Position In The Periodic Table

The elements located in the middle section of the periodic table, encompassing groups 3 through 12, constitute the d-block. These elements are characterized by the progressive filling of their d orbitals across each of the four extended periods.

Situated separately at the bottom of the periodic table are the elements forming the f-block. This block includes elements where the 4f and 5f orbitals are being filled sequentially.

Commonly, elements of the d-block are referred to as transition metals, while those of the f-block are termed inner transition metals.

According to the International Union of Pure and Applied Chemistry (IUPAC), transition metals are defined as metals possessing an incomplete d subshell either in their neutral atomic state or in their common ionic states.

Elements in Group 12, namely Zinc (Zn), Cadmium (Cd), and Mercury (Hg), have a completely filled d$^{10}$ electron configuration in both their ground state and typical oxidation states. Consequently, they are not classified as transition metals based on the IUPAC definition. However, due to their position at the end of the 3d, 4d, and 5d series, their chemical properties are often studied alongside those of transition metals.

Electronic Configurations Of The D-Block Elements

The d-block elements reside in the central portion of the periodic table, positioned between the s-block and p-block elements. The filling of d-orbitals in the second to outermost energy level ((n-1)d) gives rise to four series of transition metals: the 3d, 4d, 5d, and 6d series.

The general electronic configuration for the outer orbitals of these elements is commonly represented as $(n-1)d^{1-10} ns^{1-2}$. Here, (n-1) refers to the inner d orbitals which can accommodate between one and ten electrons, while the outermost ns orbital typically holds one or two electrons.

It is important to note that this general configuration has exceptions. These exceptions arise because the energy difference between the (n-1)d and ns orbitals is very small. Additionally, electron configurations resulting in half-filled (d$^5$) or completely filled (d$^{10}$) d orbitals exhibit enhanced stability.

Notable examples in the 3d series include Chromium (Cr) and Copper (Cu). Chromium adopts the configuration 3d$^5$ 4s$^1$ instead of the expected 3d$^4$ 4s$^2$. Similarly, Copper has the configuration 3d$^{10}$ 4s$^1$ instead of 3d$^9$ 4s$^2$. These deviations maximise the number of half-filled or fully filled orbitals for greater stability.

Element	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
Z	21	22	23	24	25	26	27	28	29	30
4s	2	2	2	1	2	2	2	2	1	2
3d	1	2	3	5	5	6	7	8	10	10
Element	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd
Z	39	40	41	42	43	44	45	46	47	48
5s	2	2	1	1	1	1	1	0	1	2
4d	1	2	4	5	6	7	8	10	10	10
Element	La	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg
Z	57	72	73	74	75	76	77	78	79	80
6s	2	2	2	2	2	2	2	1	1	2
5d	1	2	3	4	5	6	7	9	10	10
Element	Ac	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn
Z	89	104	105	106	107	108	109	110	111	112
7s	2	2	2	2	2	2	2	2	1	2
6d	1	2	3	4	5	6	7	8	10	10

As noted earlier, elements like Zn, Cd, Hg, and Cn have a general electronic configuration of $(n-1)d^{10} ns^2$. Since their d orbitals are completely filled in both their ground state and typical ionic states, they do not fit the definition of transition elements.

A key characteristic is that the d orbitals extend more outwards compared to s and p orbitals. This makes them highly susceptible to influence from their surroundings and allows them to significantly affect surrounding atoms or molecules. The presence of partially filled d orbitals (d$^{1}$ to d$^{9}$) is responsible for many characteristic properties of transition elements, including:

• Exhibiting a wide range of oxidation states.
• Forming coloured ions.
• Engaging in complex formation with various ligands.
• Displaying catalytic activity.
• Showing paramagnetic behaviour.

Unlike non-transition elements, transition elements in a horizontal series often exhibit greater similarities in properties. However, some similarities also exist within groups.

Answer:

General Properties Of The Transition Elements (D-Block)

Physical Properties

Most transition elements exhibit characteristic metallic properties. These include high tensile strength, the ability to be drawn into wires (ductility), shaped by hammering (malleability), excellent thermal and electrical conductivity, and a typical metallic lustre.

With the notable exceptions of Zinc (Zn), Cadmium (Cd), Mercury (Hg), and Manganese (Mn), transition metals typically possess one or more standard metallic crystal structures at normal temperatures.

Series	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
1st (3d)	hcp	hcp	bcc	bcc	X (bcc)	bcc	ccp	ccp	ccp	X (hcp)
Series	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd
2nd (4d)	hcp	hcp	bcc	bcc	hcp	hcp	ccp	ccp	ccp	X (hcp)
Series	La	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg
3rd (5d)	hcp	hcp	bcc	bcc	hcp	hcp	ccp	ccp	ccp	X (bcc, ccp)
Series	Ac	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn
4th (6d)	hcp	hcp	bcc	bcc	hcp	hcp	ccp	ccp	ccp	X (bcc)

(bcc = body centred cubic; hcp = hexagonal close packed; ccp = cubic close packed; X = a typical metal structure, variations indicated in parenthesis).

Most transition metals (again, excluding Zn, Cd, and Hg) are characterized by their significant hardness and low tendency to vaporize (low volatility). They generally exhibit high melting and boiling points.

The exceptionally high melting points are attributed to the involvement of a large number of electrons from the inner $(n-1)d$ orbitals, in addition to the outermost $ns$ electrons, in forming strong interatomic metallic bonds.

Across a horizontal series of transition metals, the melting points tend to increase, reaching a maximum around the middle of the series (elements with d$^5$ configuration), before declining as the atomic number increases. Exceptions are observed for Mn and Tc, which show anomalous values.

Transition metals also possess high enthalpies of atomisation. The peak values around the middle of each series suggest that having one unpaired electron per d orbital is particularly favourable for strong interatomic bonding interactions. Generally, a larger number of valence electrons available for bonding correlates with stronger resulting bonds and higher enthalpy of atomisation.

Since the enthalpy of atomisation is a factor influencing the standard electrode potential, metals with very high enthalpies of atomisation (meaning very high boiling points) tend to be less reactive or 'noble'.

When comparing different series, the elements of the second (4d) and third (5d) series generally have higher enthalpies of atomisation than their corresponding elements in the first (3d) series. This difference contributes to the more frequent occurrence of metal-metal bonding in compounds of heavier transition metals.

Variation In Atomic And Ionic Sizes Of Transition Metals

In a given series, ions with the same charge show a gradual decrease in radius as the atomic number increases from left to right. This trend is explained by the increasing nuclear charge (+1 for each increment in atomic number) accompanied by the addition of electrons into the inner d orbitals. The shielding effect provided by a d electron for other electrons in the same shell is not as effective as the shielding by electrons in inner shells. Consequently, the effective nuclear attraction experienced by the outermost electrons increases, leading to a contraction in ionic size.

A similar decreasing trend is observed in the atomic radii within a series, although the variation is relatively small.

An interesting comparison arises between the atomic sizes of elements in different series. Atomic radii generally increase from the first (3d) series to the second (4d) series. However, the radii of elements in the third (5d) series are remarkably similar to those of their corresponding elements in the second (4d) series.

This unexpected similarity is a direct consequence of the Lanthanoid contraction. This phenomenon occurs because the 4f orbitals are filled *before* the 5d series begins (from atomic number 58, Cerium, to 71, Lutetium). The filling of the 4f orbitals results in a steady decrease in atomic radii across the lanthanoid series. This contraction effectively offsets the expected increase in size that would normally occur when moving from the 4d to the 5d series.

The significant outcome of the lanthanoid contraction is that elements of the second and third d series have nearly identical radii (e.g., Zirconium (Zr) with 160 pm and Hafnium (Hf) with 159 pm). This similarity in size leads to very similar physical and chemical properties between corresponding elements in the 4d and 5d series, making their separation challenging.

The cause of the lanthanoid contraction is analogous to the contraction observed within a regular transition series: the imperfect shielding of valence electrons by inner electrons. Specifically, the shielding provided by one 4f electron for another is less effective than that provided by a d electron. As the nuclear charge increases along the 4f series, the effective nuclear attraction on the entire 4f$^n$ electron cloud increases, causing a regular decrease in size.

The combined effect of decreasing metallic radius (due to increasing atomic number and contraction) and increasing atomic mass results in a general increase in the density of transition elements across a series, particularly noticeable from Titanium (Z=22) to Copper (Z=29).

Element	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
Metallic/ionic radii/pm M	164	147	135	129	137	126	125	125	128	137
Metallic/ionic radii/pm M$^{2+}$	–	–	79	82	82	77	74	70	73	75
Metallic/ionic radii/pm M$^{3+}$	73	67	64	62	65	65	61	60	–	–
Density/g cm$^{-3}$	3.43	4.1	6.07	7.19	7.21	7.8	8.7	8.9	8.9	7.1

Ionisation Enthalpies

The ionisation enthalpy is the energy required to remove an electron from a gaseous atom or ion. There is a general increase in ionisation enthalpy across each series of transition elements from left to right. This increase is primarily due to the rise in nuclear charge as atomic number increases, while electrons are added to the inner (n-1)d orbitals.

However, the increase in successive ionisation enthalpies (first, second, third, etc.) for transition elements is less steep compared to that observed in non-transition elements across a period. This implies that the energy cost of removing varying numbers of electrons from the $(n-1)d$ orbitals is relatively less prohibitive, contributing to their ability to show multiple oxidation states.

The variation in ionisation enthalpy along a transition series is also much smaller than the variation observed along a period for non-transition elements.

While the first ionisation enthalpy generally increases, the magnitude of increase for the second and third ionisation enthalpies tends to be higher across the series for successive elements.

