Lecture 16: Electrochemistry-- The Big Picture Electrochemistry follows the adventures of the electron e− Recall that we first discussed the electron when it came up as a fundamental particle back when discussing quantum mechanics. Also recall that when learning about configurations for atoms and molecules that we obsessed over it. But as we went on to discuss thermodynamics and properties of the bulk, we put it away. It is back now, demanding its own chapter and perhaps its own consideration in thermodynamic terms. After all, we spent six weeks on the fundamental particle called the proton and it has nothing on the electron from a reactivity perspective. How do we know this? Think about it from a charge density perspective as we compare some particles of different charge and size Note that while the mass of an electron is quite small—2000 times smaller than a proton and 50,000 times smaller than a sodium ion: It has the same amount of charge which means that its charge density is vastly larger than any other particle. Seeing as how we have on many occasions we have used charge density to rank reactivity, having stumbled across a particle with 2000 times the charge density of the next guy in line, we should be willing to examine it more carefully. Type of particle e- H+ H Na+ mass per particle 8.3 x 10-27 g 1.6 x 10-24 g 1.6 x10-24 g 3.7 x 10-23 g mass per mole 1/2000 g/mole 1 g/mole 1 g/mole 23 g/mole charge per particle 1.614x10-19 C 1.614 x 10-19 C 0 1.614 x 10-19 C charge per mole 9.6 x 104 C/mole = 1 F 9.6 x 104 C/mole = 1 F 0 9.6 x 104 C/mole = 1 F electronTime out for charge: Everyone understands mass—you step on scales every day to see how much you weigh and are used to buying goods “by the pound or gram.” And that is why we can talk about the mass of a mole of an atom and you just look at the periodic table and rattle off “uranium weighs 238 grams/mole.” But charge is not so familiar. You are used to the notion of charge in the form of static electricity or getting a shock, but have no real quantitative feel for it. Until now. Very simply: • The unit for charge is the Coulomb in the same way the unit of mass is grams. and • In the same way you can assign a mass to a mole of a compound—like water is 18 grams/mole, you can assign a mole of charge to something. And the best part is that a mole of charge always has exactly the same value: 1 mole of charge = 1 Faraday = 9.6 x104 C/mole And this value is the same whether you have a mole of e- or a mole of H+ or a mole of Na+The chemistry of the electron—oxidation/reduction reactions: We have just finished a unit on equilibria in which we looked at reactions in which we followed the movement of a proton and reactions in which we followed the movement of cations and anions: Acid base reactions solubility reactions H+ proton Na+ or Cl- cations or anions NH3 + HNO3 NH4+ + NO3-- NaCl + AgNO3 NaNO3 + AgCl Note that in both cases, the oxidation number does not change: for example, in the acid base reaction, oxidation numbers of H = +1, N = -3 and O = -2 are assigned for every atom on both sides of the reaction. In contrast, in electrochemistry we focus on reactions in which the formal oxidation number assigned to an atoms in a reaction changes. In every oxidation reduction reaction, at least one atom is oxidized (increases in oxidation number) while at least one atom is reduced (decreases in oxidation number.) For example, 0 +1 -2 +1 -2 +1 0 note the Na atom oxidation number increases while the H atom oxidation number decreases Na + H 2O Na+ + OH- + H2 0 0 +1 –2 note the H atom oxidation number increases while the O atom oxidation number decreases 2H 2 + O2 = 2H 2O What we do in electrochemistry is follow the flow of electrons along this chemical reaction path.Time out—if you can’t assign oxidation numbers as easily as you breathe, you will fail the electrochemistry material. So here is a refresher brought to you by your friendly neighborhood chemistry professor: Assigning oxidation numbers What do oxidation numbers tell us? Very simply, oxidation numbers tell us where the electron density is in a compound—kind of like all that work we did with quantum mechanics at the beginning of 301, but with none of the huge equations and complicated rules. So in the ion dichromate +6 -2 Cr2O7-2 the electrons are more densely packed around the O though than Cr and you can tell this because the more negative the oxidation number, the greater the electron density. Note surprisingly this result is consistent with everything you learned about oxygen loving electrons and metals not wanting them, So what are the rules for oxidation number: rather than try to memorize all the possible combinations, learn the following hierarchy of rules: • Rule 0: Free elements have an oxidation number of 0. Example: Oxygen in O2 has 0 oxidation number • Rule 1: Assigning oxidation numbers to the groups, from the left. the common formal charge on an element increases, +1, +2, +3.... From the right, skipping the noble gases, the common formal charge on an element increases, -1, -2,-3. So alkali metals are +1. Halogens are -1, etc. Complications involving hydrides, H, and oxides, O, with other elements in compounds: • Rule 2: Alkali metals are always +1 except for rule 0. Example Na in NaO, is +1. • Rule 3: Hydrogen is always +1 except for rules 0 and 2. Example, H is -1 in NaH but +1 in H2O. • Rule 4:. O is always -2 except for rules 0, 2, and 3. Example: O is -0.5 in NaO2, but -2 in H2O. • Rule 5: In every redox problem you do, assign the oxidation numbers in order using rules 0- 4. There will usually only be one element left to assign by difference. Example, what is the oxidation number of P in K2HPO4?Example of setting up problems to determine oxidation numbers for unknowns. What is the oxidation number for P in K2HPO4? charge on the molecule = 0 = (ox. # of K)(# of K atoms) + (ox. # of H)(# of H atoms) + (ox. # of P)(# of P atoms) +
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