O Lvl Chem - Reducing Agent
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I'm really confused with whether Potassium or Lithium is a stronger reducing agent.
I've read about Standard Reduction Potential(althought I don't understand much) and it says that Lithium with the lowest reduction potential is the stronger reducing agent.However, according to the Reactivity Series, Potassium is more reactive than Lithium and thus is a stronger reducing agent.
Can anyone clarify this with me?Thanks in advance.
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Originally posted by Xavolon:
I'm really confused with whether Potassium or Lithium is a stronger reducing agent.
I've read about Standard Reduction Potential(althought I don't understand much) and it says that Lithium with the lowest reduction potential is the stronger reducing agent.However, according to the Reactivity Series, Potassium is more reactive than Lithium and thus is a stronger reducing agent.
Can anyone clarify this with me?Thanks in advance.
It so happens that this particular comparison you've brought up (or to be precise : Lithium), is problematic. Yes, K is indeed more reactive than Li (due to less endothermic 1st ionization energy of K due to greater distance between the +vely charged nucleus and valence shell, and greater shielding effect from greater no. of intermediate electron shells).
However, it is also true that the standard oxidation potential of Li is more positive than that of K. This may indeed be correctly interpeted as an apparent contradiction, at A levels. As it turns out, the reason for this contradiction has to do with the high charge density of the Li+ ion, resulting in extraordinarily strong ion-permanent dipole interactions with water, in the aqueous state. This phenomenon additionally favours the oxidation of Li to Li+ in the aqueous state (over K to K+), even though the 1st ionization energy of Li to Li+ is significantly more endothermic (compared with K to K+) in the gaseous state.
Therefore, for 'O' level purposes, it would suffice to simply state that because K is indeed more reactive than Li, hence K is a stronger reducing agent than Li. For 'A' level purposes, the high achieving student should be able to describe this apparent contradiction and explain it both ways.
Other than this particular problematic case of Li, you (both 'O' and 'A' level students) can use the Standard Redox Potentials accordingly, as per the example below :
Because a reducing agent is oxidized during the reaction, thus it is more appropriate to compare oxidation potentials (rather than reduction potentials). Since you're taking the reverse direction to these equations given, accordingly you should take the -ve sign for these potentials.
Eg. Based on the image below, since the reduction potential of Ag+ to Ag is +0.8V, hence the oxidation potential of Ag to Ag+ is -0.8V.
Let's put aside Lithium for the moment, and compare Na versus K. (to show you how you can use the image above to compare reducing and oxidizing strengths, other than Li).
Since the oxidation potential of Na to Na+ is +2.71V, while the oxidation potential of K to K+ is +2.92V, we see that K thus has a more positive potential (ie. greater potential = greater tendency = greater inclination = greater propensity) to be oxidized, compared to Na.
Therefore K is a stronger reducing agent than Na.
Now try it out for yourself, say, is bromine or iodine a stronger oxidizing agent? This time round, since oxidizing agents are being reduced, you should compare reduction potentials, not oxidation potentials.