elements, specifically in hydrides and oxides.
• Table 3.9 displays some of these periodic trends in the valence of elements.
• The book discusses other periodic trends in the chemical behavior of elements separately.
• Variable Valence:
• Several elements exhibit variable valence.
• This variability is particularly notable in transition elements and actinoids.
• Detailed study of variable valence in these elements will be covered later in the material.

 

Anomalous Properties of Second Period Elements:

1. Covalent Character of First Elements in Groups 1 and 2:

• Lithium (Group 1) and beryllium (Group 2) exhibit covalent character in their compounds, unlike other alkali and alkaline earth metals.
• Contrast with the other group members that tend to form predominantly ionic compounds.

2. Diagonal Relationship:

• Similar behavior observed between lithium and magnesium, and between beryllium and aluminum.
• This relationship is referred to as diagonal relationship in the periodic properties.

3. Reasons for Different Chemical Behavior:

• Attributed to small size, large charge/radius ratio, and high electronegativity of the first group member.
• The small size and higher electronegativity influence bonding and reactivity patterns.

4. Valence Orbitals:

• First member of a group has only four valence orbitals (2s and 2p) available for bonding.
• Second member of the group possesses nine valence orbitals (3s, 3p, 3d).

5. Maximum Covalency:

• Due to limited valence orbitals, the first group member's maximum covalency is 4 (e.g., boron can only form BF₄).
• Other group members can expand their valence shells, accommodating more than four electron pairs (e.g., aluminum forms AlF₆³⁻).

6. Formation of Multiple Bonds:

• First member of p-block elements displays enhanced ability to form pπ - pπ multiple bonds.
• Examples include C=C, C≡C, N=N, N≡N.
• Also forms multiple bonds with other second period elements, like C=O, C=N, C≡N, N=O.

 

1. Atomic Radii Change in Transition Metals (3d Series):
• The change in atomic radii among transition metals (3d series) is notably smaller compared to representative elements in the same period.
• This trend also holds true for inner-transition metals (4f series), where the change in atomic radii is even smaller.
2. Ionization Enthalpies and Electropositivity:
• Transition metals' ionization enthalpies fall between those of s-block and p-block elements.
• Consequently, these metals exhibit lower electropositivity when compared to group 1 and 2 metals.
3. Group Trends: Increase in Atomic and Ionic Radii:
• Within a group, there's a consistent increase in atomic and ionic radii as the atomic number rises.
• This increase leads to a gradual decrease in ionization enthalpies.
• Electron gain enthalpies also generally decrease (with exceptions in some third period elements), particularly in the case of main group elements.
4. Group Trends: Metallic and Non-Metallic Character:
• Down a group, the trend shows an increase in metallic character and a decrease in non-metallic character.
• This shift in properties can be linked to the elements' reducing and oxidizing behaviors, which will be explained later.
5. Transition Elements' Exceptional Trend:
• In contrast to main group elements, transition elements exhibit a reverse trend.
• This anomalous trend can be rationalized by considering factors like atomic size and ionization enthalpy.

 

Top of Form