Write a balanced equation for these redox reactions. (a) The...

Write a balanced equation for these redox reactions. (a) The oxidation of nitrite ion to nitrate ion by permanganate ion, $\mathrm{MnO}_{4}^{-}$, in acidic solution $\left(\mathrm{MnO}_{4}^{-}\right.$ ion is reduced to $\mathrm{Mn}^{2+}$ ). (b) The reaction of manganese(II) ion and permanganate ion in basic solution to form solid manganese dioxide. (c) The oxidation of ethanol by dichromate ion in acidic solution, producing chromium(III) ion, acetaldehyde $\left(\mathrm{CH}_{3} \mathrm{CHO}\right),$ and water as products.

Write a balanced equation for these redox reactions.
(a) The oxidation of nitrite ion to nitrate ion by permanganate ion, $\mathrm{MnO}_{4}^{-}$, in acidic solution $\left(\mathrm{MnO}_{4}^{-}\right.$ ion is reduced to $\mathrm{Mn}^{2+}$ ).
(b) The reaction of manganese(II) ion and permanganate ion in basic solution to form solid manganese dioxide.
(c) The oxidation of ethanol by dichromate ion in acidic solution, producing chromium(III) ion, acetaldehyde $\left(\mathrm{CH}_{3} \mathrm{CHO}\right),$ and water as products.

Key Concepts

Redox Reactions

Redox reactions involve a transfer of electrons between substances. In these reactions, one species is oxidized (loses electrons) while another is reduced (gains electrons), leading to changes in their oxidation states. Understanding how electrons are transferred is central to writing and balancing redox reactions.

Half-Reaction Method

The half-reaction method is a systematic technique used to balance redox reactions. It involves breaking down the overall reaction into two separate half-reactions – one for oxidation and one for reduction. Each half-reaction is balanced first for mass and charge, and then they are recombined to give the fully balanced overall reaction.

Balancing in Acidic and Basic Media

The medium of the reaction (acidic or basic) plays a crucial role in balancing redox equations. In acidic conditions, H+ ions and water molecules are used to balance hydrogen and oxygen atoms, whereas in basic conditions, OH- ions and water molecules are utilized. This adjustment ensures that both mass and charge are balanced in the final equation.

Stoichiometry and Charge Balance

Balancing redox reactions requires careful attention to stoichiometry – ensuring that the number of atoms for each element and the overall charge is the same on both sides of the equation. This includes balancing complex reactions with multiple reactants and products by correctly accounting for all electron transfers and the distribution of ions.

Transcript

00:02 So we will try to write the redock equation and balance it according to the statement that is given.

00:18 Number one is nitrite to nitrate and 02 minus and per manganate to manganese.

00:43 Mn -04 minus and nitrate and 2, manganese 2 plus.

01:12 And this needs to be in acidic condition, which means we need hydrogen ion.

01:22 So we checked the oxidation state.

01:26 Nitrogen is from positive 3 to positive 5 so it increased by 2 and then manganese from positive 7 to positive 2 drops by 5 so in order to balance we need to multiply by 5 for the nitrogen multiply by 2 for manganese.

02:24 That way they both get 10.

02:28 Okay, so nitrogen nitrite has by as the coefficient and so it's nitrate and then permanganate and manganate and manganese 2 plus has a 2.

02:47 Now we're going to decide the hydrogen ion that we need to add.

02:54 Okay, so we'll check the charge on the left side is 5 minus plus 2 minus, there's 7 minus on the left side.

03:08 And on the right side is 5 plus and 5 minus 3 plus.

03:22 So it's negative 1.

03:28 And in order to balance it, we're going to need to add 6 hydrogen ion.

03:39 That way, and then 6 plus, that is equal to 1 minus.

03:48 Okay.

03:49 And which means on the right side, we're going to add 3 water molecules.

04:01 We'll check the hydrogen.

04:04 We have 6, 6, the oxygen will have 10 here, plus 2 times 4 is 8, so it's 18 on the left side.

04:14 On the right side, 13 plus 3 is 18, so it's balanced.

04:20 So we add 6 hydrogen ion here and 3 water molecules on the right side.

04:31 The next one we have is manganese.

04:39 Mn 2 plus and permangonates producing manganase 4 plus oxide.

05:10 This is a manganese at different oxidation state.

05:20 Either losing electron and gaining electron to produce manganese 4 plus.

05:29 Okay.

05:30 So one we have is from positive 2 to positive 4.

05:41 The charges increase by 2.

05:45 Okay.

05:46 The other one is from permanganate, which is positive 7 drops to positive 4.

05:56 So it goes down by 3.

05:59 So that way we know we need to multiply by three.

06:04 This one we need to multiply by two so that they both have six electrons.

06:10 So multiplied by three is this one.

06:18 Okay, multiplied by two is the permanganate.

06:23 And then on the right side, then we're going to have five magnesium dioxide.

06:36 And then this reaction happens in the basic condition, which means we're gonna add a hydroxide ion.

06:47 So let's check the charges on both sides.

06:50 This one is 6 plus and 2 minus, so an overall is 4 plus.

07:00 And on the right side is 0, which means on the right left side, we're going to need to add the hydroxide ion to balance the 4 plus charges.

07:12 So we're going to add 4 hydroxide ion...

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