Is a salt bridge required when using the standard hydrogen electrode to measure the potential of another species?

The differing reactivities of metals

When metals react, they give away electrons and form positive ions. This particular topic sets about comparing the ease with which a metal does this to form hydrated ions in solution - for example, Mg2+(aq) or Cu2+(aq).

We might want to compare the ease with which these two changes take place:

Everybody who has done chemistry for more than a few months knows that magnesium is more reactive than copper. The first reaction happens much more readily than the second one. What this topic does is to try to express this with some numbers.

Looking at this from an equilibrium point of view

Suppose you have a piece of magnesium in a beaker of water. There will be some tendency for the magnesium atoms to shed electrons and go into solution as magnesium ions. The electrons will be left behind on the magnesium.

In a very short time, there will be a build-up of electrons on the magnesium, and it will be surrounded in the solution by a layer of positive ions. These will tend to stay close because they are attracted to the negative charge on the piece of metal.

Some of them will be attracted enough that they will reclaim their electrons and stick back on to the piece of metal.

A dynamic equilibrium will be established when the rate at which ions are leaving the surface is exactly equal to the rate at which they are joining it again. At that point there will be a constant negative charge on the magnesium, and a constant number of magnesium ions present in the solution around it.

Simplifying the diagram to get rid of the "bites" out of the magnesium, you would be left with a situation like this:

Don't forget that this is just a snapshot of a dynamic equilibrium. Ions are continually leaving and rejoining the surface.

How would this be different if you used a piece of copper instead of a piece of magnesium?

Copper is less reactive and so forms its ions less readily. Any ions which do break away are more likely to reclaim their electrons and stick back on to the metal again. You will still reach an equilibrium position, but there will be less charge on the metal, and fewer ions in solution.

If we write the two reactions as equilibria, then what we are doing is comparing the two positions of equilibrium.

The position of the magnesium equilibrium . . .

. . . lies further to the left than that of the copper equilibrium.

Notice the way that the two equilibria are written. By convention, all these equilibria are written with the electrons on the left-hand side of the equation. If you stick with this convention without fail, you will find that it makes the rest of this topic much easier to visualise.

Everything else concerning electrode potentials is simply an attempt to attach some numbers to these differing positions of equilibrium.

In principle, that is quite easy to do. In the magnesium case, there is a lot of difference between the negativeness of the metal and the positiveness of the solution around it. In the copper case, the difference is much less.

This potential difference could be recorded as a voltage - the bigger the difference between the positiveness and the negativeness, the bigger the voltage. Unfortunately, that voltage is impossible to measure!

It would be easy to connect a voltmeter to the piece of metal, but how would you make a connection to the solution? By putting a probe into the solution near the metal? No - it wouldn't work!

Any probe you put in is going to have a similar sort of equilibrium happening around it. The best you could measure would be some sort of combination of the effects at the probe and the piece of metal you are testing.

Understanding the ideas behind a reference electrode

Suppose you had an optical device for measuring heights some distance away, and wanted to use it to find out how tall a particular person was. Unfortunately, you can't see their feet because they are standing in long grass.

Although you can't measure their absolute height, what you can do is to measure their height relative to the convenient post. Suppose that in this case, the person turned out to be 15 cm taller than the post.

You could repeat this for a range of people . . .

. . . and come up with a set of results like this:

Although you don't know any of their absolute heights, you can usefully rank them in order, and do some very simple sums to work out exactly how much taller one is than another. For example, C is 5 cm taller than A.

This turns out to be exactly what we need to do with the equilibria we started talking about. We don't actually need to know the absolute position of any of these equilibria. Going back to the magnesium and copper equilibria:

All we need to know is that the magnesium equilibrium lies further to the left than the copper one. We need to know that magnesium sheds electrons and forms ions more readily than copper does.

That means that we don't need to be able to measure the absolute voltage between the metal and the solution. It is enough to compare the voltage with a standardised system called a reference electrode.

The system used is called a standard hydrogen electrode.

Measuring standard electrode potentials (standard redox potentials)

Have you ever wondered what is inside a battery? You already know electric current results from the movement of electrons. So what are batteries made of so that they can provide flowing electrons? Read on to learn!