The irregular trend observed in the first ionisation enthalpy values for the 3d series has minor chemical significance. It can be understood by considering how removing an electron alters the relative energies of the 4s and 3d orbitals. When d-block elements form ions, electrons are removed from the $ns$ orbitals before the $(n-1)d$ orbitals.

Moving across the 3d series, the nuclear charge increases, but the additional electrons occupy the inner 3d orbitals. These 3d electrons provide some shielding to the outer 4s electrons from the increasing nuclear charge. This shielding is somewhat more effective than the shielding among electrons within the same shell. Consequently, the atomic radii decrease less significantly, and the ionisation energies show only a slight increase across the series.

For doubly or more highly charged ions, the configuration is $d^n$ without any 4s electrons. A steady increase in second and third ionisation enthalpies is generally expected as the effective nuclear charge increases, and one d electron is not very effective at shielding another d electron due to the shape and spatial distribution of d orbitals.

However, the trend of steadily increasing second and third ionisation enthalpies is disrupted for the formation of Mn$^{2+}$ and Fe$^{3+}$ ions, respectively. Both these ions have a stable d$^5$ configuration (Mn$^{2+}$ is [Ar] 3d$^5$, Fe$^{3+}$ is [Ar] 3d$^5$). Similar disruptions occur at corresponding elements in the later transition series.

The variation in ionisation enthalpy for a configuration $d^n$ is influenced by three main factors:

1. Attraction of electrons towards the nucleus.
2. Repulsion between electrons.
3. Exchange energy.

Exchange energy provides stability to an electronic state. It is roughly proportional to the number of possible pairs of electrons with parallel spins in degenerate orbitals. According to Hund's rule, the lowest energy state occurs when maximum orbitals are singly occupied with parallel spins. The loss of exchange energy upon ionisation increases stability, making ionisation more difficult.

There is no loss of exchange energy when forming a d$^6$ configuration. Mn$^+$ has 3d$^5$4s$^1$, and Cr$^+$ has 3d$^5$. Removing an electron from Mn$^+$ forms Mn$^{2+}$ (3d$^5$) which is stable. Removing an electron from Cr$^+$ forms Cr$^{2+}$ (3d$^4$) which loses the d$^5$ stability. Hence, the second ionisation enthalpy (IE2) of Mn (going from Mn$^+$ to Mn$^{2+}$) is lower than that of Cr (going from Cr$^+$ to Cr$^{2+}$).

Similarly, Fe$^{2+}$ has a d$^6$ configuration, and Mn$^{2+}$ has a d$^5$ configuration. The third ionisation enthalpy (IE3) involves removing an electron from Fe$^{2+}$ (d$^6$ $\rightarrow$ d$^5$) and Mn$^{2+}$ (d$^5$ $\rightarrow$ d$^4$). Removing an electron from the stable d$^5$ of Mn$^{2+}$ requires more energy than removing one from d$^6$ of Fe$^{2+}$ (which gains d$^5$ stability). Therefore, the third ionisation enthalpy of Fe is lower than that of Mn.

The lowest common oxidation state is +2. The energy required to form M$^{2+}$ ions from the solid metal involves the enthalpy of atomisation plus the sum of the first and second ionisation enthalpies. The second ionisation enthalpy is a significant factor. Notably high values for the sum of IE1 + IE2 are seen for Cr and Cu because forming M$^+$ ions involves achieving stable d$^5$ and d$^{10}$ configurations, respectively (Cr: 3d$^5$4s$^1$, Cu: 3d$^{10}$4s$^1$). The value for Zn is low because forming Zn$^+$ (3d$^{10}$4s$^1$) followed by Zn$^{2+}$ (3d$^{10}$) involves removing a 4s electron first, leading to a stable d$^{10}$ configuration.

The trend in third ionisation enthalpies (IE3) is simpler as it doesn't involve the 4s orbital factor. High IE3 values are observed for removing an electron from stable d$^5$ (Mn$^{2+}$) and d$^{10}$ (Zn$^{2+}$) ions. The high IE3 values for Cu, Ni, and Zn explain why obtaining oxidation states greater than +2 for these elements is difficult.

While ionisation enthalpies provide some insight into the relative stability of oxidation states, predicting stability is complex and not easily generalised.

Element	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
Ionisation enthalpy/DiHV/kJ mol–1 I	631	656	650	653	717	762	758	736	745	906
Ionisation enthalpy/DiHV/kJ mol–1 II	1235	1309	1414	1592	1509	1561	1644	1752	1958	1734
Ionisation enthalpy/DiHV/kJ mol–1 III	2393	2657	2833	2990	3260	2962	3243	3402	3556	3837

Answer:

Oxidation States

One of the most characteristic features of transition elements is their ability to display a variety of oxidation states in their compounds. This arises from the participation of both the $(n-1)d$ and $ns$ electrons in bonding, as the energies of these orbitals are quite similar.

Element	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
Common Oxidation States	+3	+2, +3, +4	+2, +3, +4, +5	+2, +3, +4, +5, +6	+2, +3, +4, +5, +6, +7	+2, +3, +4, +6	+2, +3, +4	+2, +3, +4	+1, +2	+2

(Most common oxidation states are in bold type).

The elements exhibiting the largest number of oxidation states are typically found in or near the middle of the series. Manganese (Mn), for instance, shows oxidation states ranging from +2 to +7.

At the beginning and end of the series, the number of accessible oxidation states is limited. Elements like Scandium (Sc) and Titanium (Ti) at the start have only a few d electrons to lose or share. Scandium primarily exhibits +3, and Ti(IV) is more stable than Ti(III) or Ti(II).

At the end of the series, elements like Copper (Cu) and Zinc (Zn) have almost or completely filled d subshells. Zinc exhibits only the +2 oxidation state, as no d electrons are involved in achieving this state. Copper commonly shows +1 and +2.

The maximum oxidation states with reasonable stability often correspond to the sum of the $s$ and $d$ electrons, particularly up to Manganese (e.g., Ti(IV) in TiO$_2$, V(V) in VO$_2^+$, Cr(VI) in CrO$_4^{2-}$, Mn(VII) in MnO$_4^-$). Beyond Manganese, there is a rapid decrease in the stability of higher oxidation states, with more typical states being Fe(II, III), Co(II, III), Ni(II), Cu(I, II), and Zn(II).

A key characteristic of transition elements' variable oxidation states is that they often differ by a unit of one (e.g., V(II), V(III), V(IV), V(V)). This contrasts with non-transition elements, where oxidation states typically differ by two (e.g., P(III), P(V) or S(II), S(IV), S(VI)).

Within groups (specifically groups 4 to 10), an interesting trend is observed, which is opposite to the inert pair effect seen in p-block elements. In d-block groups, the higher oxidation states are favoured by heavier members. For example, in Group 6, Molybdenum(VI) and Tungsten(VI) are more stable than Chromium(VI). This explains why acidified dichromate (Cr(VI)) is a strong oxidising agent, while MoO$_3$ and WO$_3$ are not.

Low oxidation states (like 0 or +1) are typically observed in complex compounds containing ligands that can accept electron density from the metal via $\pi$ bonding ($\pi$-acceptor ligands), in addition to forming $\sigma$ bonds. For instance, in metal carbonyls like Ni(CO)$_4$ and Fe(CO)$_5$, the oxidation state of the metal is zero.

Answer:

Trends In The M2+/M Standard Electrode Potentials

The standard electrode potential (E°) for the reduction of a divalent metal ion to the solid metal (M$^{2+}$/M) is influenced by the energy changes involved in transforming the solid metal into gaseous atoms (enthalpy of atomisation, $\Delta_{a}H^{\circ}$), ionising the gaseous atoms to gaseous ions (first and second ionisation enthalpies, $\Delta_{i}H^{\circ}_1 + \Delta_{i}H^{\circ}_2$), and hydrating the gaseous ions to form aqueous ions (enthalpy of hydration, $\Delta_{hyd}H^{\circ}$).

The overall energy change relates to the standard electrode potential. A more negative E° indicates a greater tendency for the metal to be oxidised (or the ion to be reduced with more difficulty).

Element (M)	$\Delta_{a}H^{\circ}$ (M)	$\Delta_{i}H^{\circ}_1$	$\Delta_{i}H^{\circ}_2$	$\Delta_{hyd}H^{\circ}$ (M$^{2+}$)	E$^{\circ}$/V M$^{2+}$/M
Ti	469	656	1309	-1866	-1.63
V	515	650	1414	-1895	-1.18
Cr	398	653	1592	-1925	-0.90
Mn	279	717	1509	-1862	-1.18
Fe	418	762	1561	-1998	-0.44
Co	427	758	1644	-2079	-0.28
Ni	431	736	1752	-2121	-0.25
Cu	339	745	1958	-2121	+0.34
Zn	130	906	1734	-2059	-0.76

Copper (Cu) exhibits unique behaviour with a positive E° value (+0.34 V) for the Cu$^{2+}$/Cu couple. This positive value means that the reduction of Cu$^{2+}$ to Cu is thermodynamically favoured over the oxidation of Cu to Cu$^{2+}$ under standard conditions compared to other transition metals. This is why copper is unable to displace hydrogen from non-oxidising acids like dilute HCl or H$_2$SO$_4$. Copper only reacts with oxidising acids (like nitric acid or hot concentrated sulphuric acid), where the acid itself is reduced.

The high energy required to convert solid copper to aqueous Cu$^{2+}$ ions is primarily due to the relatively high enthalpy of atomisation of Cu and the sum of its first and second ionisation enthalpies ($\Delta_aH^{\circ} + \Delta_iH^{\circ}_1 + \Delta_iH^{\circ}_2$). This energy input is not fully offset by the hydration enthalpy of the Cu$^{2+}$ ion ($\Delta_{hyd}H^{\circ}$).

