Chemistry, asked by RotilaMae23, 2 months ago

which species in each pair is a better reducing agent under standard-state conditions?
a) Na or Li
b) H2 or I2
c) Fe2+ or Ag
d) Br- or Co2+​

Answers

Answered by taara2307
0

Step 1 of 17

(a)

Identify the oxidation and reduction half-reaction by assigning the oxidation numbers.

Separate the unbalanced reaction into half-reactions.

Step 2 of 17

Balance each half-reaction for mass, excluding O and H.

Balance both reactions for O by adding.

Balance both half for H by adding.

Step 3 of 17

Balance the total charge of both half-reactions by adding electrons.

Multiply the half-reactions to make the numbers of electrons the same in both.

Add the half-reactions back together, cancelling electrons.

Therefore, the balanced equation is as follows:

Step 4 of 17

(b)

Identify the oxidation and reduction half-reaction by assigning the oxidation numbers.

Separate the unbalanced reaction into half-reactions.

Step 5 of 17

Balance each half-reaction for mass, excluding O and H.

Balance both reactions for O by adding.

Balance both half for H by adding

Step 6 of 17

Balance the total charge of both half-reactions by adding electrons.

Multiply the half-reactions to make the numbers of electrons the same in both.

Add the half-reactions back together, cancelling electrons.

Therefore, the balanced equation is as follows:

Step 7 of 17

(c)

Identify the oxidation and reduction half-reaction by assigning the oxidation numbers.

Separate the unbalanced reaction into half-reactions.

Step 8 of 17

Balance each half-reaction for mass, excluding O and H.

Balance both reactions for O by adding.

Balance both half for H by adding.

Step 9 of 17

Balance the total charge of both half-reactions by adding electrons.

Multiply the half-reactions to make the numbers of electrons the same in both.

Add the half-reactions back together, cancelling electrons.

Step 10 of 17

For each ion in the final equation, add one ion to each side of the equation, combining and ions to produce.

Therefore, the balanced equation is as follows:

Step 11 of 17

(d)

Identify the oxidation and reduction half-reaction by assigning the oxidation numbers.

Separate the unbalanced reaction into half-reactions.

Step 12 of 17

Balance each half-reaction for mass, excluding O and H.

Balance both reactions for O by adding.

Balance both half for H by adding.

Step 13 of 17

Balance the total charge of both the half reactions by adding electrons.

Multiply the half-reactions to make the numbers of electrons the same in both.

Add the half-reactions back together, cancelling electrons.

Step 14 of 17

For each ion in the final equation, add one ion to each side of the equation, combining and ions to produce.

Therefore, the balanced equation is as follows :]

Similar questions