• In this article, you will learn about Standard Electrode Potential of a species
• You will learn about reduction potential and oxidation potential of a species.
• Oxidation state or oxidation number of a species.
• What is meant by oxidation or reduction of a species in a chemical reaction.
• Half-reactions or half-cells.
• Redox reactions.
• Galvanic cell.
• How chemical energy can be converted to electrical energy.

Consider this equation below.

This equation shows equilibrium between Cu2+ ions and ground state Cu. These are the two oxidation states of Copper - +2 and 0. The Copper ion has to gain 2 electrons to get to its ground state with 0 net charge.

The measure of the ability of a species to gain or lose electrons is called Standard Electrode Potential (E°).

The Oxidation State or Oxidation Number is the number of electrons a species has to gain or lose to form bonds with other species.

A species gets oxidized when it loses electrons. The oxidation state (charge on ion) increases.

Conversely, a species gets reduced when it gains electrons. Oxidation state decreases.

Electrode potential is measured in Volts. Standard electrode potential for Cu2+ is +0.34 volts.

Let's consider the Chlorine atom. You know that Chlorine has 7 electrons in its outermost electron shell, and only needs 1 more to have a completely filled stable electron shell. This means Chlorine has a high tendency to gain 1 electron and to exist with an oxidation state of -1. Since Standard Electrode Potential is a measure of the ability of a species to gain or lose an electron, Chlorine has a high Standard Electrode Potential.

E° for Chlorine is E° = +1.36V.

A species that gains electrons is said to undergo reduction and the electrode potential measured for these reactions is called the reduction potential of that species. Therefore, the reduction potential for Chlorine is +1.36V. The reduction potential of Cu2+ is +0.34V.

Reduction Potential is a measure of the ability of a species to gain electrons and get reduced in the process.

Conversely, Oxidation Potential is a measure of the ability of a species to lose electrons and get oxidized in the process.

Numerically, Oxidation Potential is the negative of Reduction Potential.

Let us see for Vanadium.

V2+ has a reduction potential of -1.2V. What does the negative sign imply? It implies that Vanadium is more likely to lose electrons than gain them i.e., more likely to get oxidized than get reduced. Therefore, Vanadium has the oxidation potential of +1.20V.

Species with highly positive reduction potentials are good oxidizing agents, since they can oxidize other species easily.

Conversely, species with highly negative reduction potentials are good reducing agents, since they can reduce the other species easily.

The 3 reactions which are written above (Cu2+, Cl2, V2+) are called half equations, or half cells. They only show the reduction side of a chemical reaction. But, in a chemical reaction, oxidation and reduction is happening in tandem. Therefore, these reactions are called redox reactions (reduction + oxidation) (Cool name, no?).

Generally, half equations are written as reduction equations, and standard electrode potential is written as standard reduction potential. This is done to avoid confusion and maintain uniformity. Since oxidation potential is just the negative of reduction potential, only one of them needs to be calculated.

Electrode Potential Table

Standard electrode potentials, E°, can be listed as an electrochemical series. This is also called the electrode potential Table. This table consists of half reactions of species undergoing reduction and stating the reduction potential of that half reaction. It is the IUPAC convention to always write the half reactions as reduction reactions i.e., The species are always shown to be gaining electrons. This is why the standard electrode potential of any species is the same as its standard reduction potential. This is done to establish a standard when comparing electrode potentials of any two species.

Standard electrode potential, E°, refers to conditions of 298 K, 100 kPa and 1.00 mol dm−3 solution of ions.

The reduction potential of H+ is the reference point for all other species, therefore, it is considered 0. Electrode potential for all half reactions is actually measured with reference to H+.

Redox Reactions

Redox reactions are chemical reactions where both an oxidation reaction and a reduction reaction are taking place in tandem.

Let us consider the half equations for Copper and Vanadium.

What If we combined these 2 half equations or half cells? One of the two (either Copper or Vanadium) will have to be the one that gives electrons and the other will have to be the one that accepts electrons. You have already seen that Copper has a tendency to accept electrons (get reduced), and Vanadium has a tendency to lose electrons (get oxidized). So naturally, the half-reaction for Copper will go in the forward direction, and the half-reaction for Vanadium will go in the reverse direction. In other words, Cu2+ will get reduced to Cu, and V will get oxidized to V2+.

we can't have both half reactions giving a receiving electrons because the electrons in the equation have to be balanced too. The electrons have to come from either of these two, they can't come from thin air!