Across the first transition series from left to right, there is a general trend towards less negative E° values for the M$^{2+}$/M couple. This trend is largely associated with the general increase in the sum of the first and second ionisation enthalpies across the series.

However, the E° values for Manganese (Mn), Nickel (Ni), and Zinc (Zn) are more negative than the general trend would suggest. The more negative values for Mn and Zn are related to the relative stability of the d$^5$ configuration in Mn$^{2+}$ and the d$^{10}$ configuration in Zn$^{2+}$. For Nickel, the relatively more negative E° value is associated with its particularly high (most negative) enthalpy of hydration for Ni$^{2+}$.

Trends In The M3+/M2+ Standard Electrode Potentials

Examining the E° values for the redox couple M$^{3+}$/M$^{2+}$ provides insight into the relative stability of the +3 and +2 oxidation states. A more positive E° value for M$^{3+}$/M$^{2+}$ means that the M$^{3+}$ ion is more easily reduced to M$^{2+}$, implying M$^{3+}$ is a stronger oxidising agent and M$^{2+}$ is relatively stable.

Referring to Table 8.2 (reproduced below for convenience):

Element	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
Standard electrode potential E$^{\circ}$/V M$^{3+}$/M$^{2+}$	–	–0.37	–0.26	–0.41	+1.57	+0.77	+1.97	–	–	–

The very low E° value for the Sc$^{3+}$/Sc$^{2+}$ couple (not listed, but Sc only stable as Sc$^{3+}$) reflects the exceptionally high stability of the Sc$^{3+}$ ion, which has a noble gas electronic configuration ([Ar] 3d$^0$). It does not readily accept an electron to form Sc$^{2+}$.

Zinc has no stable +3 state, so the Zn$^{3+}$/Zn$^{2+}$ potential is not relevant. The hypothetical removal of an electron from Zn$^{2+}$ (3d$^{10}$) is very difficult, hence a very high positive potential if we considered Zn$^{3+}$/Zn$^{2+}$.

The comparatively high positive value (+1.57 V) for the Mn$^{3+}$/Mn$^{2+}$ couple indicates that Mn$^{3+}$ is easily reduced to Mn$^{2+}$. This is because Mn$^{2+}$ has a stable d$^5$ configuration. Thus, Mn$^{3+}$ is a strong oxidising agent.

The comparatively low positive value (+0.77 V) for the Fe$^{3+}$/Fe$^{2+}$ couple suggests that Fe$^{3+}$ is also relatively stable (due to its d$^5$ configuration), but less so than Mn$^{2+}$. Fe$^{3+}$ can be reduced to Fe$^{2+}$, but it is a weaker oxidising agent compared to Mn$^{3+}$.

The very high positive value (+1.97 V) for the Co$^{3+}$/Co$^{2+}$ couple shows that Co$^{3+}$ is a very strong oxidising agent, readily reduced to Co$^{2+}$. Co$^{2+}$ (d$^7$) is significantly more stable in aqueous solution than Co$^{3+}$ (d$^6$).

The low value for the V$^{3+}$/V$^{2+}$ couple (–0.26 V) suggests that V$^{2+}$ is quite stable. This stability is related to the half-filled t$_{2g}$ level in the octahedral crystal field (a concept from Unit 9).

Answer:

Trends In Stability Of Higher Oxidation States

The highest oxidation states for transition metals are most frequently observed when they combine with highly electronegative elements like Fluorine (F) and Oxygen (O). This is because fluorine and oxygen can effectively oxidise the metal to its maximum possible state.

Groups	3	4	5	6	7	8	9	10	11	12
Oxidation Number	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
+6				CrF$_6$
+5			VF$_5$	CrF$_5$
+4		TiX$_4$	VX$^{\text{I}}_4$	CrX$_4$	MnF$_4$
+3	ScX$_3$	TiX$_3$	VX$_3$	CrX$_3$	MnF$_3$	FeX$^{\text{I}}_3$	CoF$_3$
+2		TiX$_2$	VX$_2$	CrX$_2$	MnX$_2$	FeX$_2$	CoX$_2$	NiX$_2$	CuX$_2$	ZnX$_2$
+1									CuX$^{\text{III}}$

Key: X = F $\rightarrow$ I; X$^I$ = F $\rightarrow$ Br; X$^{II}$ = F, Cl; X$^{III}$ = Cl $\rightarrow$ I

The highest oxidation states achieved in simple halides are Ti(IV) in TiX$_4$, V(V) in VF$_5$, and Cr(VI) in CrF$_6$. Manganese's +7 state is not seen in simple halides, though MnO$_3$F exists. Beyond Manganese, achieving a +3 state in halides becomes less common, primarily limited to FeX$_3$ and CoF$_3$.

Fluorine's ability to stabilise high oxidation states is linked to either high lattice energies in ionic compounds (like CoF$_3$) or strong bond enthalpies in covalent compounds (like VF$_5$, CrF$_6$).

Fluorides can also exhibit instability in low oxidation states, such as VX$_2$ (X = Cl, Br, I) and CuX. Conversely, almost all Copper(II) halides are known, except CuI$_2$. This is because Cu$^{2+}$ ions oxidise iodide ions (I$^-$) to iodine (I$_2$) in solution:

$2\text{Cu}^{2+}\text{(aq)} + 4\text{I}^-\text{(aq)} \rightarrow 2\text{CuI(s)} + \text{I}_2\text{(aq)}$

Many Copper(I) compounds are unstable in aqueous solution and undergo disproportionation, where the ion is simultaneously oxidised and reduced:

$2\text{Cu}^+\text{(aq)} \rightarrow \text{Cu}^{2+}\text{(aq)} + \text{Cu(s)}$

The reason Cu$^{2+}$(aq) is more stable than Cu$^+$(aq) is that the hydration enthalpy ($\Delta_{hyd}H^{\circ}$) of Cu$^{2+}$ is significantly more negative than that of Cu$^+$. This large negative hydration enthalpy for Cu$^{2+}$ provides enough energy to more than compensate for the relatively high second ionisation enthalpy needed to form Cu$^{2+}$ from Cu$^+$.

Oxygen is even more effective than fluorine at stabilising the highest oxidation states. In oxides, the highest oxidation number matches the group number up to Manganese (from Sc$_2$O$_3$ for Group 3 to Mn$_2$O$_7$ for Group 7). Beyond Group 7, stable oxides do not show the maximum possible oxidation state; for instance, the highest stable oxide of Iron is Fe$_2$O$_3$ (+3 state), although unstable ferrates(VI) (FeO$_4^{2-}$) exist in alkaline media.

Oxocations also play a role in stabilising high oxidation states, such as V(V) in VO$_2^+$, V(IV) in VO$^{2+}$, and Ti(IV) in TiO$^{2+}$.

Oxygen's superiority in stabilising high oxidation states is likely due to its ability to form multiple bonds with the metal atom. In the covalent oxide Mn$_2$O$_7$, each Mn atom is linked by a Mn-O-Mn bridge, and the molecule involves multiple bonds between Mn and terminal O atoms.

Tetrahedral oxoanions [MO$_4$]$^{n-}$ are known for several metals in high oxidation states, including V(V), Cr(VI), Mn(V), Mn(VI), and Mn(VII).

Groups	3	4	5	6	7	8	9	10	11	12
Oxidation Number	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
+7					Mn$_2$O$_7$
+6				CrO$_3$
+5			V$_2$O$_5$
+4		TiO$_2$	V$_2$O$_4$	CrO$_2$	MnO$_2$
+3	Sc$_2$O$_3$	Ti$_2$O$_3$	V$_2$O$_3$	Cr$_2$O$_3$	Mn$_2$O$_3$	Fe$_2$O$_3$
Mixed Oxides					Mn$_3$O$_4$*	Fe$_3$O$_4$*	Co$_3$O$_4$*
+2		TiO	VO	(CrO)	MnO	FeO	CoO	NiO	CuO	ZnO
+1									Cu$_2$O

* mixed oxides

Answer:

Chemical Reactivity And EV Values

Transition metals exhibit a wide range of chemical reactivity. Many of them are electropositive enough to dissolve in mineral acids, while others are considered 'noble' because they are unreactive towards single acids.

Most metals of the first (3d) series, with the exception of copper, are relatively reactive. They can be oxidised by 1 M H$^+$ ions, although their reaction rate with oxidising agents like hydrogen ions can sometimes be slow. For instance, Titanium and Vanadium are practically passive to dilute non-oxidising acids at room temperature.

The E° values for the M$^{2+}$/M couple (listed in Table 8.2) generally indicate a decreasing tendency for metals to form stable divalent cations across the series from left to right. This general trend towards less negative E° values is related to the increasing sum of the first and second ionisation enthalpies.

As discussed previously, the E° values for Mn, Ni, and Zn are more negative than expected based on this general trend. The increased stability of d$^5$ in Mn$^{2+}$ and d$^{10}$ in Zn$^{2+}$ contributes to their more negative potentials, while for Nickel, it is related to the highly negative hydration enthalpy of Ni$^{2+}$.

Looking at the E° values for the M$^{3+}$/M$^{2+}$ redox couple (Table 8.2), it is evident that Mn$^{3+}$ and Co$^{3+}$ ions are the strongest oxidising agents in aqueous solutions because they are easily reduced to the more stable Mn$^{2+}$ and Co$^{2+}$ ions respectively.

