Voltaic Cells — a Refresher Remember from the first course that we can some7mes ‘catch’ the electrons passing from one substance to another by separa7ng the reactants by a wire and a salt bridge. One of the substances will have a greater affinity for electrons and electrons will flow from the substance with the weaker affinity to the substance with the greater affinity. The electrode where reduc7on occurs is the cathode and the electrode where oxida7on occurs is the anode. Ions pass over the salt bridge to balance the charge of the electrons flowing through the wire. A voltmeter measures the ‘pressure’ of electrons trying to get from one side to the other in units of volts. This pressure is referred to in scien7fic terms as a poten7al difference (but is more commonly known as voltage). We saw in the first course that a Zn-­‐Cu cell has a poten7al difference of 1.100 V and a Cu-­‐Ag cell has a poten7al difference of 0.462V. In the Zn-­‐Cu cell, Cu is the cathode, because Cu2+ ions receive electrons from Zn but in the Cu-­‐Ag cell, Cu is the anode, dona7ng electrons to Ag+ ions. We can, therefore, make a Zn-­‐Ag cell where Ag+ ions receive electrons from Zn. As we might expect, the poten7al difference for this cell is the sum of the poten7al differences of the Zn-­‐Cu and the Cu-­‐Ag cells: 1.562 V. Calcula7ng Cell Poten7als We have seen that we can calculate the poten7al difference of a theore7cal voltaic cell if we already know the poten7al differences of the two half cells with a third, common electrode : direc7on of electron ‘pressure’ Zn2+/Zn 1.100 V Cu2+/Cu 0.462 V Ag+/Ag 1.562 V Therefore, if we choose one electrode as a standard by which all other electrodes are measured, we only have to make actual measurements for substances with this ‘standard electrode’. All other combina7ons can be calculated from a table of standard electrode poten7als. For historical (and apparently arbitrary) reasons, chemists chose the hydrogen cell as the standard reference electrode, i.e. the cell where 2H+ + 2e– —> H2. By measuring the poten7al of half cells with reference to the hydrogen cell, we can calculate the theore7cal poten7al difference of any combina7on of half-­‐cells and determine which would be the cathode and which the anode: We can see from this diagram that a Li-­‐Au cell would have a 2H+/ H2 Li+/Li 3.045 V poten7al difference of 4.545 V that oxygen cannot take 2H+/ H2 1.229 V O2/2O2– and electrons from gold because the Au+/Au electron pressure is in the wrong 2H+/ H2 1.50 V direc7on. Notes on Standard Reduc7on Poten7als (To understand this slide, you will have to make reference to the table on the next slide) The following table from the textbook lists some common standard reduc1on poten1als. The first point to remember is that the hydrogen cell is defined as being zero. Next, to prevent confusion, all poten7als are given for reduc7ons by the standard hydrogen cell, i.e. where the hydrogen cell pushes electrons towards the other cells. For many cells (e.g. the Li-­‐H2 cell) this is actually the wrong direc7on i.e. 2Li+ + H2 —> 2Li +2H+ is the wrong way round. In these cases, the poten7al difference has a nega7ve value. The real direc7on of electron flow is defined as giving a posi7ve cell poten7al so the real cell, where lithium metal reduces protons (i.e. 2Li +2H+ —> 2Li+ + H2) has a posi7ve cell poten7al of 3.045 V. real direc1on of electron ‘pressure’ Li+/Li –3.045 V 2H+/ H2 Fe2+/Fe –0.44 V 2H+/ H2 2H+/ H2 1.229 V O2/2O2– Standard Reduc7on Poten7als This table appears in the textbook. (The Standard Hydrogen Cell) The standard electrode poten1al is the poten7al difference between any half cell and a HCl(aq)/
