Chem 30 Unit 1: REDOX - Building A Redox Table
Given experimental evidence, it is possible to make ones own redox table. This is an important skill to develop: you will need to use it in a lab later this week and it almost always gets tested on the big D.
In this hand out, you will see 3 worked problems. I’ve used one method for the first two problems, and a second method for the last one. Use whichever one you find most helpful. After the worked problems, you have some homework problems to work on.
Question:
Solution:
Step 1: Let’s look at Test 1. Write out a net ionic reaction.
CoCl2(aq) + Zn(s) Co(s) + ZnCl2(aq)
Co2+(aq) + 2Cl- (aq) + Zn(s)Co(s) + Zn2+(aq) + 2Cl-(aq)
Co2+(aq) + Zn(s) Co(s) + Zn2+(aq)
Step 2: Identify the reducing agent (reducing agents are oxidized, so they will lose electrons) and the oxidizing agent (oxidizing agents are reduced, so they will gain electrons).
Co2+(aq) + Zn(s) Co(s) + Zn2+(aq)
OA RA
Step 3: Recall that for a reaction to take place spontaneously, the OA must be above the RA. Since this reaction took place in the lab spontaneously, place cobalt higher than zinc in the table, like below.
Co2+(aq)
+
Zn(s)
Step 4: Repeat steps 1-3 with Test 2.
Mg(NO3)2(aq) + Zn(s) Mg(s) + Zn(NO3)2(aq)
Mg2+(aq) + Zn(s) Mg(s) + Zn2+(aq)
RA OA
Zn(s)
+
Mg2+(aq)
Step 5: You now see that cobalt is above zinc and that zinc is above magnesium. You can now complete the table by writing each half-reaction in the appropriate order.
Co2+(aq) + 2e- Co(s)
Zn2+(aq) + 2e- Zn(s)
Mg2+(aq) + 2e- Mg(s)
You can check your work by referencing the redox table in the data booklet. Hey, you could even copy it directly…but if you did, you’d never learn anything about redox. So on second thought, don’t do that.
Here’s another worked problem. This one has more tests in it, but the steps are the same.
Step 1: Write out the net ionic reactions. Step 2: Label RA and OA. Step 3: decide whether reaction is spontaneous and use the spontaneity rule to order the metals.
Ra(s) + Al3+ (aq) Ra2+ (aq) + Al(s)
RA OA
Al3+(aq)
+
Ra(s)
Al(s) + Co3+(aq) Al3+(aq) + Co2+(aq)
RA OA
Co3+(aq)
+
Al(s)
From the first two tests, we know aluminum is above radium, and that cobalt is above aluminum. The last test tells us that potassium will not react with any metals given. I.e., potassium can not take electrons from either metal to turn them into ions. It is therefore the weakest OA and must go on the bottom of the list.
Step 5: Write out the table.
Co3+(aq) + 3e- Co(s)
Al3+(aq) + 3e- Al(s)
Ra2+(aq) + 2e- Ra(s)
K+(aq) + e- K(s)
You can also build a redox table by remembering the trends present in these tables. The strongest oxidizing agents are always on the top left. A strong oxidizing agent is one that is really good at being reduced, or at gaining electrons. You can now place the metal which gains electrons best on the top of the table and fill down from there.
Solution:
Test 1: R2+(aq)is able to take two electrons from X(s). Therefore, R is a better oxidizing agent than X. Place R above X in the final table.
Test 2: Q2+(aq) could not take electrons from X(s). Therefore, Q must be below X on the table.
Test 3: X2+(aq) is able to take two electrons from M(s). Therefore, X must be above M on the table.
Test 4: Q2+(aq) could not take electrons from M(s). Therefore, M must be above Q on the table.
R2+(aq) + 2e- R(s)
X2+(aq) + 2e- X(s)
M2+(aq) + 2e- M(s)
Q2+(aq) + 2e- Q(s)
Practice Problems:
1. The results from six lab tests are recorded below. Write the net ionic reactions for each test and use them to build a redox table. The nitrate ion, NO3-, is a spectator ion.
Test 1: A strip of cobalt, Co(s), was placed in a solution of 1.0 mol/L Pb(NO3)2(aq). The Co(s) started to dissolve to Co2+(aq) and crystals of lead started to form.
Test 2: A bar of lead, Pb(s), began to dissolve in 1.0 mol/L Cu(NO3)2(aq), forcing copper out of solution.
Test 3: A rod of copper, Cu(s), was placed in a solution of 1.0 mol/L Co(NO3)2(aq). No reaction occurred.
Test 4: A rod of copper, Cu(s), was placed in a 1.0 mol/L solution of iodine, I2(l). Copper began to dissolve.
Test 5: No reaction occurred when a strip of silver, Ag(s), was placed in a 1.0 mol/L solution of I2(l).
Test 6:Silver crystals formed when a cobalt strip, Co(s), was immersed in a 1.0 mol/L AgNO3(aq).
2. For the reduction table established in question 1, name:
a) the best oxidizing agent
b) the worst oxidizing agent
c) the best reducing agent
d) the worst reducing agent
3. Using the hypothetical reactions below, establish a redox table.
Test 1: X3+(aq) + D(s) D3+(aq) + X(s)
Test 2: X3+(aq) + T(s) no reaction
Test 3: T3+(aq) + Z(s) Z3+(aq) + T(s)
Test 4: D3+(aq) + Z(s) Z3+(aq) + D(s)
4. Use the standard reduction table (in data booklet) to answer the following.
a) Will cobalt, Co(s) dissolve in 1.0 mol hydrochloric acid?
b) Can zinc, Zn(s) reduce the copper (II) ion, Cu2+(aq)?
c) Can an aqueous solution of chlorine, Cl2(g), oxidize zinc, Zn(s)?
5. For each of the chemical changes below, state whether the initial substance was oxidized or reduced.
a) Ba(s) is changed to Ba2+(aq).
b) Br2(l) is changed to 2Br-(aq)
c) Fe2+(aq) is changed to Fe3+(aq)
d) Sn4+(aq) is changed to Sn2+(aq)
e) Cr3+(aq) is changed to Cr(s)
f) Te2-(aq) is changed to Te(s)
6. When chlorine gas is passed over hot cesium, the two elements combine vigorously to produce a white solid. The exothermic reaction releases energy as a brilliant flash of light.
a) What is the compound product?
b) Write the net ionic reaction.
c) Write the two half-reactions.
d) Which is the oxidizing agent and which is the reducing agent?
7. Five spontaneous reactions are given below. Use these reactions to build a reduction table of the five elements involved.
a) Br2(l) + 2Fe2+(aq) 2Br-(aq) + 2Fe3+(aq)
b) 2Fe3+(aq) + 2I-(aq) I2(aq) + 2Fe2+(aq)
c) Cl2(g) + 2Br-(aq) 2Cl-(aq) + Br2(l)
d) Hg(l) + Br2(l) 2Br-(aq) + Hg2+(aq)
e) Hg2+(aq) + 2Fe2+(aq) 2Fe3+(aq) + Hg(l)
8. For the reduction table established in #7, which substance is the best oxidizing agent and which substance is the best reducing agent?