Giving the following reaction -

Electrons and the net charge on both sides are balanced.

When combining 2 half equations, Balancing the charge on both sides of the equation is of utmost importance!

This whole equation will also have a net electrode potential value. The net electrode potential value for the reaction is the difference between the electrode potential of the reduction reaction and electrode potential of the oxidation reaction. -

For the Copper-Vanadium cell, Copper undergoes the reduction reaction, and Vanadium undergoes the oxidation reaction.

The total E° measured for a cell is called Electromotive Force of the cell, or simply EMF.

Electromotive Force (EMF) is defined as an electrical action produced by a non-electrical source.

Let us take another example - the Zinc-Copper battery. Consider the half-reactions for Copper and Zinc -

What do you think will happen if we combine these 2 equations? Which of the two (Zinc and Copper) will undergo oxidation and which will undergo reduction? This is decided by the sign of the electrode potentials. Since E° for Copper half-reaction is positive, it will go in the forward direction (Copper will get reduced).

Thus, giving us the redox reaction -

We can also calculate the E° for the combined redox reaction.

Electrochemical cells can convert either chemical energy to electrical energy, or vice versa. An electrochemical cell that converts chemical energy to electrical energy is called a Galvanic cell. We will only discuss about Galvanic cells.

You have seen how 2 half reactions can combine to give a redox reaction. You have also seen that these redox reactions have a net voltage. Thus, redox reactions can be used to form an electrochemical cell called galvanic cell, which can produce electric current. They combine the electrode potentials of the two half reactions form an electrical circuit with a net EMF.

To make a galvanic cell, you will need -

• 2 containers to contain electrolyte of the 2 half reactions.
• 2 electrodes.
• A salt bridge. Salt bridge is needed to facilitate the movement of ions from one electrolyte to the other.
• A wire to connect the two electrodes.
• A voltmeter to measure the voltage across he cell.

An Electrode is a conductor which makes contact with the non-metallic parts in an electric circuit.

For the Zn-Cu galvanic cell, you need -

• a Zinc electrode and a copper electrode.
• Zinc Sulphate (ZnSO4) and Copper Sulphate (CuSO4) solutions for the electrolytes.
• Potassium Chloride (KCl) solution for the salt bridge.

You already know that in the Zn-Cu cell, Copper undergoes reduction and Zinc undergoes oxidation i.e., Copper will gain electrons and Zinc will lose electrons. Zn atoms from the solid electrode will release electrons and get dissolved as Zn2+ ions in the ZnSO4 electrolyte solution. The electrons released by Zn atoms will travel through the contact wire over to the Cu electrode. Cu2+ ions from the CuSO4 electrolyte will gain these electrons and convert into solid Cu with the electrode. This whole process is facilitated due to the electric potential difference between the two half cells, which can be measured in the voltmeter attached to the wire connecting the 2 electrodes.

1. In a Galvanic cell, the electrode at which oxidation takes place is called anode. The electrode at which reduction takes place is called cathode.
2. Electrons flow from Anode to Cathode. Current flows from Cathode to Anode.

Copper sulphate solution provides Cu2+ ions for the Copper half cell. As more and more Zn atoms release electrons and get dissolved in the electrolyte, you will see the Zn electrode get thinner with time, just as the Cu electrode gets thicker.

The salt bridge helps balance the net charge in both electrolytes of the 2 half cells. As more cations are released into the ZnSO4 solution, K+ ions from the salt bridge permeate into the solution to balance the net charge. Similarly, Cl- ions permeate into the CuSO4 solution.

Recall the definition for electrode potential you read at the beginning of this article. Now that you have understood the concept, let us define electrode potential, and revise the definition of standard electrode potential.In a half cell, due to the separation of charges between the electrode and the electrolyte, the electrode may be positively or negatively charged with respect to the electrolyte. This charge difference is what causes the potential difference between the electrode and the electrolyte. This potential difference is the electrode potential. The electrode potential is called standard electrode potential when the concentration of all species involved in a half cell is unity.