H2(g) (2H+ + 2e– —> H2) half cell. This half cell is known as the hydrogen cell and at 25 °C and concentra7ons of 1M H+, it is known as the standard hydrogen cell. The hydrogen half-­‐cell is made by dipping an electrode coated in fine pla7num par7cles, known as pla7num black, into a 1M HCl solu7on. Hydrogen gas is then bubbled over the electrode, which either reduces the H+ ions to H2 gas or oxidises the H2 gas to H+, depending on the cell poten7al. (In reality, hydrogen cells are too unreliable for modern measurements and the standard hydrogen cell is now only a theore7cal model of a perfect hydrogen cell) Using Standard Reduc7on Poten7als Let’s take, for example, a zinc/silver cell, i.e. Zn/Zn2+ v. Ag/Ag+. Method 1: First, write two half equa7ons but write the reduc7on with the most nega7ve poten7al as an oxida7on (i.e. reverse the half reac7on of the species being oxidised). This step helps you understand the chemistry of the reac7on. It also helps to make a balanced ionic equa7on to check your chemistry is correct but it is not necessary. Zn2+ + 2e–—> Zn –0.763 V => Zn —> Zn2+ + 2e– 0.763 V Ag+ + 2e– —> Ag0 +0.799 V Zn + 2Ag+ —> Zn2+ + 2Ag0 — This is the balanced ionic equa7on. It is only necessary for understanding the chemistry. Next, add the standard reduc7on poten7al of the oxidised species to the standard reduc7on poten7al of the reduced species. Note: we do not change the reduc7on poten7al according to stoichiometry, so, for example, we do not double the reduc7on poten7al of silver because two equivalents are used in the balanced equa7on. Standard cell poten7al for a Zn-­‐Ag cell = 0.799 V + 0.763 V = 1.56 V Calcula7ng cell poten7als — Shortcut We can see there is a shortcut to calcula7ng cell poten7als — just subtract the poten7al of the more nega7ve reduc7on (the species that is actually being oxidised) from the less nega7ve reduc7on (the real reduc7on), in other words red – ox. It is important, however, that you understand why the subtrac7on is this way round. Ag+ + 2e– —> Ag0 0.799 V O2 + 4e– —> 2O2– 1.229 V The silver poten7al is ‘more’ nega7ve so the cell poten7al is: 1.229 V – 0.799 V = 0.430 V This poten7al is posi7ve, so we know that oxygen can oxidise silver but it is not very high, which suggests that the oxida7on might be quite slow.* *(The cell poten7al is actually a thermodynamic property (i.e. it is directly related to ∆G) and not a kine7c property. Although a small cell poten7al (and therefore a small value of ∆G) suggests the reac7on might be slow, we cannot be certain un7l we perform a proper kine7c analysis.) The Meaning of Standard Poten7als A posi7ve standard poten7al means that a process is spontaneous, i.e. electrons will spontaneously flow the way the reac7on has been wrinen. This is the basis of voltaic cells. Conversely, a nega7ve standard poten7al means that a process is non-­‐spontaneous and we will have to add a poten7al difference of at least that amount to drive the process. This is the basis of electroly7c cells. Zn0 + Cu2+ —> Cu0 + Zn2+ standard cell p.d. = +1.100 V, spontaneous, power is generated. Zn2+ + Cu0 —> Cu2+ + Zn0 standard cell p.d. = –1.100 V, non-­‐spontaneous, power must be applied to the cell from outside to drive this reac7on. However, cell poten7als tell us far more than simply what happens in electrical cells. They tell us about how strongly substances hold on to their electrons. We can apply our knowledge of cell poten7als to understand chemical reac7ons. For instance, we can see from the above calcula7ons/measurements that zinc metal will spontaneously give electrons to copper ions. In other words, zinc is a reducing agent for copper, and copper is an oxidising agent for zinc. We can therefore use cell poten7als to calculate which elements are capable of reducing or oxidising certain other elements. Applica7ons of Standard Reduc7on Poten7als Standard reduc7on poten7als help us to understand which species are likely to react in redox reac7ons and which species will be the reducing agent and which will be the oxidising agent. For instance, we might wish to know if we can make a metal container that can contain sulphuric acid. The table of standard reduc7on poten7als tells us that the reac7on: 2Al + 6H+ —> 2Al3+ + 3H2 has a cell poten7al of 1.66 V and the reac7on: Fe + 2H+ —> Fe2+ + H2 has a cell poten7al of 0.44 V. These posi7ve cell poten7als tell us that aluminium and steel will react with sulphuric acid to produce H2 gas. However, the cell poten7al for the reac7on: Cu + 2H+ —> Cu2+ + H2 has a cell poten7al of –3.337 V, in other words, copper will not react with acids (in fact, the reac7on goes the other way and hydrogen could be used to precipitate Cu0 out of a solu7on of Cu2+ ions). This means we could line the inside of a steel container with copper to protect it from the acid. (In prac7ce, steel containers for acids are lined with plas7c).