Question

Consider the following balanced redox reaction (do not include state of matter in your answers):2CrO2−(aq) + 2H2O(l) + 6ClO−(aq) →2CrO42−(aq) + 3Cl2(g) + 4OH−(aq)(a) Which species is being oxidized? (b) Which species is being reduced? (c) Which species is the oxidizing agent? (d) Which species is the reducing agent? (e) From which species to which does electron transfer occur?

Answer 1

Answer:

E i think

Explanation:

Answer 2

Answer:

Explanation:

Hello,

At first, it is necessary to establish the oxidation states for all the species involved in the reaction:

$2(Cr^{+3}O_2^{-2})^{-}(aq) + 2H_2^{+1}O^{-2}(l) + 6(Cl^{+1}O^{-2})^{-1}(aq) -->2(Cr^{+6}O_4^{-2})^{-2}(aq) + 3Cl_2^0(g) + 4(O^{-2}H^{+1})^-(aq)$

In such a way, we answer to the questions as follows:

(a) Oxidized species is the one that has an increase in its oxidation state, this is chromium which passes from +3 to +6.

(b) Reduced species is the one that has a decrease in its oxidation state, this is chlorine which passes from +1 to 0.

(c) Oxidizing agent is the same reduced species, this is chlorine as it removes electrons.

(d) Reducing agent is the same oxidized species, this is chromium as it gains electrons.

(e) Electron transfer occur from chlorine to chromium as chlorine removes electrons and chromium gains them.

Best regards.