The blood alcohol \(\left(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{OH}\right)\) level can be determined by titrating a sample of blood plasma with an acidic potassium dichromate solution, resulting in the production of \(\mathrm{Cr}^{3+}(a q)\) and carbon dioxide. The reaction can be monitored because the dichromate ion \(\left(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\right)\) is orange in solution, and the \(\mathrm{Cr}^{3+}\) ion is green. The balanced equation is \(16 \mathrm{H}^{+}(a q)+2 \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(a q) \longrightarrow\) $$ 4 \mathrm{Cr}^{3+}(a q)+2 \mathrm{CO}_{2}(g)+11 \mathrm{H}_{2} \mathrm{O}(l) $$ This reaction is an oxidation-reduction reaction. What species is reduced, and what species is oxidized? How many electrons are transferred in the balanced equation above?

Step by step solution

01

Determining the oxidation states of all atoms in each compound

To find the oxidized and reduced species in the reaction, we need to determine the oxidation states of all atoms in each compound. The oxidation states for the atoms in the given compounds are: \(H^+\): +1 \(Cr_2O_7^{2-}\): \(Cr\) - +6, \(O\) - -2 \(C_2H_5OH\): \(C\) - -2, \(H\) - +1, \(O\) - -2 \(Cr^{3+}\): +3 \(CO_2\): \(C\) - +4, \(O\) - -2 \(H_2O\): \(H\) - +1, \(O\) - -2

02

Identifying the species undergoing reduction and oxidation

Now, compare the oxidation states of each species before and after the reaction. We look for which species gained electrons (reduction) and which species lost electrons (oxidation): \(- Cr_2O_7^{2-}\): \(Cr\) changes from +6 to +3, meaning it gains electrons (since 2 chromium atoms are there, 6 electrons are gained in total). \(- C_2H_5OH\): \(C\) changes from -2 to +4, meaning it loses electrons (since 2 carbon atoms are there, 6 electrons are lost in total). The species that undergoes reduction is \(Cr_2O_7^{2-}\), and the species that undergoes oxidation is \(C_2H_5OH\).

03

Determining the number of electrons transferred

To find the number of electrons transferred, we look at the change in oxidation states between the reactants and products for the oxidized and reduced species: \(- Cr_2O_7^{2-}\) gains 6 electrons (from +6 to +3 for each \(Cr\) atom) \(- C_2H_5OH\) loses 6 electrons (from -2 to +4 for each \(C\) atom) So, a total of 6 electrons are transferred in this redox reaction. Both the oxidized and reduced species and the number of electrons transferred have been determined. The species reduced is \(Cr_2O_7^{2-}\) and the species oxidized is \(C_2H_5OH\). A total of 6 electrons are transferred in the balanced equation.

Unlock Step-by-Step Solutions & Ace Your Exams!

• Full Textbook Solutions

  Get detailed explanations and key concepts

• Unlimited Al creation

  Al flashcards, explanations, exams and more...

• Ads-free access

  To over 500 millions flashcards

• Money-back guarantee

  We refund you if you fail your exam.

Start your free trial

Over 30 million students worldwide already upgrade their learning with Vaia!