Electrochemistry and battery technology Contents

Size: px

Start display at page:

Download "Electrochemistry and battery technology Contents"

Jeffery Howard
5 years ago
Views:

1 Electrochemistry and battery technology Contents Introduction Redox overview voltaic cells, electrolytic cells, fuel cells, Primary and secondary batteries. Other batteries; Construction, working and applications of Lithium, Li-MnO2 and Li-ion batteries Cell potential 1

2 Introduction Electrochemical systems The essence of a chemical reaction is the exchange of atoms and electrons between molecular species. In most reactions, the exchange of the electrons is not manifest because it occurs in an intimate mixture of reactants and products. However, in a device known as an electrochemical cell, participants in a reaction are physically separated in a manner that makes it possible for the electron transfer process to be observed, measured, and utilized for doing useful work. The study of chemical reactions with this emphasis on the electron-transfer process is called electrochemistry. An obvious example for an electrochemical product that has been well known for generations is the lead acid battery of the cars. An other products are fuel cells and a hand held sensor for monitoring the glucose diabetes patients 2

3 Overview of Redox Reactions Oxidation is the loss of electrons and reduction is the gain of electrons. These processes occur simultaneously. Oxidation results in an increase in O.N. while reduction results in a decrease in O.N. The oxidizing agent takes electrons from the substance being oxidized. The oxidizing agent is therefore reduced. The reducing agent takes electrons from the substance being oxidized. The reducing agent is therefore oxidized. 3

4 Redox reactions terminologies Oxidation reduction reactions involve a transfer of electrons Zn(s) + 2H + (aq) Zn 2+ (aq) + H 2 (g) 4

5 Memory Aid Oxidation Involves Loss Reduction Involves Gain OIL RIG 5

6 Electrochemical Cells A voltaic cell uses a spontaneous redox reaction to generate electrical energy. - The system does work on the surroundings. A electrolytic cell uses electrical energy to drive a nonspontaneous reaction - The surroundings do work on the system. Both types of cells are constructed using two electrodes placed in an electrolyte solution. The anode is the electrode at which oxidation occurs. The cathode is the electrode at which reduction occurs. 6

7 General characteristics of (A) voltaic and (B) electrolytic cells. 7

8 Construction of a Voltaic Cell Each half-reaction takes place in its own half-cell, so that the reactions are physically separate. Each half-cell consists of an electrode in an electrolyte solution. The half-cells are connected by the external circuit. A salt bridge completes the electrical circuit. The salt bridge completes the electrical circuit and allows ions to flow through both half-cells. As Zn is oxidized at the anode, Zn2+ ions are formed and enter the solution. Cu2+ ions leave solution to be reduced at the cathode. The salt bridge maintains electrical neutrality by allowing excess Zn2+ ions to enter from the anode, and excess negative ions to enter from the cathode. A salt bridge contains nonreacting cations and anions, often K+ and NO3-, dissolved in a gel. 8

9 Connected this way the reaction starts. Electricity travels in a complete circuit H + e - MnO - Fe

10 A voltaic cell based on the zinc-copper reaction. Oxidation halfreaction Zn(s) Zn2+(aq) + 2eoccurs at the anode, which is therefore the source of e-. Overall (cell) reaction Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) Reduction halfreaction Cu2+(aq) + 2e- Cu(s) occurs at the cathode, where the e- are used up. 10

11 Operation of the Voltaic Cell The anode produces e - by the oxidation of Zn(s). The anode is the negative electrode in a voltaic cell. Electrons flow through the external wire from the anode to the cathode, where they are used to reduce Cu 2+ ions. The cathode is the positive electrode in a voltaic cell. Over time, the Zn(s) anode decreases in mass and the [Zn 2+ ] in the electrolyte solution increases. Over time, the [Cu 2+ ] in this half-cell decreases and the mass of the Cu(s) cathode increases. 11

12 Memory Aid ANode, OXidation; REDuction, CAThode AN OX and a RED CAT: 12

The Pain of a Dental Voltaic Cell Have you ever felt a jolt of pain when biting down with a filled tooth on a scrap of foil left on a piece of food? Here s the reason.

