Redox reactions are spontaneous when an element can successfully displace another element within a compound. And when we say the term "displace," that means to remove an element from its compound, thereby reducing it. So to displace an element means I'm actually reducing that element. And we're going to say here we utilize what's called an activity series chart to determine if an element can displace another element. Now what we're going to say here is we're going to say an element that is higher in the activity series will displace an element lower in the activity series chart. And what we need to also remember is that recall an oxidizing agent is reduced, and a reducing agent will be oxidized. So if we take a look here at this activity series chart, we have lithium, potassium, calcium, all the way down to gold down here. And what we can say here is up near lithium, that's where we have the strongest reducing agent, which means that we have the greatest propensity for oxidation. So we want to be oxidized a lot. And up here is where we have the weakest, by default, the weakest oxidizing agent, which means reduction does not want to happen up here. Okay. So we're going to say at this end of the activity series chart, this is the greatest tendency to lose electrons. If we do the opposite end, down here we have the weakest reducing agent, which means that we have the least likely to be oxidized. So oxidation is very weak here. And down here we have the strongest oxidizing agent, which means reduction really wants to happen down here. So down here on the activity series chart is where we have the greatest tendency to gain electrons. So just remember, when looking at this activity series chart, an element that's higher up will displace another element that's below it. So if we had sodium, it could displace zinc from another compound. Why? Because sodium is higher up than zinc is.
Spontaneous Redox Reactions - Online Tutor, Practice Problems & Exam Prep
Spontaneous Redox Reactions occur when an element displaces another element within a compound.
Activity Series Chart
Spontaneous Redox Reactions Concept 1
Video transcript
Spontaneous Redox Reactions Example 1
Video transcript
Based on the activity series chart, determine if the following reaction represents a spontaneous redox reaction. Basically, we're looking to see if calcium can displace this silver. And we know that we're trying to compare those two because they're the elements on the activity series chart mentioned earlier. Now, take a look. Is calcium higher up than silver? If it is, then we can successfully displace the silver, essentially kicking it out. So, if you look, calcium is indeed higher up. Right? So, step 1, we locate the monoatomic element on the activity series chart. We found out where both are located. And remember, step 2, if it is higher on the activity series chart, it will displace the element within the nearby compound. So, silver is going to get kicked out, metals by themselves exist as solids at room temperature except for mercury. Plus, now calcium and chlorine are going to combine together to give us our new ionic compound. Remember, calcium is in group 2A, so it has a 2+ charge. Chlorine is in group 7A, so it has a 1− charge. We crisscross the numbers to give us our ionic compound, which would come out to be CaCl2, and here we can make it a solid if we want. The states are not that important. What's important is that calcium can successfully displace silver within this redox reaction, and as a result, it is spontaneous. Now, if we wanted to balance this, we'd say we have 2 chlorines here and only 1 here. So, I put a 2 here, and then as a consequence, I'd also have to put a 2 here for this silver. So these would be our coefficients for our balanced redox reaction: 1:2:2:1 as the coefficients.
Which element is the best reducing agent?
Determine whether which of the following redox reactions will occur spontaneously in the forward direction?
a) Ni(s) + Zn2+(aq) → Ni2+(aq) + Zn(s)
b) Fe(s) + Pb4+(aq) → Fe2+(aq) + Pb(s)
c) Al(s) + Ag+(aq) → Al3+(aq) + Ag(s)
d) Pb(s) + Mn2+ (aq) → Pb2+(aq) + Mn(s)
Suppose you wanted to cause Ni2+ ions to come out of solution as solid Ni. Which metal could you use to accomplish this?