Conversely, ions like Ti$^{2+}$, V$^{2+}$, and Cr$^{2+}$ are strong reducing agents. They readily lose electrons and are strong enough reductants to liberate hydrogen gas from dilute acids. For example:

$2\text{Cr}^{2+}\text{(aq)} + 2\text{H}^+\text{(aq)} \rightarrow 2\text{Cr}^{3+}\text{(aq)} + \text{H}_2\text{(g)}$

Answer:

Answer:

Magnetic Properties

When exposed to a magnetic field, substances exhibit different behaviours. The main types are diamagnetism (repelled by the field) and paramagnetism (attracted by the field). Substances that are very strongly attracted are called ferromagnetic (an extreme form of paramagnetism, like Fe, Co, Ni).

Many transition metal ions are paramagnetic. This property arises from the presence of unpaired electrons in their d orbitals. Each unpaired electron possesses a magnetic moment due to its spin and orbital motion.

For compounds of the first transition series (3d elements), the contribution of orbital angular momentum to the magnetic moment is often negligible or "quenched" due to interaction with the surrounding environment (ligand field). Therefore, the magnetic moment is primarily determined by the number of unpaired electrons and can be calculated using the 'spin-only' formula:

$\mu = \sqrt{n(n + 2)}$

where:

• $\mu$ is the magnetic moment, measured in Bohr Magnetons (BM).
• $n$ is the number of unpaired electrons in the ion or atom.

A single unpaired electron has a spin-only magnetic moment of $\sqrt{1(1+2)} = \sqrt{3} \approx 1.73$ BM.

The magnetic moment value increases with the number of unpaired electrons. Therefore, experimentally measured magnetic moments are valuable indicators of the number of unpaired electrons present in a transition metal species.

Ion	Configuration	Unpaired electron(s) (n)	Magnetic moment Calculated (BM)	Magnetic moment Observed (BM)
Sc$^{3+}$	3d$^0$	0	0	0
Ti$^{3+}$	3d$^1$	1	1.73	1.75
Tl$^{2+}$	3d$^2$	2	2.84	2.76
V$^{2+}$	3d$^3$	3	3.87	3.86
Cr$^{2+}$	3d$^4$	4	4.90	4.80
Mn$^{2+}$	3d$^5$	5	5.92	5.96
Fe$^{2+}$	3d$^6$	4	4.90	5.3 – 5.5
Co$^{2+}$	3d$^7$	3	3.87	4.4 – 5.2
Ni$^{2+}$	3d$^8$	2	2.84	2.9 – 3.4
Cu$^{2+}$	3d$^9$	1	1.73	1.8 – 2.2
Zn$^{2+}$	3d$^{10}$	0	0	0

(Observed values are typically for hydrated ions in solution or solid state).

Answer:

Formation Of Coloured Ions

Many transition metal ions in solution or in solid compounds are coloured. This is a characteristic property linked to the presence of partially filled d orbitals.

The colour arises from the absorption of light energy that causes an electron to transition or be excited from a lower energy d orbital to a higher energy d orbital within the same subshell. This process is called a d-d transition.

In the presence of ligands or counterions, the d orbitals of the metal ion are split into different energy levels (as explained by Crystal Field Theory in Unit 9). The energy difference between these levels corresponds to specific frequencies of electromagnetic radiation.

When white light passes through a transition metal compound, the ion absorbs light of certain frequencies corresponding to the energy needed for d-d transitions. The colours that are *not* absorbed are transmitted or reflected, and these are the colours that we perceive.

The energy required for d-d transitions typically falls within the visible region of the electromagnetic spectrum. The specific colour observed is the complementary colour of the light that was absorbed. For instance, if a substance absorbs red light, it will appear green.

The frequency of light absorbed (and thus the colour observed) depends on the energy splitting of the d orbitals, which in turn is influenced by factors like the nature of the metal ion, its oxidation state, and the nature of the ligands surrounding it.

Ions with completely empty (d$^0$) or completely filled (d$^{10}$) d subshells cannot undergo d-d transitions because there are no electrons to excite or no vacant higher-energy d orbitals to receive them. Consequently, such ions are typically colourless.

Configuration	Example	Colour
3d$^0$	Sc$^{3+}$	colourless
3d$^0$	Ti$^{4+}$	colourless
3d$^1$	Ti$^{3+}$	purple
3d$^1$	V$^{4+}$	blue
3d$^2$	V$^{3+}$	green
3d$^3$	V$^{2+}$	violet
3d$^3$	Cr$^{3+}$	violet
3d$^4$	Mn$^{3+}$	violet
3d$^4$	Cr$^{2+}$	blue
3d$^5$	Mn$^{2+}$	pink
3d$^5$	Fe$^{3+}$	yellow
3d$^6$	Fe$^{2+}$	green
3d$^6$	Co$^{3+}$	blue
3d$^7$	Co$^{2+}$	pink
3d$^8$	Ni$^{2+}$	green
3d$^9$	Cu$^{2+}$	blue
3d$^{10}$	Zn$^{2+}$	colourless

(Colours listed are for aquated ions, where water molecules act as ligands).

(The image would visually show examples like aqueous solutions of V$^{4+}$ (blue), V$^{3+}$ (green), Mn$^{2+}$ (pink), Fe$^{3+}$ (yellow/brown), Co$^{2+}$ (pink), Ni$^{2+}$ (green), and Cu$^{2+}$ (blue)).

Formation Of Complex Compounds

Transition metals are known for forming a vast number of complex compounds (also called coordination compounds). In these compounds, a central metal ion (usually a transition metal) is bonded to a number of surrounding ions or neutral molecules, which are called ligands. These complexes have distinct properties from the constituent ions or molecules.

Examples include [Fe(CN)$_6$]$^{3-}$, [Fe(CN)$_6$]$^{4-}$, [Cu(NH$_3$)$_4$]$^{2+}$, and [PtCl$_4$]$^{2-}$. The formation and properties of complex compounds are studied in detail in Unit 9.

The ability of transition metals to form a large variety of complexes is attributed to three main factors:

1. Relatively smaller size of the metal ions.
2. High ionic charges (or high charge density) on the metal ions.
3. Availability of vacant d orbitals of appropriate energy to accept electron pairs from the ligands (for coordinate covalent bond formation).

Catalytic Properties

Many transition metals and their compounds exhibit significant catalytic activity, meaning they can increase the rate of chemical reactions without being consumed themselves. This property is closely related to two features of transition elements:

1. Their ability to exist in multiple oxidation states.
2. Their capacity to form complex compounds.

Examples of transition metals or their compounds used as catalysts include Vanadium(V) oxide (V$_2$O$_5$) in the Contact Process for sulfuric acid production, finely divided Iron (Fe) in the Haber's Process for ammonia synthesis, and Nickel (Ni) in catalytic hydrogenation reactions.

In heterogeneous catalysis (where the catalyst is in a different phase, typically solid, than the reactants), surface catalysis by transition metals involves the formation of temporary bonds between the reactant molecules and the atoms on the catalyst surface. First-row transition metals utilise their 3d and 4s electrons for this bonding. This interaction helps in two ways:

• It increases the concentration of reactants on the catalyst surface.
• It weakens the existing bonds within the reacting molecules, thereby lowering the activation energy required for the reaction to occur.

Furthermore, the ability of transition metal ions to change their oxidation states allows them to participate in the reaction pathway by forming intermediate compounds, making them more effective catalysts. For instance, Iron(III) ions (Fe$^{3+}$) catalyse the reaction between iodide ions (I$^-$) and persulphate ions (S$_2$O$_8^{2-}$):

$2\text{I}^-\text{(aq)} + \text{S}_2\text{O}_8^{2-}\text{(aq)} \rightarrow \text{I}_2\text{(aq)} + 2\text{SO}_4^{2-}\text{(aq)}$

This reaction is believed to occur via the following steps involving the interconversion of iron oxidation states:

$2\text{Fe}^{3+}\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow 2\text{Fe}^{2+}\text{(aq)} + \text{I}_2\text{(aq)}$

$2\text{Fe}^{2+}\text{(aq)} + \text{S}_2\text{O}_8^{2-}\text{(aq)} \rightarrow 2\text{Fe}^{3+}\text{(aq)} + 2\text{SO}_4^{2-}\text{(aq)}$

The Fe$^{3+}$ catalyst is regenerated at the end of the reaction cycle.

Formation Of Interstitial Compounds

Interstitial compounds are formed when small non-metal atoms like Hydrogen (H), Carbon (C), or Nitrogen (N) get trapped and occupy the voids or interstitial spaces within the crystal lattice of a metal. These compounds are well-known for transition metals because their crystal lattices often have appropriate sized interstitial holes.

These compounds are typically non-stoichiometric, meaning the ratio of the non-metal atom to the metal atom is not fixed and can vary within a range (e.g., TiC$_{0.98}$ or TiH$_{1.7}$). They do not conform to typical ionic or covalent bonding rules; their bonding is complex and includes contributions from metallic bonding.

Examples include TiC, Mn$_4$N, Fe$_3$H, VH$_{0.56}$, and TiH$_{1.7}$. The formulas reflect the composition range rather than fixed valency.

Interstitial compounds formed by transition metals possess several characteristic physical and chemical properties that differ significantly from the parent metals:

1. They have high melting points, often considerably higher than those of the pure transition metals.
2. They are remarkably hard; some metal borides, for instance, are nearly as hard as diamond.
3. They generally retain metallic conductivity, although it may be lower than the parent metal.
4. They are typically very chemically inert, being resistant to attack by acids, bases, and other reagents.

