Aluminum metal is oxidized in aqueous base with water serving as the oxidizing agent. Al(s) + H2O(1) → Al(OH)4- + H2(g) (basic solution) Write a balanced equation for this reaction in basic solution.
First of all, Aluminum is oxidized from 0 to +3 oxidation states. Hydrogen is reduced from +1 to 0. We need to balance the half-reactions. Remember to first balance the elements (that are not O or H); then balance the O's with H2O; then balance the H's with H+; then balance the charges with e-'s. Finally, make sure the electrons balance for each of the half-reactions. Then you add them up. 4 H2O + Al -- Al(OH)4- + 4H+ + 3 e- 2 e- + 2 H+ + H2O -- H2 + H2O I'll multiply the first equation by 2 and the second by 3 so we exchange the same number of electrons. 8 H2O + 2 Al -- 2 Al(OH)4- + 8 H+ + 6 e- 6 e- + 6H+ + 3H2O -- 3 H2 + 3H2O --------------------------------------... 8 H2O + 2 Al -- 2 Al(OH)4- + 2H+ + 3 H2 Since it's in a basic solution, we will hydrolyze the H+ 8 H2O + 2 Al -- 2 Al(OH)4- + 2H+ + 3 H2 2H+ + 2OH- -- 2 H2O --------------------------------------... 6 H2O + 2 Al + 2 OH- -- 2 Al(OH)4- + 3 H2
Aluminium is an alloy not a metal.
solutions b, c, and d - chlorine, bromine, and permanganate anion all characterize powerful oxidizers. In an elementary diatomic halogen molecule like chlorine or bromine, the oxidation state of each and every atom is 0. even as they react with molecular magnesium, as an social gathering (Mg, whose oxidation state is likewise 0), the chlorine atoms each and each and every get rid of an electron from the metal and for that reason have an oxidation state of -a million. they have been decreased, and are for that reason oxidizing brokers. Magnesium loses 2 electrons and for that reason has an oxidation state of +2. this is been oxidized (with techniques from an oxidizing agent) and is for that reason a reducing agent.
You need to add OH- ions on the left. 2Al + 2OH- + 6H2O ---- 2Al(OH4)- + 3H2
