7.5 – Electrochemical Cells

1. 7.5 – Electrochemical Cells Unit 7 – Redox Reactions & Electrochemistry

2. Introduction • During redox reactions, electrons pass from one substance to another. The flow of electrons - electric current - can be harnessed to do work. Electrochemistry is the branch of chemistry that deals with the conversion between chemical and electrical energy. • There are two major branches of electrochemistry: • 1. Electrochemical Cells: • the energy released by a spontaneous chemical reaction is converted into electrical energy. • Example: batteries. • 2. Electrolytic Cells: • electrical energy is used to cause a non-spontaneous reaction to occur. • Examples: recharging batteries, electroplating. • We will spend the remainder of the course studying Electrochemical Cells

3. Electrochemical Cells • The basic unit of all batteries is the electrochemical cell (also called a voltaic cell or galvanic cell). Electrochemical cells convert the energy of a spontaneous redox reaction into electricity. This will be accomplished as the electrons that are released from the oxidation half-reaction are passed to the reduction reaction which will absorb the electrons. • We will create an electrochemical cell based on the following redox reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) • This reactions involves two half-reactions: • Oxidation: Zn → Zn +2 + 2 e- • Reduction: Cu2+ + 2e- → Cu • In order for electrical work to be done by this reaction, we need to have the electrons travel through an external circuit. If we simply placed a piece of zinc metal in a solution containing copper(II) ions, a reaction would occur but electricity would not be created.

4. How to set up an Electrochemical Cell • 1) Begin by getting 2 beakers into which we will place metal strips in electrolytic solutions • solutions that conduct electricity due to the presence of ions. • In one place a strip of zinc metal in a Zn(NO3)2 solution. • In the other place a strip of copper metal in a Cu(NO3)2 solution.     • Each beaker represents one of the two half cells. But because there is no way for electrons to move from one beaker to the other, our redox reaction cannot yet occur.

5. How to set up an Electrochemical Cell • 2) We need to connect our two half-cells which we need to do in two ways.     • First we will connect the two metal strips, our electrodes, with some wire. We'll also place a voltmeter here so we can detect the electric current once we are up and running. This will be our external circuit. • Second we add a salt bridge. A salt bridge is a U-shaped tube that contains an electrolytic solution (we'll use KNO3). This electrolytic solution will allow ions to flow between the two beakers. This is our internal circuit.

6. How to set up an Electrochemical Cell • 3) The zinc half-cell undergoes oxidation. Here, the solid zinc electrode disintegrates, forming zinc ions and releasing electrons. By definition, the half-reaction that undergoes oxidation in an electrochemical cell is called the anode. • The anode is the source of electrons, making it the negative post of the electrochemical cell. • Anode = oxidation • “An ox” • The copper half-cell undergoes reduction. Here, copper ions from the electrolytic solution become deposited on the copper electrode, forming more solid copper. Electrons are required for this to occur. By definition, the half-reaction that undergoes reduction in electrochemical cells is called the cathode. • The cathode is the positive post of the electrochemical cell as it consumes electrons. • Cathode = reduction • “Red cat”

7. How to set up an Electrochemical Cell • 4) It is important to understand the roles of the external circuit and the salt bridge.     • External circuit: • this is where the electrical work is done as electrons travel from one half-cell to the other. The electrons are produced at the zinc anode, where oxidation occurs. The electrons then travel through the wire of the external circuit to the copper cathode. The electrons are then available for the copper ions (from the Cu(NO3)2 solution) and solid copper is produced.     • Internal circuit: • At the anode, Zn2+ ions are being produced and go into solution. This causes a build-up of positive ions in this solution. If this electrical imbalance is not corrected the reaction cannot continue. The excess positive charge attracts the negative NO3- ions (anions) from the salt bridge, thereby keeping the solution electrically neutral.

8. How to set up an Electrochemical Cell • At the cathode the opposite occurs. As positive Cu2+ ions are removed from solution, to form solid Cu, the solution becomes overly negative. This attracts the positive K+cations from the salt bridge, keeping this side of the cell neutral.     • Once we have the entire electrochemical cell assembled - the two half-cells (the electrodes in their electrolytic solutions), the internal circuit (the salt bridge and half-cells), and the external circuit (the wire connected the two electrodes) - the cell is complete and the redox reaction will occur. • It is important to note and remember that unless the electrons can pass from one electrode to the other the reaction will not proceed.

9. Standard Electrode Potentials • One question you may have had when we set up our zinc and copper cell was why was zinc oxidized and copper reduced, and not the other way around? In fact, for any redox reaction what determines which element is oxidized and which is reduced? • Metals, because they only have a few valence electrons, like to lose electrons. In other words they tend to be easily oxidized. But metals differ in how easily they lose electrons. A list of metals arranged in order of how easily the metal is oxidized is known as an activity series. You will remember this from our solutions unit! • The fact that different substances are oxidized more readily than others is the driving force behind electrochemical cells, and it is this force that forces electrons through the external circuit from the anode (site of oxidation) to the cathode (site of reduction). • This force is known as the potential difference or electromotive force (emfor E). • Potential difference is measured in volts (V), and thus is also referred to as the voltage of the cell. • Voltage is a measure of the tendency of electrons to flow. The higher the voltage, the greater the tendency for electrons to flow from the anode to the cathode.

10. Standard Electrode Potentials • Tables of Standard Reduction Potentials for Half-Reactions allow us to determine the voltage of electrochemical cells. These tables compare the ability of different half-reactions to compete for electrons (become reduced). Since half-reactions cannot occur on their own, all values in the table are determined by comparing a half-reaction with a hydrogen half-cell: 2H+(aq) + 2e- → H2 (g)    E° = 0.00 V • the degree symbol following the E (E°) indicates standard conditions: • temperature = 25°C; • pressure = 100 kPa; • concentration of aqueous solutions = 1 mol·L-1 • (Find this half-reaction, and other the other half-reactions described below, in the Table.) This hydrogen half-cell has been assigned a voltage of 0.00 V. If a half-reaction is better at competing for electrons than this half-cell, that half-reaction will undergo reduction and the hydrogen will be oxidized. That other half-reaction will then be assigned a positive voltage.

11. Electrode Potentials – Example 1 • If copper and hydrogen half-cells are joined together we find that the copper half-cell will gain electrons from the hydrogen half-cell. Thus the copper half-cell is given a positive voltage and given a relative value of +0.34 V: Cu2+(aq) + 2e- → Cu(s)    E° = 0.34 V • Since both half-reactions cannot undergo reduction, we must reverse the equation of the reaction that will undergo oxidation. This will give us an electrochemical cell voltage of 0.34 V: • U

12. Electrode Potentials – Example 2 • We see in the Table of Standard Reduction Potentials that zinc has a negative E° indicating that it is not as good at competing for electrons as hydrogen.  Zn2+(aq) + 2e- → Zn(s)    E° = -0.76 V • Therefore if zinc and hydrogen are paired together in an electrochemical cell, the hydrogen would be reduced (gain the electrons) and zinc would be oxidized (losing electrons). To determine the net redox reaction as well as the voltage of the electrochemical cell we reverse the zinc equation, and also reverse it's sign before adding the equations and E° together: • mem