Alloy Formation

An alloy is a material composed of two or more metallic elements, or a metal and a non-metal, that are mixed together. Transition metals readily form alloys because they have similar atomic sizes and other metallic characteristics. Alloys are formed by metals whose atomic radii differ by no more than approximately 15 percent, allowing atoms of one metal to substitute for atoms of the other in the crystal lattice, forming homogeneous solid solutions.

Alloys formed by transition metals are often harder and have higher melting points than the constituent pure metals.

Some of the most important alloys are ferrous alloys, which are based on iron. Adding transition metals like Chromium (Cr), Vanadium (V), Tungsten (W), Molybdenum (Mo), and Manganese (Mn) to iron produces various types of steel, including stainless steel, which are crucial for construction and manufacturing.

Alloys involving transition metals and non-transition metals are also industrially significant. Examples include Brass (an alloy of Copper and Zinc) and Bronze (an alloy of Copper and Tin).

Some Important Compounds Of Transition Elements

Oxides And Oxoanions Of Metals

Transition metal oxides are typically formed by reacting the metals directly with oxygen at high temperatures. Most transition metals, with the exception of Scandium, form monoxides with the formula MO, which are ionic in nature.

As discussed earlier, the highest oxidation state observed in oxides often corresponds to the group number, up to Manganese (Group 7). For example, Sc$_2$O$_3$ (+3), TiO$_2$ (+4), V$_2$O$_5$ (+5), CrO$_3$ (+6), and Mn$_2$O$_7$ (+7).

Beyond Group 7, achieving the highest possible oxidation state in simple oxides becomes less common. For instance, the highest stable oxide of iron is Fe$_2$O$_3$ (+3), not FeO$_3$ (+6).

Besides simple oxides, transition metals can form oxocations that stabilise higher oxidation states, such as VO$_2^+$ (V in +5 state), VO$^{2+}$ (V in +4 state), and TiO$^{2+}$ (Ti in +4 state).

A notable trend in the oxides of transition metals is the relationship between oxidation state and acidic character. As the oxidation number of the metal increases, the ionic character of the oxide decreases, and the acidic character becomes predominant. For example, Mn$_2$O$_7$ (+7) is a covalent green oil and is acidic, forming permanganic acid (HMnO$_4$). CrO$_3$ (+6) is also acidic, forming chromic acid (H$_2$CrO$_4$) and dichromic acid (H$_2$Cr$_2$O$_7$).

Vanadium oxides show a gradual change in acidity with increasing oxidation state: V$_2$O$_3$ (+3) is basic, V$_2$O$_4$ (+4) is less basic, and V$_2$O$_5$ (+5) is amphoteric (reacts with both acids and bases), although mainly acidic. V$_2$O$_5$ dissolves in bases to form vanadates (like VO$_4^{3-}$) and in strong acids to form oxocations (like VO$_2^+$ or VO$^{2+}$ salts).

Similarly, CrO (+2) is basic, while Cr$_2$O$_3$ (+3) is amphoteric.

Potassium Dichromate K2Cr2O7

Potassium dichromate (K$_2$Cr$_2$O$_7$) is an important chemical widely used in the leather industry and as an oxidising agent in organic synthesis and volumetric analysis.

Dichromates are typically produced from chromates. Chromates are obtained by fusing chromite ore (FeCr$_2$O$_4$) with sodium or potassium carbonate in the presence of air:

$4\text{FeCr}_2\text{O}_4\text{(chromite)} + 8\text{Na}_2\text{CO}_3 + 7\text{O}_2 \xrightarrow{\text{fusion}} 8\text{Na}_2\text{CrO}_4\text{(yellow)} + 2\text{Fe}_2\text{O}_3 + 8\text{CO}_2$

The resulting yellow solution of sodium chromate (Na$_2$CrO$_4$) is filtered and then acidified, usually with sulfuric acid. Acidification converts the chromate ions to orange dichromate ions:

$2\text{Na}_2\text{CrO}_4 + 2\text{H}^+ \rightarrow \text{Na}_2\text{Cr}_2\text{O}_7\text{(orange)} + 2\text{Na}^+ + \text{H}_2\text{O}$

Sodium dichromate (Na$_2$Cr$_2$O$_7 \cdot 2\text{H}_2\text{O}$) can be crystallised from this solution. Since sodium dichromate is more soluble in water than potassium dichromate, potassium dichromate (K$_2$Cr$_2$O$_7$) is usually prepared by treating the sodium dichromate solution with potassium chloride. Due to its lower solubility, potassium dichromate precipitates out:

$\text{Na}_2\text{Cr}_2\text{O}_7 + 2\text{KCl} \rightarrow \text{K}_2\text{Cr}_2\text{O}_7\text{(orange crystals)} + 2\text{NaCl}$

Chromates and dichromates exist in equilibrium in aqueous solution, and their interconversion is dependent on the pH of the solution. In acidic conditions, chromate ions (CrO$_4^{2-}$) are converted to dichromate ions (Cr$_2$O$_7^{2-}$). In alkaline conditions, dichromate ions are converted back to chromate ions. The oxidation state of Chromium in both chromate and dichromate ions is +6.

Interconversion reactions:

$2\text{CrO}_4^{2-}\text{(yellow)} + 2\text{H}^+ \rightleftharpoons \text{Cr}_2\text{O}_7^{2-}\text{(orange)} + \text{H}_2\text{O}$

The chromate ion (CrO$_4^{2-}$) has a tetrahedral structure. The dichromate ion (Cr$_2$O$_7^{2-}$) consists of two tetrahedra sharing a common oxygen atom, forming a bent Cr-O-Cr bridge with a bond angle of 126°.

(The image would show the structures: CrO$_4^{2-}$ as a tetrahedron with Cr at the center and O at vertices; Cr$_2$O$_7^{2-}$ as two such tetrahedra connected by a shared oxygen atom, with the Cr-O-Cr angle shown as 126°).

Sodium and potassium dichromates are powerful oxidising agents. Sodium dichromate is more soluble and often used in organic chemistry. Potassium dichromate is frequently used as a primary standard in volumetric analysis.

In acidic solution, the dichromate ion is reduced to the Chromium(III) ion (Cr$^{3+}$). The half-reaction for its oxidising action is:

$\text{Cr}_2\text{O}_7^{2-}\text{(aq)} + 14\text{H}^+\text{(aq)} + 6e^- \rightarrow 2\text{Cr}^{3+}\text{(aq)} + 7\text{H}_2\text{O(l)} \quad (E^{\circ} = +1.33 \text{ V})$

Acidified potassium dichromate can oxidise various substances. Examples of half-reactions for typical reducing agents include:

• Iodide ions (I$^-$) to Iodine (I$_2$): $6\text{I}^- \rightarrow 3\text{I}_2 + 6e^-$
• Tin(II) ions (Sn$^{2+}$) to Tin(IV) ions (Sn$^{4+}$): $3\text{Sn}^{2+} \rightarrow 3\text{Sn}^{4+} + 6e^-$
• Hydrogen sulfide (H$_2$S) to Sulfur (S): $3\text{H}_2\text{S} \rightarrow 6\text{H}^+ + 3\text{S} + 6e^-$
• Iron(II) ions (Fe$^{2+}$) to Iron(III) ions (Fe$^{3+}$): $6\text{Fe}^{2+} \rightarrow 6\text{Fe}^{3+} + 6e^-$

To obtain the balanced ionic equation for the reaction between dichromate and a reducing agent, the dichromate half-reaction is combined with the half-reaction of the reducing agent, balancing the number of electrons exchanged. For example, with Fe$^{2+}$:

$\text{Cr}_2\text{O}_7^{2-}\text{(aq)} + 14\text{H}^+\text{(aq)} + 6\text{Fe}^{2+}\text{(aq)} \rightarrow 2\text{Cr}^{3+}\text{(aq)} + 6\text{Fe}^{3+}\text{(aq)} + 7\text{H}_2\text{O(l)}$

Potassium Permanganate KMnO4

Potassium permanganate (KMnO$_4$) is another powerful oxidising agent. It is prepared by the fusion of Manganese dioxide (MnO$_2$, pyrolusite ore) with an alkali metal hydroxide (like KOH) and an oxidising agent such as KNO$_3$ or air. This process yields dark green potassium manganate (K$_2$MnO$_4$).

$2\text{MnO}_2 + 4\text{KOH} + \text{O}_2 \rightarrow 2\text{K}_2\text{MnO}_4\text{(green)} + 2\text{H}_2\text{O}$

The green manganate solution is then treated in a neutral or acidic solution where it undergoes disproportionation to produce the purple permanganate:

$3\text{MnO}_4^{2-}\text{(green)} + 4\text{H}^+ \rightarrow 2\text{MnO}_4^-\text{(purple)} + \text{MnO}_2\text{(s)} + 2\text{H}_2\text{O}$

Alternatively, potassium permanganate can be prepared commercially by alkaline oxidative fusion of MnO$_2$ followed by electrolytic oxidation of the manganate(VI) ions in alkaline solution:

$\text{MnO}_2 \xrightarrow{\text{Fused with KOH, Oxidised with air or KNO}_3} \text{MnO}_4^{2-}\text{(manganate ion)}$

$\text{MnO}_4^{2-}\text{(manganate)} \xrightarrow{\text{Electrolytic oxidation in alkaline solution}} \text{MnO}_4^-\text{(permanganate ion)}$

In the laboratory, potassium permanganate can be prepared by oxidising Manganese(II) ions (Mn$^{2+}$) using peroxodisulphate ions (S$_2$O$_8^{2-}$):

$2\text{Mn}^{2+}\text{(aq)} + 5\text{S}_2\text{O}_8^{2-}\text{(aq)} + 8\text{H}_2\text{O(l)} \rightarrow 2\text{MnO}_4^-\text{(aq)} + 10\text{SO}_4^{2-}\text{(aq)} + 16\text{H}^+\text{(aq)}$

Potassium permanganate forms dark purple (almost black) crystals. It is isostructural with KClO$_4$. It has moderate solubility in water (6.4 g per 100 g water at 293 K) and decomposes when heated above 513 K:

$2\text{KMnO}_4\text{(s)} \xrightarrow{\Delta} \text{K}_2\text{MnO}_4\text{(s)} + \text{MnO}_2\text{(s)} + \text{O}_2\text{(g)}$

Potassium permanganate is known for its intense colour and its magnetic properties (diamagnetic with temperature-dependent weak paramagnetism), which are explained by molecular orbital theory.