Electrode Potential is the potential difference between the electrode and the electrolyte in a half cell.

There is no way of measuring the electric potential of an electrolyte, therefore, electrode potential for a species is measured according to this definition -

Electrode Potential of a species is the electromotive force (emf) / E°cell of a galvanic cell built from a half cell of that species and a reference half cell (H+).

Electrode potential is called Standard Electrode Potential when the concentration of all species involved in a half cell is unity.

Disproportionation reactions are chemical reactions in which an element undergoes oxidation as well as reduction, like so -

The oxidation and reduction halves of this equation can be written as -

This redox reactions can be seen as an electrochemical cell. Therefore, the two half equations of this electrochemical cell would be -

Note that half reactions are written as reduction reactions according to IUPAC convention.

If you calculate the electrode potential of the whole cell -

The electrode potential of the disproportionation reaction of Copper is 0.36V.

In the section "Electrode Potential Table", we briefly mentioned that electrode potential for H+ is considered 0. It is the reference electrode for all other species and standard electrode potential for all species is measured against H+ , called the Standard Hydrogen Electrode (SHE). In this section, we will see how standard electrode potential is measured against the Standard Hydrogen Electrode.

First of all, how can an electrode be made out of Hydrogen? It's a tube in which Hydrogen gas is passed. A piece of Platinum serves as the electrical contact and also as a catalyst in the half reaction of Hydrogen. The tube is dipped in an electrolyte containing H+ ions. The other half of the galvanic cell consists of the electrode of which the electrode potential has to be measured.

The setup to calculate the standard electrode potential of Zinc is described in the figure -

In this cell, Zinc forms the oxidation half cell and the SHE forms the reduction half cell. That means Zinc undergoes oxidation and Hydrogen undergoes reduction. You already know that the electrode where oxidation takes place is called the anode, and the electrode where reduction takes place is called the cathode. Since Zinc forms the anode in this cell, and electrode potential of cathode 0 (Since it is a SHE), the EMF measured for this cell is negative. Mathematically, it can be calculated as -

It is important to note that Platinum only serves as an electrical contact and a catalyst. It does not contribute to the EMF of the cell.

At the beginning of this article, we defined the term "standard electrode potential". Now that we have understood everything around the term, let us redefine it.

Standard electrode potential / standard reduction potential is the ability of a species to reduce a standard hydrogen electrode at conditions of 298 K, 100 kPa and 1.00 mol dm−3 ion concentration.

A cell can be made using two species (two half reactions). The net cell potential depends on the electrode potential of the two species.

Standard cell potential, Eo, is the difference between the potentials of the reduction half and the oxidation half of a cell. It is measured at standard conditions of 298 K, 100 kPa and 1.00 mol dm−3 ion concentration.

Electrode Potential - Key takeaways

• Standard Electrode Potential is a measure of the ability of a species to gain or lose electrons.
• Species which gains electrons is said to be reduced. Species which loses electrons is said to be oxidized.
• Reduction potential is the ability of a species to get reduced. It is the ability of a species to oxidize other species.
• Oxidation potential is the negative of reduction potential. It is the ability of a species to get oxidized. It is its ability to reduce other species.
• Species with highly positive reduction potential are good oxidizing agents. Species with highly negative reduction potentials are good reducing agents.
• Redox reactions are reactions where a reduction reaction and an oxidation reaction is taking place in tandem.
• Each half of a redox reaction (oxidation/reduction) is called a half reaction or half cell.
• Galvanic cells are electrochemical cells which can convert chemical energy to electrical energy.
• Galvanic cells are based on redox reactions. They combine the electrode potentials of the two electrodes of the two half reactions to form an electrical circuit
• The potential difference between the electrode and the electrolyte is called Electrode Potential.
• In a Galvanic cell, the electrode at which oxidation takes place is called anode. The electrode at which reduction takes place is called cathode.
• Electrons flow from anode to cathode. Current flows from cathode to anode.
• Electrode potential for all species is calculated against the Standard Hydrogen Electrode (SHE). It is the reference electrode and its potential is considered to be 0.