The increase in density from titanium (Z = 22) to copper (Z = 29) in the first series of transition elements can be explained by several factors:

1. Atomic Mass: As you move from titanium to copper in the periodic table, the atomic mass generally increases due to the addition of more protons, neutrons, and electrons. Since density is mass per unit volume, an increase in atomic mass tends to increase density.

2. Atomic Radius: While the atomic radius generally decreases across a period in the periodic table due to increasing effective nuclear charge, the increase in atomic mass across the transition metals offsets this effect to some extent. As you move from titanium to copper, the increase in atomic mass generally outweighs the decrease in atomic radius, contributing to the increase in density.

3. Crystal Structure: Transition metals typically have a close-packed crystal structure, which means that their atoms are densely packed together in a regular pattern. Changes in atomic size and mass can influence how tightly packed these atoms are, affecting the density of the material.

4. Electron Configuration: Transition metals have complex electron configurations, with electrons occupying different sublevels within the d-block. Changes in electron configuration can influence the interactions between atoms and hence affect the density of the material.

5. Transition Metals' Special Properties: Transition metals often exhibit unique properties such as high melting points, hardness, and metallic bonding characteristics, all of which can influence the density of the elements in this series.

Overall, the increase in density from titanium to copper in the first series of transition elements is a result of various interplaying factors including atomic mass, atomic radius, crystal structure, electron configuration, and special properties of transition metals.

(i) Transition elements generally form colored compounds:

The color exhibited by transition metal compounds arises from the d-d transition, which involves the movement of electrons between the d orbitals of the metal ions. Transition metals have partially filled d orbitals, which allow for the absorption of visible light. When light strikes a transition metal complex, it can promote an electron from a lower-energy d orbital to a higher-energy d orbital, resulting in the absorption of certain wavelengths of light and the reflection or transmission of others. The color observed depends on the energy difference between the d orbitals involved in the transition.

The intensity and nature of the color can be influenced by various factors such as the oxidation state of the metal ion, the ligands surrounding the metal ion, and the coordination geometry of the complex. Ligands with different electron-donating abilities can lead to different splitting patterns of the d orbitals, resulting in different absorption spectra and hence different colors.

(ii) Zinc is not regarded as a transition element:

Zinc is often not considered a transition element because it lacks partially filled d orbitals in its common oxidation states. In its most common oxidation state, +2, the 3d orbitals are completely filled, which means there are no available d electrons for d-d transitions to occur. Therefore, zinc typically forms colorless compounds.

Transition metals, by definition, have incompletely filled d orbitals in at least one oxidation state, which allows them to exhibit characteristic transition metal properties such as forming colored compounds and acting as catalysts. Since zinc does not fulfill this criterion, it is often excluded from the list of transition elements despite being located in the d-block of the periodic table.

(i) Copper (I) ion is not known in aqueous solution primarily because copper tends to exist in the +2 oxidation state in aqueous solutions. This is due to the relative stability of the Cu(II) oxidation state compared to Cu(I) in aqueous environments. The standard reduction potential for the Cu(II)/Cu(I) couple is higher than that for many other metal ions, making the Cu(II) state more stable in water. Additionally, Cu(II) ions readily hydrolyze in water, forming insoluble Cu(OH)₂, further reducing the concentration of Cu(I) ions in solution.

(ii) Actinoids exhibit a greater range of oxidation states than lanthanoids due to the presence of f-orbitals in their electron configurations. Actinoid elements have more extended series of f-orbitals available for electron configuration, leading to a greater variety of possible oxidation states. The lanthanoid series, on the other hand, have electrons filling 4f orbitals, which are relatively shielded from the outer environment by the 5s and 5p orbitals. As a result, lanthanoid elements generally exhibit fewer accessible oxidation states compared to actinoids. Additionally, the actinoid series is longer than the lanthanoid series, providing more elements with a greater variety of electron configurations and oxidation states.

Sure, let's break down each of these statements:

(i) Transition metals generally form colored compounds: Reason: The color of transition metal compounds arises due to the presence of partially filled d orbitals in the transition metal ions. When light interacts with these compounds, electrons in the d orbitals can absorb certain wavelengths of light, causing them to transition to higher energy levels. The absorbed wavelengths correspond to the complementary color of the one observed, resulting in the compound appearing colored. This phenomenon is known as d-d transition. The energy gap between the d orbitals varies depending on the metal ion and its oxidation state, leading to a wide range of colors observed in transition metal compounds.

(ii) Manganese exhibits the highest oxidation state of +7 among the 3d series of transition elements: Reason: Manganese, being a member of the 3d transition metal series, can exhibit multiple oxidation states due to the availability of its d orbitals for electron transfer. However, among the 3d series elements, manganese has the highest number of unpaired electrons available in its 3d orbitals, which allows it to achieve its highest oxidation state of +7. This occurs in compounds like potassium permanganate (KMnO4), where manganese is in the +7 oxidation state. The ability of manganese to access this high oxidation state is attributed to its electron configuration and its position within the periodic table.

These are interesting observations that can be explained by considering the electronic configurations and trends in oxidation states across transition metals.

(i) Cr2+ is reducing in nature while Mn3+ is an oxidizing agent: This can be explained by looking at the electronic configurations of Cr2+ and Mn3+.

• Cr2+ has an electronic configuration of [Ar] 3d4, where it has a half-filled d orbital. Half-filled orbitals have lower energy due to greater exchange energy, making it energetically favorable for Cr2+ to lose electrons and become Cr3+ in order to achieve a stable half-filled d orbital, thus acting as a reducing agent.

• On the other hand, Mn3+ has an electronic configuration of [Ar] 3d4, which is one electron short of achieving a stable half-filled d orbital. So, Mn3+ tends to gain an electron to achieve a stable half-filled d orbital, making it an oxidizing agent as it oxidizes other species by accepting electrons.

(ii) In a transition series of metals, the metal which exhibits the greatest number of oxidation states occurs in the middle of the series: This observation can be explained by considering the trends in the filling of d orbitals across the transition series.

• At the beginning of the transition series, elements have fewer d electrons available for oxidation, limiting the number of oxidation states they can exhibit.

• Toward the middle of the series, there's a peak in the number of oxidation states exhibited. This is because these elements have a balance between gaining and losing electrons, allowing them to exhibit a wider range of oxidation states.

• Towards the end of the series, the number of oxidation states generally decreases as elements have a higher tendency to gain electrons rather than lose them, leading to fewer oxidation states.

So, the middle of the transition series tends to have elements that can exhibit the greatest number of oxidation states due to the balance between gaining and losing electrons facilitated by their electronic configurations.