Both the manganate (MnO$_4^{2-}$) and permanganate (MnO$_4^-$) ions have a tetrahedral structure. Pi ($\pi$) bonding is involved, formed by the overlap of oxygen p orbitals with manganese d orbitals. The green manganate ion (MnO$_4^{2-}$) is paramagnetic because it has one unpaired electron. The purple permanganate ion (MnO$_4^-$) is diamagnetic as it has no unpaired electrons.

Potassium permanganate is a strong oxidising agent, but its reduction product depends on the pH of the solution:

• In strongly acidic solution, MnO$_4^-$ is reduced to Mn$^{2+}$ (Manganese(II), +2 state). $E^{\circ} = +1.52 \text{ V}$
• In neutral or faintly alkaline solution, MnO$_4^-$ is reduced to MnO$_2$ (Manganese(IV), +4 state). $E^{\circ} = +1.69 \text{ V}$ for MnO$_4^-$ to MnO$_2$ (in acid) or $E^{\circ} = +0.56 \text{ V}$ for MnO$_4^-$ to MnO$_4^{2-}$ (in alkaline).
• In strongly alkaline solution, MnO$_4^-$ is reduced to MnO$_4^{2-}$ (Manganate(VI), +6 state).

The high positive reduction potential in acidic medium (+1.52 V) suggests that acidic permanganate should be a very strong oxidising agent, capable of oxidising water ($2\text{H}_2\text{O} \rightarrow \text{O}_2 + 4\text{H}^+ + 4e^-$, $E^{\circ} = +1.23 \text{ V}$). While thermodynamically favourable, the reaction with water is kinetically very slow unless catalysed by Mn$^{2+}$ ions or at elevated temperatures.

Important oxidising reactions of KMnO$_4$:

1. In Acid Solutions:

• Oxidation of Iodide (I$^-$) to Iodine (I$_2$): $10\text{I}^- + 2\text{MnO}_4^- + 16\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 8\text{H}_2\text{O} + 5\text{I}_2$
• Oxidation of Iron(II) (Fe$^{2+}$) to Iron(III) (Fe$^{3+}$): $5\text{Fe}^{2+}\text{(green)} + \text{MnO}_4^- + 8\text{H}^+ \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O} + 5\text{Fe}^{3+}\text{(yellow)}$
• Oxidation of Oxalate (C$_2$O$_4^{2-}$) or Oxalic acid (H$_2$C$_2$O$_4$) to Carbon dioxide (CO$_2$) (requires heating to ~333 K): $5\text{C}_2\text{O}_4^{2-} + 2\text{MnO}_4^- + 16\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 8\text{H}_2\text{O} + 10\text{CO}_2$
• Oxidation of Hydrogen sulfide (H$_2$S) to precipitated Sulfur (S): $5\text{S}^{2-} + 2\text{MnO}_4^- + 16\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 8\text{H}_2\text{O} + 5\text{S(s)}$
• Oxidation of Sulfurous acid (H$_2$SO$_3$) or Sulfite (SO$_3^{2-}$) to Sulfate (SO$_4^{2-}$): $5\text{SO}_3^{2-} + 2\text{MnO}_4^- + 6\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 3\text{H}_2\text{O} + 5\text{SO}_4^{2-}$
• Oxidation of Nitrite (NO$_2^-$) to Nitrate (NO$_3^-$): $5\text{NO}_2^- + 2\text{MnO}_4^- + 6\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 5\text{NO}_3^- + 3\text{H}_2\text{O}$

2. In Neutral or Faintly Alkaline Solutions:

• Oxidation of Iodide (I$^-$) to Iodate (IO$_3^-$): $2\text{MnO}_4^- + \text{H}_2\text{O} + \text{I}^- \rightarrow 2\text{MnO}_2\text{(s)} + 2\text{OH}^- + \text{IO}_3^-$
• Oxidation of Thiosulphate (S$_2$O$_3^{2-}$) to Sulfate (SO$_4^{2-}$): $8\text{MnO}_4^- + 3\text{S}_2\text{O}_3^{2-} + \text{H}_2\text{O} \rightarrow 8\text{MnO}_2\text{(s)} + 6\text{SO}_4^{2-} + 2\text{OH}^-$
• Oxidation of Manganous salts (Mn$^{2+}$) to MnO$_2$ (catalysed by ZnSO$_4$ or ZnO): $2\text{MnO}_4^- + 3\text{Mn}^{2+} + 2\text{H}_2\text{O} \rightarrow 5\text{MnO}_2\text{(s)} + 4\text{H}^+$

Note: Permanganate titrations cannot be performed accurately in the presence of hydrochloric acid (HCl) because HCl is oxidised to chlorine gas by permanganate.

Uses: Beyond its extensive use in analytical chemistry for titrations, potassium permanganate is a versatile oxidising agent in preparative organic chemistry. Its strong oxidising power is also utilised for bleaching materials like wool, cotton, silk, and other textile fibres, and for decolorising oils.

The Lanthanoids

The f-block of the periodic table is divided into two series: the Lanthanoids and the Actinoids. The lanthanoids comprise the fourteen elements that follow Lanthanum (La) in the periodic table (Cerium, Ce, to Lutetium, Lu).

Although Lanthanum (Z=57) is formally a d-block element, its properties are very similar to those of the lanthanoids. Therefore, it is conventionally discussed along with the lanthanoids, and the general symbol Ln is often used to represent any element in this series, including Lanthanum.

Lanthanoids show greater similarity among themselves compared to the variation seen among elements in a typical transition series. This is mainly because their characteristic properties are governed by the filling of the deeply buried 4f orbitals, while the outer electron configuration (5d$^x$ 6s$^2$) remains relatively constant.

They predominantly exhibit a single stable oxidation state (+3), which allows for the study of how small, progressive changes in size and nuclear charge affect the properties of otherwise similar elements.

Electronic Configurations

Atoms of lanthanoid elements generally have a 6s$^2$ electron configuration in their outermost shell, while the occupancy of the inner 4f and occasionally 5d levels varies.

The most stable oxidation state for all lanthanoids is +3. The electronic configurations of the tripositive ions (Ln$^{3+}$) are primarily of the form 4f$^n$, where n increases from 1 (for Ce$^{3+}$) to 14 (for Lu$^{3+}$) with increasing atomic number.

Atomic Number	Name	Symbol	Electronic configuration (M)	Electronic configuration (Ln$^{2+}$)	Electronic configuration (Ln$^{3+}$)	Electronic configuration (Ln$^{4+}$)
57	Lanthanum	La	5d$^1$6s$^2$	5d$^1$	4f$^0$
58	Cerium	Ce	4f$^1$5d$^1$6s$^2$	4f$^2$	4f$^1$	4f$^0$
59	Praseodymium	Pr	4f$^3$6s$^2$	4f$^3$	4f$^2$	4f$^1$
60	Neodymium	Nd	4f$^4$6s$^2$	4f$^4$	4f$^3$	4f$^2$
61	Promethium	Pm	4f$^5$6s$^2$	4f$^5$	4f$^4$
62	Samarium	Sm	4f$^6$6s$^2$	4f$^6$	4f$^5$
63	Europium	Eu	4f$^7$6s$^2$	4f$^7$	4f$^6$
64	Gadolinium	Gd	4f$^7$5d$^1$6s$^2$	4f$^7$5d$^1$	4f$^7$
65	Terbium	Tb	4f$^9$6s$^2$	4f$^9$	4f$^8$	4f$^7$
66	Dysprosium	Dy	4f$^{10}$6s$^2$	4f$^{10}$	4f$^9$	4f$^8$
67	Holmium	Ho	4f$^{11}$6s$^2$	4f$^{11}$	4f$^{10}$
68	Erbium	Er	4f$^{12}$6s$^2$	4f$^{12}$	4f$^{11}$
69	Thulium	Tm	4f$^{13}$6s$^2$	4f$^{13}$	4f$^{12}$
70	Ytterbium	Yb	4f$^{14}$6s$^2$	4f$^{14}$	4f$^{13}$
71	Lutetium	Lu	4f$^{14}$5d$^1$6s$^2$	4f$^{14}$5d$^1$	4f$^{14}$

Atomic And Ionic Sizes

A defining characteristic of lanthanoids is the overall decrease in atomic and ionic radii from Lanthanum (Z=57) to Lutetium (Z=71). This phenomenon is known as the Lanthanoid contraction.

The decrease in atomic radii (derived from metallic structures) is not entirely smooth, but the decrease in the radii of the tripositive ions (Ln$^{3+}$) is more regular.

The lanthanoid contraction is caused by the imperfect shielding of the nuclear charge by the 4f electrons. As the atomic number increases across the series, the nuclear charge increases by one unit at each step, and an electron is added to a 4f orbital. Electrons within the same subshell do not shield each other very effectively. Specifically, the shielding provided by one 4f electron for another electron is less effective than that of a d electron. This results in a gradual increase in the effective nuclear charge experienced by the outer electrons, pulling the electron cloud closer to the nucleus and causing a contraction in size.

