Redox reactions are spontaneous when an element can successfully displace another element within a compound. When we say the term 'displace', that means to remove an element from its compound and thereby reducing it. So to displace an element means I'm actually reducing that element. We utilize what's called an activity series chart to determine if an element can displace another element. We say an element that is higher in the activity series will displace an element lower in the activity series chart. It's also important to recall that an oxidizing agent is reduced and a reducing agent will be oxidized. If we take a look here at this activity series chart, we have elements like lithium, potassium, calcium, all the way down to gold. At the top near lithium, we have the strongest reducing agent, indicating the greatest propensity for oxidation; essentially, these elements want to be oxidized a lot. Conversely, this also means that at the top, we have the weakest oxidizing agent, implying that reduction does not want to happen. At this end of the activity series chart, we observe the greatest tendency to lose electrons. On the opposite end, we have the weakest reducing agent, making it least likely to be oxidized, meaning oxidation is very weak here. At the lower end, we have the strongest oxidizing agent, indicating that reduction really wants to happen and here we see the greatest tendency to gain electrons. Remember, when looking at this activity series chart, an element higher up will displace another element that's below it. For example, if we had sodium, it could displace zinc from another compound because sodium is higher up than zinc.
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Spontaneous Redox Reactions - Online Tutor, Practice Problems & Exam Prep
Redox reactions occur spontaneously when a higher activity series element displaces a lower one, effectively reducing it. The activity series chart ranks elements by their tendency to lose electrons, with lithium and potassium as strong reducing agents. An oxidizing agent is reduced, while a reducing agent is oxidized. Understanding this hierarchy is crucial for predicting reaction outcomes, such as sodium displacing zinc due to its higher position in the activity series. This concept is foundational in electrochemistry and helps in analyzing oxidation-reduction reactions.
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. We're looking to see if calcium can displace silver. We need to compare the two because they're the elements on the activity series chart above. Look and see if calcium is higher than silver. It is, indicating calcium can successfully displace the silver, ejecting it from the compound. So step 1, we locate the monoatomic element on the activity series chart. We found out where both are located. Remember, step 2, if it is higher on the activity series chart, it will displace the element within the nearby compound. Thus, silver will be kicked out, as metals by themselves exist as solids at room temperature, except for mercury. Now, calcium and chlorine combine to form a new ionic compound. Calcium belongs to group 2a, so it has a 2+ charge, and chlorine belongs to group 7a, so it has a 1- charge. We crisscross these numbers to form the ionic compound, CaCl2, which we can denote as a solid. The states are not crucial here; what's important is that calcium can successfully displace silver in this redox reaction, making it spontaneous. To balance this, we note there are two chlorines here and only one here, so we put a 2 here, and as a consequence, we'd also put a 2 here for the silver. These are the coefficients for our balanced redox reaction: 1, 2, 2, 1.
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?