Part C: Concept Checks

Transcript

0:00Introduction

Hello everyone! Okay, a small puzzle before we start. Of all the alkali metals, lithium holds its outer electron the most tightly. It has the highest ionisation energy in its group. So in water, lithium should be the most unwilling to give that electron away, right? And yet, lithium has the most negative E naught of them all. It is the strongest reducing agent in water! How can that be? Hold on to that question. Today: the concept checks of our electrochemistry problem bank, six short why questions. Questions like these come up often in JEE as statements to judge, so let's begin.

0:39Check 1: the sign of the anode

So, check one. On screen is the Daniell cell, a galvanic cell. The zinc rod on the left is the anode, marked with a minus sign, and the copper rod on the right is the cathode, marked with a plus sign. The circle marked V, between them, is a voltmeter. Okay? Now tell me: is the anode always the negative electrode? No, not always! In a galvanic cell like this one, yes, the anode is negative and the cathode is positive. The little arrows on top, marked e minus, show the electrons flowing through the wire, from the anode to the cathode. Right? But the word always is the trap. In an electrolytic cell, an external source drives the reaction, and the anode is joined to its positive terminal. So there, the anode is positive, and the cathode is negative. Clear? So what never changes? In both cells, oxidation takes place at the anode, and reduction at the cathode. You can see it at the bottom of the figure: at the anode, zinc gives up two electrons and becomes Z n two plus; at the cathode, Cu two plus takes two electrons and becomes copper. So name an electrode by its reaction, never by its sign.

1:55Check 2: intensive or extensive?

Got it? Check two. E cell, the E with the small word cell below it, is the cell potential. Suppose you double the whole cell reaction, every coefficient times two. Does E cell double too? No, it stays the same! E cell is intensive: it does not depend on how much reaction you write. Okay? Now delta G, the little triangle delta meaning the change in Gibbs energy. Delta G equals minus n F E, where n is the number of electrons transferred in the equation as written, and capital F is the Faraday constant. And delta G is extensive. Why? Look at n. Double the reaction, and twice as many electrons are transferred, so n doubles. Then delta G doubles too, but E stays the same. So never multiply E naught, the standard value with the small circle on top, by the coefficients. Students lose easy marks here!

2:54Check 3: why agar in the salt bridge?

Check three. Look at the figure again: the bent tube in the middle is the salt bridge. It holds K C l, potassium chloride, KNO three, potassium nitrate, or N H four N O three, ammonium nitrate, set in agar agar. In our figure, it is potassium nitrate. Okay? Now think. Why agar? Why not simply fill the tube with the salt solution? Because agar sets into a gel, like a jelly. The gel holds the electrolyte in place, so the two solutions cannot mix: zinc sulphate stays on the left, and copper sulphate on the right. Right? Now, a trap. If the gel stops the solutions from mixing, does it also stop the ions from moving? No, no! The ions still move through the gel. That is how the salt bridge carries the current and completes the circuit. It stops the solutions mixing, not the ions moving. Got it?

3:50Check 4: why KCl for the cell constant?

Check four: the conductivity cell. Its cell constant, G star, the G with a little star, equals l over A: the distance between the electrodes, over their area. And it also equals kappa, the Greek letter for conductivity, times R, the resistance. Okay? Now, the resistance R we can measure easily. So, to get G star, what else must we know? Kappa, exactly! We need a solution whose conductivity is already known. Right? And that solution is K C l. The conductivities of K C l solutions are known accurately, at many concentrations and many temperatures. So fill the cell with K C l, measure the resistance, and kappa times R gives G star. Clear? There is more. K C l is a strong electrolyte, and its two ions, K plus and C l minus, have nearly equal mobilities: they move at nearly the same speed. And that is also what a salt bridge needs, which is why K C l turns up there too. Nice, right?

4:52Example 1: heat a solution vs a metal, conductivity change

Now, two examples from Part C, okay? Example one. Why does the conductivity of an electrolyte solution increase with temperature, while the conductivity of a metal decreases? Pause the video, and try it first. Done? Okay, let's check. Heat makes particles move faster, so both should conduct better, right? No, no, careful! First ask: who carries the charge? In a solution, the ions carry it. When the solution gets hotter, its viscosity falls: the liquid flows more easily, so the ions move faster. Okay? And for a weak electrolyte, there is a bonus. On heating, it ionises more, so there are more ions to carry the current. Right? Now the metal. Here, electrons carry the charge. Heating makes the metal lattice vibrate more strongly, and these lattice vibrations scatter the moving electrons more. More scattering means more resistance. Clear? So, heat a solution, and it conducts better. Heat a metal, and it conducts worse. Same heat, opposite results, because the charge carriers are different. Perfect!

6:05Example 2: why Li has the most negative E0 of alkali metals

And now our puzzle from the start, as Example two. Why does lithium have the most negative E naught of the alkali metals, despite its high ionisation energy? Pause the video and try it. Hint: think about the water. Done? Okay, let's check. From the table in Chapter three: for the lithium plus, lithium couple, E naught is minus three point zero four volt, the capital V for volt. For the potassium plus, potassium couple, it is minus two point nine three volt. Okay? These are standard reduction potentials, measured at two hundred and ninety eight kelvin, with one molar solutions, and one bar for gases. The more negative E naught is, the stronger the metal is as a reducing agent: the more readily it gives away its electron in water. Clear? Now the trap. Lithium has the highest ionisation energy in the group. It holds its electron the most tightly. So it should be the weakest reducing agent, right? No, careful! If ionisation energy alone decided it, caesium, C s, at the bottom of the group, would win, not lithium. So something else is going on, and that something is the water. With me? In water, the lithium ion gets surrounded by water molecules, and that releases energy: the hydration enthalpy. The lithium ion is very small, so the water molecules crowd in very close, and its hydration enthalpy is very large. Okay? Compared with the other alkali metals, lithium pays more to remove its electron, but it gets back much more when its small ion is hydrated. Pay a bigger cost, get back an even bigger reward. This extra hydration enthalpy more than offsets the higher ionisation energy, and that is why lithium's E naught is the most negative. Got it? So lithium is the strongest reducing agent in water. And our puzzle from the start? Solved! The highest ionisation energy, and still the strongest reducer, all because of the water. That's it!

8:08Quick revision: all six electrochemistry concept checks

So, quick revision. At the anode, always oxidation; its sign depends on the cell. E is intensive, but delta G, equal to minus n F E, is extensive. Agar holds the salt bridge as a gel: no mixing, but the ions still move. And K C l sets the cell constant because its conductivity is known and its ions have nearly equal mobilities. On heating, solutions conduct better and metals worse. And hydration makes lithium the strongest reducer in water. Now go and try the two practice problems on this page. All the best!