13 The Pain of a Dental Voltaic Cell Have you ever felt a jolt of pain when biting down with a filled tooth on a scrap of foil left on a piece of food? Here s the reason. The aluminum foil acts as an active anode (E of Al = 1.66 V), saliva as the electrolyte, and the filling (usually a silver/tin/mercury alloy) as an inactive cathode. O2 is reduced to water, and the short circuit between the foil in contact with the filling creates a current that is sensed by the nerve of the tooth. 13

14 Electrochemical Processes in Batteries A battery consists of self-contained voltaic cells arranged in series, so their individual voltages are added. A primary battery cannot be recharged. The battery is dead when the cell reaction has reached equilibrium. A secondary battery is rechargeable. Once it has run down, electrical energy is supplied to reverse the cell reaction and form more reactant. 14

15 Alkaline battery Anode (oxidation): Zn(s) + 2OH - (aq) ZnO(s) + H 2 O(l) + 2e - Cathode (reduction): MnO 2 (s) + 2H 2 O(l) + 2e - Mn(OH) 2 (s) + 2OH - (aq) Overall (cell) reaction: Zn(s) + MnO 2 (s) + H 2 O(l) ZnO(s) + Mn(OH) 2 (s) E cell = 1.5 V 15

16 Silver button battery Anode (oxidation): Zn(s) + 2OH - (aq) ZnO(s) + H 2 O(l) + 2e - Cathode (reduction): Ag 2 O(s) + H 2 O(l) + 2e - 2Ag(s) + 2OH - (aq) Overall (cell) reaction: Zn(s) + Ag 2 O(s) ZnO(s) + 2Ag(s) E cell = 1.6 V The mercury battery uses HgO as the oxidizing agent instead of Ag 2 O and has cell potential of 1.3 V. 16

Lithium battery The primary lithium battery is widely used in watches, implanted medical devices, and remote-control devices. Anode (oxidation): 3.

17 Lithium battery The primary lithium battery is widely used in watches, implanted medical devices, and remote-control devices. Anode (oxidation): 3.5Li(s) 3.5Li e - Cathode (reduction): AgV 2 O Li e - Li 3.5 V 2 O 5.5 Overall (cell) reaction: AgV 2 O Li(s) Li 3.5 V 2 O

18 Lead-acid battery The cell generates electrical energy when it discharges as a voltaic cell. Anode (oxidation): Pb(s) + HSO 4- (aq) PbSO 4 (s) + H + (aq) + 2e - Cathode (reduction): PbO 2 (s) + 3H + (aq) + HSO 4- (aq) + 2e - PbSO 4 (s) + 2H 2 O(l) Overall (cell) reaction (discharge): PbO 2 (s) + Pb(s) + H 2 SO 4 (aq) 2PbSO 4 (s) + 2H 2 O(l) E cell = 2.1 V Overall (cell) reaction (recharge): 2PbSO 4 (s) + 2H 2 O(l) PbO 2 (s) + Pb(s) + H 2 SO 4 (aq) The lead-acid car battery is a secondary battery and is rechargeable. 18

19 Lithium-ion battery Anode (oxidation): Li x C 6 (s) xli + + xe - + C 6 (s) Cathode (reduction): Li 1-x Mn 2 O 4 (s) + xli + + xe - LiMn 2 O 4 (s) Overall (cell) reaction: Li x C 6 (s) + Li 1-x Mn 2 O 4 (s) LiMn 2 O 4 (s) E cell = 3.7 V The secondary (rechargeable) lithium-ion battery is used to power laptop computers, cell phones, and camcorders. 19

20 Electrolytic Cells An electrolytic cell uses electrical energy from an external source to drive a nonspontaneous redox reaction. Cu(s) Cu 2+ (aq) + 2e - Sn 2+ (aq) + 2e - Sn(s) Cu(s) + Sn 2+ (aq) Cu 2+ (aq) + Sn(s) [anode; oxidation] [cathode; reduction] E cell = V and ΔG = 93 kj As with a voltaic cell, oxidation occurs at the anode and reduction takes place at the cathode. An external source supplies the cathode with electrons, which is negative, and removes then from the anode, which is positive. Electrons flow from cathode to anode. 20

21 Electrolysis and batteries 21

Zn(s) Zn 2+ (aq) + 2e - Cu 2+ (aq) + 2e - Cu(s) Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq)

22 The tin-copper reaction as the basis of a voltaic and an electrolytic cell. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Zn(s) Zn 2+ (aq) + 2e - Cu 2+ (aq) + 2e - Cu(s) Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq) voltaic cell Cu(s) Cu 2+ (aq) + 2e - Zn 2+ (aq) + 2e - Zn(s) Zn 2+ (aq) + Cu(s) Zn(s) + Cu 2+ (aq) electrolytic cell 22

23 Electroplating 23

24 Fuel Cells Fuel cells are electrochemical devices that convert the chemical energy of a reaction directly into electrical energy. The basic physical structure or building block of a fuel cell consists of an electrolyte layer in contact with a porous anode and cathode on either side. This is very different from conventional combustion based power plant which convert chemical energy to thermal energy, then thermal energy to kinetic energy, and only then kinetic energy to electrical energy. Using hydrogen as a fuel (which can be extracted from hydrocarbon fuels or renewable sources) a fuel cell electro-chemically oxidizes the hydrogen using oxygen from the air generating electricity and some heat. The fuel cell makes more efficient use of the fuel and produces fewer pollutants e.g. reduced nitrogen oxides and carbon dioxide emissions,. So they are environmentally friendly products 24