Atomic Number	Name	Symbol	Radii/pm (Ln)	Radii/pm (Ln$^{3+}$)
57	Lanthanum	La	187	106
58	Cerium	Ce	183	103
59	Praseodymium	Pr	182	101
60	Neodymium	Nd	181	99
61	Promethium	Pm	181	98
62	Samarium	Sm	180	96
63	Europium	Eu	199	95
64	Gadolinium	Gd	180	94
65	Terbium	Tb	178	92
66	Dysprosium	Dy	177	91
67	Holmium	Ho	176	89
68	Erbium	Er	175	88
69	Thulium	Tm	174	87
70	Ytterbium	Yb	173	86
71	Lutetium	Lu	–	–

(The image would show a graph depicting the decrease in ionic radii of Ln$^{3+}$ ions from La to Lu).

The lanthanoid contraction has significant implications for the chemistry of the elements that follow the lanthanoids in the periodic table, specifically the third (5d) transition series. The cumulative contraction causes the atomic radii of the 5d transition metals (from Hafnium, Hf, onwards) to be very similar to those of their corresponding elements in the 4d series. For example, Zr (4d, 160 pm) and Hf (5d, 159 pm) have nearly identical radii due to the lanthanoid contraction. This similarity in size accounts for their natural co-occurrence and the difficulty encountered in separating them.

Oxidation States

The most common and stable oxidation state for lanthanoids is +3. However, some lanthanoids can occasionally exhibit +2 and +4 oxidation states, particularly in solid compounds or specific solutions.

These variations from the typical +3 state are often linked to the extra stability associated with achieving empty (4f$^0$), half-filled (4f$^7$), or completely filled (4f$^{14}$) f subshells. This contributes to irregularities in trends like ionisation enthalpies and oxidation states.

• Cerium (Ce) commonly exhibits the +4 state (4f$^0$ configuration), which is favoured by its noble gas configuration. However, Ce(IV) is a strong oxidising agent (E° for Ce$^{4+}$/Ce$^{3+}$ is +1.74 V), readily reverting to the more stable +3 state (4f$^1$). Despite the positive potential suggesting it could oxidise water, the reaction is very slow, making Ce(IV) a useful analytical reagent.
• Praseodymium (Pr), Neodymium (Nd), Terbium (Tb), and Dysprosium (Dy) can also show the +4 state, primarily in their oxides (MO$_2$).
• Europium (Eu) forms the +2 ion (Eu$^{2+}$) by losing its two 6s electrons, resulting in a stable 4f$^7$ (half-filled) configuration. However, Eu$^{2+}$ is a strong reducing agent, readily oxidised to the +3 state (4f$^6$).
• Similarly, Ytterbium (Yb) can form the +2 ion (Yb$^{2+}$) with a stable 4f$^{14}$ (completely filled) configuration, making it a reductant.
• Terbium (Tb) in the +4 state (Tb$^{4+}$) has a 4f$^7$ (half-filled) configuration, making it an oxidant.
• Samarium (Sm) can exhibit both +2 and +3 oxidation states.

In essence, the availability of stable f$^0$, f$^7$, or f$^{14}$ configurations drives some lanthanoids to adopt +2 or +4 states besides their common +3 state, and these states often exhibit strong reducing or oxidising properties as they tend to return to the +3 state.

General Characteristics

Lanthanoids share many similar physical and chemical properties:

• All lanthanoids are silvery-white, soft metals. The hardness generally increases with increasing atomic number across the series, with Samarium (Sm) being relatively hard, like steel.
• Their melting points typically range between 1000 K and 1200 K, although Samarium (Sm) has an unusually high melting point of 1623 K.
• They possess characteristic metallic structures and are good conductors of heat and electricity.
• Properties like density show smooth variations across the series, with some exceptions for Eu and Yb, and occasionally Sm and Tm.

Many trivalent lanthanoid ions (Ln$^{3+}$) are coloured in both solid compounds and aqueous solutions, with the exception of La$^{3+}$ (4f$^0$) and Lu$^{3+}$ (4f$^{14}$) which are colourless. The colour is due to f-f electronic transitions, similar to d-d transitions in transition metals. However, the absorption bands for f-f transitions are characteristically narrow because the 4f orbitals are deeply shielded from the external environment.

Lanthanoid ions (except those with f$^0$ or f$^{14}$ configurations, like La$^{3+}$, Ce$^{4+}$, Yb$^{2+}$, Lu$^{3+}$) are paramagnetic due to the presence of unpaired electrons in the 4f orbitals.

The first ionisation enthalpies (IE1) of lanthanoids are around 600 kJ mol$^{-1}$, and the second (IE2) are about 1200 kJ mol$^{-1}$, comparable to Calcium. Variations in the third ionisation enthalpies (IE3) are influenced by the stability of empty, half-filled, or completely filled 4f orbitals, similar to the exchange energy effects in d-block elements. Lanthanum, Gadolinium (4f$^7$), and Lutetium (4f$^{14}$) show abnormally low third ionisation enthalpies, reflecting the stability of their resulting f$^0$, f$^7$, and f$^{14}$ configurations in the +3 state.

In terms of chemical behaviour, the earlier members of the lanthanoid series are quite reactive, similar to Calcium. As the atomic number increases, their reactivity decreases, and they behave more like Aluminium. The standard electrode potentials (E°) for the half-reaction Ln$^{3+}$(aq) + 3e$^-$ $\rightarrow$ Ln(s) are generally in the range of –2.2 V to –2.4 V, with Europium (Eu) being an exception at –2.0 V. This indicates they are quite electropositive metals.

General reactions of lanthanoids (Ln):

• React with Hydrogen (H$_2$) when gently heated to form hydrides.
• React with Carbon (C) at high temperatures (e.g., 2773 K) to form carbides with various formulas (Ln$_3$C, Ln$_2$C$_3$, LnC$_2$).
• React with Nitrogen (N$_2$) when heated to form nitrides (LnN).
• React with Water (H$_2$O) to liberate Hydrogen and form hydroxides: $2\text{Ln(s)} + 6\text{H}_2\text{O(l)} \rightarrow 2\text{Ln(OH)}_3\text{(aq)} + 3\text{H}_2\text{(g)}$
• Dissolve in dilute acids to liberate Hydrogen: $2\text{Ln(s)} + 6\text{H}^+\text{(aq)} \rightarrow 2\text{Ln}^{3+}\text{(aq)} + 3\text{H}_2\text{(g)}$
• Burn in Oxygen (O$_2$) to form oxides (Ln$_2$O$_3$).
• React with Sulfur (S) when heated to form sulfides (Ln$_2$S$_3$).
• Burn in Halogens (X$_2$) to form halides (LnX$_3$).

The hydroxides, Ln(OH)$_3$, are distinct compounds (not just hydrated oxides) and are basic in nature, similar to alkaline earth metal hydroxides.

(The image would depict a central 'Ln' symbol with arrows pointing outwards to indicate reactions with H$_2$, C, N$_2$, H$_2$O, acids, O$_2$, S, halogens, showing the resulting compounds).

The primary application of lanthanoids is in the production of alloy steels, particularly for plates and pipes. A notable alloy is mischmetall, which is composed of about 95% lanthanoid metals and 5% iron, with small amounts of other elements. Mischmetall is used in Magnesium-based alloys for items like bullets and shell casings, as well as for lighter flints.

Mixed oxides of lanthanoids serve as catalysts in processes such as petroleum cracking. Some individual lanthanoid oxides are used as phosphors, materials that emit light when energised, finding applications in television screens and fluorescent surfaces.

The Actinoids

The Actinoids form the second series of inner transition elements, comprising the fourteen elements that follow Actinium (Ac) in the periodic table (Thorium, Th, to Lawrencium, Lr).

Actinium (Z=89) is formally a d-block element, similar to Lanthanum for the lanthanoids. However, due to its close resemblance to the actinoids, it is typically included in discussions of this series.

A significant characteristic of actinoids is that all of them are radioactive. The earlier members of the series have relatively long half-lives, making them easier to study and handle in larger quantities. However, the later actinoids are highly unstable with very short half-lives, sometimes only a few minutes. These later elements can only be prepared and studied in extremely small quantities, which makes their chemistry more challenging to investigate thoroughly.

The chemistry of actinoids is considerably more complex than that of lanthanoids, largely because they exhibit a wider range of oxidation states.

Electronic Configurations

Actinoid elements are generally believed to have a 7s$^2$ electron configuration in their outermost shell, with variable occupancy of the inner 5f and 6d subshells.

The fourteen electrons are formally added to the 5f orbitals. While Thorium (Z=90) has a configuration [Rn] 6d$^2$ 7s$^2$ (or 5f$^0$ 6d$^2$ 7s$^2$), the filling of 5f orbitals effectively begins from Protactinium (Z=91) onwards and is completed at element 103, Lawrencium (Z=103).

Similar to lanthanoids, there are irregularities in the electronic configurations of some actinoids. These deviations are again related to the enhanced stability associated with the empty (5f$^0$), half-filled (5f$^7$), and completely filled (5f$^{14}$) 5f orbitals.

Examples of configurations showing these stabilities include Americium (Am), [Rn] 5f$^7$ 7s$^2$, and Curium (Cm), [Rn] 5f$^7$ 6d$^1$ 7s$^2$.

Although the 5f orbitals resemble the 4f orbitals in their angular wave functions, the 5f orbitals are less effectively shielded from the nucleus and protrude further outwards compared to the 4f orbitals. Consequently, the 5f electrons are less 'buried' and can participate in chemical bonding to a much greater extent than the 4f electrons.

