Amphoteric character of V2O5

Question

In my NCERT textbook, in the chapter 'The d- and f-block elements', the last paragraph under oxides and oxoanions of metals goes like this:

$(...)$ $\ce{V2O5}$ is, however, amphoteric though mainly acidic and it gives $\ce{VO4^3-} \space\text{&}\space\ce{VO2+}$ salts (...)

What do they actually mean by the phrase amphoteric though mainly acidic?

Wikipedia clearly states that,

Vanadium(V) oxide (vanadia) is the inorganic compound with the formula $\ce{V2O5}$. Commonly known as vanadium pentoxide, it is a brown/yellow solid, although when freshly precipitated from an aqueous solution, its colour is deep orange. Because of its high oxidation state, it is both an amphoteric oxide and an oxidizing agent. From the industrial perspective, it is the most important compound of vanadium, being the principal precursor to alloys of vanadium, and is a widely used industrial catalyst.

$\ce{V2O5}$ has more of an acidic character than a basic character, doesn't it? How does amphotericity exactly work in practice for vanadium (V) oxoions? Is it difficult to make $\ce{VO2+}$?

Answer 1

I will attempt to answer the question and will focus on $\ce{V2O5}$'s acidic character. I will first sum up what Poutnik said:

• The phrase "amphoteric though mainly acidic" means the species has more of an acidic character than a basic character.
• the rule of thumb is that with increasing metal oxidation state, the oxide basicity gets weaker and acidity stronger. See: Why is Cr2O3 amphoteric but CrO not?

Now focusing on $\ce{V2O5}$. Yes, it is amphoteric but it has a more acidic character because:

1. It dissolves slightly in water to give a pale yellow, acidic solution. It reacts with strong non-reducing acids to form solutions containing dioxovanadium(V) centers. It is also soluble in strong acids (like $\ce{H2SO4}$), where the pervanadyl ion is observed:

$$\ce{V2O5 + 2 H+ -> 2 VO2+ + H2O}$$

2. It can be a strong oxidant. It react with $\ce{HCl}$ where the $\ce{V(V)}$ is reduced to $\ce{V(IV)}$ with evolution of chlorine:

$$\ce{V2O5 + 6HCl -> 2VOCl2 + Cl2 + 3H2O}$$

3. It tends to react with alkalis to form various vanadate species. The predominant species depends on pH and concentration of alkali. The more pH is lowered (acidifying the solution), the vanadate ion is more protonated and it forms protonated oxovanadate ions (more broadly isopolyanions). Simultaneously, the species generated tends to polymerize. These are called polyoxovanadates species. The reaction scheme is very complicated but it has been explained in the Wikipedia article:

Dissolution of vanadium pentoxide in strongly basic aqueous solution gives the colourless $\ce{VO3^4-}$ ion. On acidification, this solution's colour gradually darkens from orange to red at around pH 7. Brown hydrated $\ce{V2O5}$ precipitates around pH 2, redissolving to form a light yellow solution containing the $\ce{[VO2(H2O)4]+}$ ion. The number and identity of the oxyanions that exist between pH 13 and 2 depend on pH as well as concentration. For example, the protonation of vanadate initiates a series of condensations to produce polyoxovanadate ions:

pH 9–12: $\ce{HVO4^2−, V2O7^4-}$
pH 4–9: $\ce{H2VO4-, V4O12^4-, HV10O28^5-}$
pH 2–4: $\ce{H3VO4, H2V10O28^4−}$

The overall scheme can be thus summed up as below:

$$\ce{\underset{colorless}{[VO4]^3-} ->[pH-12] \underset{colorless}{[VO3.OH]^2-} ->[pH-10] \underset{colorless}{[V2O6.OH]^3-}->[pH-9] \underset{orange}{[V3O9]^3-} ->[pH-7] \underset{red}{[V5O14]^3-} ->[pH-6.5] \underset{brown ppt.}{[V2O5].(H2O)_n} ->[pH-2.2] [V10O28]^6- ->[pH<1] \underset{pale-yellow}{[VO2]^+}}$$

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Closing points:

1. If we move down the group, $\ce{Nb2O5}$ and $\ce{Ta2O5}$ are technically amphoteric but they are too unreactive to be considered amphoteric. They are not attacked by acids, except $\ce{HF}$ where they form fluoro complexes. They form niobates and tantales when they are reacted (more specifically fused) with alkali. They do have weak acidic properties and the niobates/tantalates are decomposed by weak acids or $\ce{CO2}$
2. Group 4 elements also follow similar trends where $\ce{TiO2}$ is more acidic than basic and form titanates like $\ce{Na2TiO3}$ and $\ce{Na2Ti2O5}$. $\ce{ZrO2}$ and $\ce{HfO2}$ are more basic than acidic (but they are rather unreactive).