25 Hydrogen fuel cell It consists of three components - a cathode, an anode, and an electrolyte sandwiched between the two. Oxygen from the air flows through the cathode A fuel gas containing hydrogen, such as methane, flows past the anode. Negatively charged oxygen ions migrate through the electrolyte membrane react with the hydrogen to form water, The reacts with the methane fuel to form hydrogen (H2) & carbon dioxide (CO2). This electrochemical reaction generates electrons, which flow from the anode to an external load and back to the cathode, Anode (oxidation): 2H 2 (g) 4H + (aq) + 4e - Cathode (reduction): O 2 (g) + 4H + (aq) + 4e - 2H 2 O(g) Overall (cell) reaction: 2H 2 (g) + O 2 (g) 2H 2 O(g) E cell = 1.2 V 25

26 26

27 Fuel cells can be big! Or very small Fuel cell technology, power space vehicles and electric cars GM 27

28 June 2016, the Toyota Mirai is available for retail sale in Japan, California, the UK, Denmark, Germany, Belgium, and Norway. A modified Ford Focus) 28

29 Cell Potential A voltaic cell converts the DG of a spontaneous redox reaction into the kinetic energy of electrons. The cell potential (E cell ) of a voltaic cell depends on the difference in electrical potential between the two electrodes. Cell potential is also called the voltage of the cell or the electromotive forces (emf). E cell > 0 for a spontaneous process. 29

30 Voltages of Some Voltaic Cells Voltaic Cell Voltage (V) Common alkaline flashlight battery 1.5 Lead-acid car battery (6 cells 12 V) 2.1 Calculator battery (mercury) 1.3 Lithium-ion laptop battery 3.7 Electric eel (~5000 cells in 6-ft eel = 750 V) Nerve of giant squid (across cell membrane)

31 Measuring the standard cell potential of a zinc-copper cell. The standard cell potential is designated E cell and is measured at a specified temperature with no current flowing and all components in their standard states. By convention, all standard electrode potentials refer to the half-reaction written as a reduction. The standard cell potential depends on the difference between the abilities of the two electrodes to act as reducing agents. E cell = E cathode (reduction) - E anode (oxidation) 31

32 E cell E E Where: E + = the reduction potential for the ½ cell reaction at the positive electrode E + = electrode where reduction occurs (cathode) E - = the reduction potential for the ½ cell reaction at the negative electrode E - = electrode where oxidation occurs (anode) 32

33 Sample Problem Calculating an Unknown E half-cell from E cell PROBLEM: A voltaic cell houses the reaction between aqueous bromine and zinc metal: Br 2 (aq) + Zn(s) Zn 2+ (aq) + 2Br - (aq) E cell = 1.83 V. Calculate E bromine, given that E zinc = V PLAN: E cell is positive, so the reaction is spontaneous as written. By dividing the reaction into half-reactions, we see that Br 2 is reduced and Zn is oxidized; thus, the zinc half-cell contains the anode. We can use the equation for E cell to calculate E bromine. SOLUTION: Br 2 (aq) + 2e - 2Br - (aq) [reduction; cathode] Zn(s) Zn 2+ (aq) + 2e - [oxidation; anode] E zinc = V E cell = E cathode E anode 1.83 = E bromine (-0.76) = E bromine E bromine = 1.07 V 33

Electrochemical cells Two types of electrochemical cells: 1. A battery has all of its chemicals stored inside, and it converts those chemicals into electricity too.

34 Electrochemical cells Two types of electrochemical cells: 1. A battery has all of its chemicals stored inside, and it converts those chemicals into electricity too. This means that a battery eventually "goes dead" and you either throw it away or recharge it. 2. Fuel cells: To convert the chemicals hydrogen and oxygen into water, and in the process it produces electricity. Chemicals constantly flow into the cell so it never goes dead 34

35 Differences between Voltaic and Electrolytic Cells. Galvanic cell, Voltaic This is an electrochemical power source. Cathode Positive (cations) Anode Negative (anions) Electrolytic cell. This is an electrochemical substance producer. Produce chemical Cathode Negative (cations) Anode Positive (anions) The cell does work by releasing free energy from a spontaneous reaction to produce electricity. The cell does work by absorbing free energy from a source of electricity to drive a non-spontaneous reaction Examples: - Battery - Fuel cell Examples: -Electro synthesis -Electro plating 35

Chapter 21: Electrochemistry: Chemical Change and Electrical Work

Chapter 21: Electrochemistry: Chemical Change and Electrical Work CHEM 1B: GENERAL CHEMISTRY Instructor: Dr. Orlando E. Raola Santa Rosa Junior College 21-1 Chapter 21 Electrochemistry: Chemical Change