Atomic Number	Name	Symbol	Electronic configuration (M)	Electronic configuration (M$^{3+}$)	Electronic configuration (M$^{4+}$)
89	Actinium	Ac	6d$^1$7s$^2$	5f$^0$
90	Thorium	Th	6d$^2$7s$^2$	5f$^1$	5f$^0$
91	Protactinium	Pa	5f$^2$6d$^1$7s$^2$	5f$^2$	5f$^1$
92	Uranium	U	5f$^3$6d$^1$7s$^2$	5f$^3$	5f$^2$
93	Neptunium	Np	5f$^4$6d$^1$7s$^2$	5f$^4$	5f$^3$
94	Plutonium	Pu	5f$^6$7s$^2$	5f$^5$	5f$^4$
95	Americium	Am	5f$^7$7s$^2$	5f$^6$	5f$^5$
96	Curium	Cm	5f$^7$6d$^1$7s$^2$	5f$^7$	5f$^6$
97	Berkelium	Bk	5f$^9$7s$^2$	5f$^8$	5f$^7$
98	Californium	Cf	5f$^{10}$7s$^2$	5f$^9$	5f$^8$
99	Einstenium	Es	5f$^{11}$7s$^2$	5f$^{10}$	5f$^9$
100	Fermium	Fm	5f$^{12}$7s$^2$	5f$^{11}$	5f$^{10}$
101	Mendelevium	Md	5f$^{13}$7s$^2$	5f$^{12}$	5f$^{11}$
102	Nobelium	No	5f$^{14}$7s$^2$	5f$^{13}$	5f$^{12}$
103	Lawrencium	Lr	5f$^{14}$6d$^1$7s$^2$	5f$^{14}$	5f$^{13}$

Ionic Sizes

Actinoids also exhibit a gradual decrease in the size of their atoms and M$^{3+}$ ions across the series, similar to the lanthanoids. This effect is known as the Actinoid contraction.

However, the contraction from one element to the next in the actinoid series is generally greater than in the lanthanoid series. This larger contraction is attributed to the even poorer shielding provided by the 5f electrons compared to the 4f electrons. The 5f electrons are less effective at shielding the increasing nuclear charge, leading to a stronger pull on the outer electrons and a more pronounced decrease in size across the series.

Atomic Number	Name	Symbol	Radii/pm (M$^{3+}$)	Radii/pm (M$^{4+}$)
89	Actinium	Ac	111
90	Thorium	Th		99
91	Protactinium	Pa		96
92	Uranium	U	103	93
93	Neptunium	Np	101	92
94	Plutonium	Pu	100	90
95	Americium	Am	99	89
96	Curium	Cm	99	88
97	Berkelium	Bk	98	87
98	Californium	Cf	98	86
99	Einstenium	Es	–	–
100	Fermium	Fm	–	–
101	Mendelevium	Md	–	–
102	Nobelium	No	–	–
103	Lawrencium	Lr	–	–

Oxidation States

Actinoids display a greater range and variability of oxidation states compared to lanthanoids. This is attributed, in part, to the relatively comparable energies of the 5f, 6d, and 7s electron shells, allowing more electrons to be involved in bonding.

	Ac	Th	Pa	U	Np	Pu	Am	Cm	Bk	Cf	Es	Fm	Md	No	Lr
Oxidation States	3	3, 4	3, 4, 5	3, 4, 5, 6	3, 4, 5, 6, 7	3, 4, 5, 6, 7	3, 4, 5, 6	3, 4	3, 4	3	3	3	3	3	3

The most common oxidation state for actinoids is +3, similar to lanthanoids. However, elements in the first half of the series frequently exhibit higher oxidation states.

For example, the maximum oxidation state increases from +4 in Thorium (Th) to +5 in Protactinium (Pa), +6 in Uranium (U), and +7 in Neptunium (Np) and Plutonium (Pu). This trend reverses in the later elements, with the maximum oxidation state decreasing.

While actinoids have more compounds in the +3 state than in the +4 state, both +3 and +4 ions of actinoids tend to undergo hydrolysis in aqueous solution.

Due to the uneven distribution and variation in oxidation states, particularly between the early and later actinoids, it is often less straightforward to systematically review their chemistry based solely on oxidation states compared to the more consistent +3 state of lanthanoids.

General Characteristics And Comparison With Lanthanoids

Actinoid metals generally have a silvery appearance. They exhibit a wider variety of crystal structures compared to lanthanoids. This structural variability is linked to the more significant irregularities in their metallic radii across the series.

Actinoids are highly reactive metals, especially when in a finely divided state. They react with boiling water to form a mixture of oxide and hydride, and readily combine with most non-metals at moderate temperatures.

Dilute hydrochloric acid attacks all actinoid metals. However, nitric acid has less effect on most actinoids, likely due to the formation of a protective oxide layer on the surface (passivation). Alkaline solutions generally do not react with actinoid metals.

The magnetic properties of actinoids are more complex than those of lanthanoids. While the general trend in magnetic susceptibility with the number of unpaired 5f electrons is somewhat similar to the lanthanoids, the actinoids typically show lower magnetic moment values.

Evidence suggests that the ionisation enthalpies of the early actinoids are lower than those of the early lanthanoids. This is expected because the 5f orbitals, when they begin to be filled, are less penetrating into the inner electron core than the 4f orbitals. This means the 5f electrons are more effectively shielded from the nucleus compared to the 4f electrons in corresponding lanthanoids.

Because the outer electrons in actinoids are less tightly bound, they are more readily available for participation in chemical bonding, contributing to the wider range of oxidation states observed.

Comparing actinoids and lanthanoids:

Characteristic	Lanthanoids	Actinoids
Electronic Configuration	[Xe] 4f$^{0-14}$ 5d$^{0-1}$ 6s$^2$	[Rn] 5f$^{0-14}$ 6d$^{0-1}$ 7s$^2$
Primary Oxidation State	+3 (most stable)	+3 (most common, but wider range of higher states exist)
Other Oxidation States	Occasionally +2 and +4 (linked to f$^0$, f$^7$, f$^{14}$ stability)	More common +4, +5, +6, +7, especially in early members (due to comparable energy levels of 5f, 6d, 7s)
Contraction	Lanthanoid contraction (gradual decrease in size)	Actinoid contraction (greater decrease in size per element due to poorer 5f shielding)
Shielding	4f electrons provide better shielding than 5f	5f electrons provide poorer shielding than 4f
Bonding Involvement	4f electrons are deeply buried, less involved in bonding beyond +3 state	5f electrons are less buried, more involved in bonding, contributing to higher oxidation states
Radioactivity	Mostly non-radioactive (except Pm)	All are radioactive
Reactivity	Generally reactive, similar to alkaline earth metals (early members) or Al (later members)	Highly reactive, particularly when finely divided
Chemical Similarity	More similar to each other, especially in +3 state	Show greater variability, similarity mainly among elements in the second half of the series

The behaviour similar to the close-knit properties of lanthanoids is more apparent in the second half of the actinoid series. However, even the early actinoids show similarities with each other and gradual property variations without necessarily changing oxidation state.

Both the lanthanoid and actinoid contractions have lasting effects on the sizes and properties of the elements that succeed them in their respective periods (5d and 6d transition series). The lanthanoid contraction is currently considered more significant in its impact because the chemistry of elements after the actinoids (transactinide elements) is much less understood due to their extreme radioactivity and short lifetimes.

Answer:

Answer:

Some Applications Of D- And F-Block Elements

D-block and f-block elements and their compounds have numerous important applications:

• Structural Materials: Iron and various types of steel (alloys of iron with transition metals like Cr, Mn, Ni, V, W, Mo) are indispensable in construction due to their strength and durability. Their production involves reducing iron oxides and adding alloying elements.
• Pigments and Materials: Titanium dioxide (TiO$_2$) is widely used as a white pigment in paints, paper, and plastics. Manganese dioxide (MnO$_2$) is used in dry battery cells. Zinc (Zn) and Nickel/Cadmium (Ni/Cd) are also essential for the battery industry.
• Coinage: Elements of Group 11, Copper (Cu), Silver (Ag), and Gold (Au), have historically been used as coinage metals. While modern coins may be alloys (like Cu/Ni) or coated metals (like copper-coated steel), these elements remain significant for their value and properties.
• Catalysis: Many transition metals and their compounds are vital catalysts in the chemical industry.
  • Vanadium(V) oxide (V$_2$O$_5$) catalyses the oxidation of sulfur dioxide (SO$_2$) in the Contact Process to produce sulfuric acid.
  • A catalyst system based on Titanium tetrachloride (TiCl$_4$) and Triethylaluminium (Al(CH$_3$)$_3$), known as the Ziegler-Natta catalyst, is used for manufacturing polyethylene (polythene).
  • Iron (Fe) catalysts are used in the Haber-Bosch process for the synthesis of ammonia (NH$_3$) from nitrogen and hydrogen.
  • Nickel (Ni) catalysts are commonly used in the hydrogenation of vegetable oils to solid fats.
  • Palladium(II) chloride (PdCl$_2$) is a catalyst in the Wacker process for the oxidation of ethene to ethanal (acetaldehyde).
  • Nickel complexes are employed in the polymerisation of alkynes and other organic compounds like benzene.
• Photography: The photographic industry relies on the light-sensitive properties of Silver bromide (AgBr).
• Nuclear Energy: Inner transition elements, particularly Actinoids like Thorium (Th), Protactinium (Pa), and Uranium (U), are crucial as sources of nuclear energy in modern power generation and other applications.

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer:

Answer: