Chapter 3 - Primary galvanic cells and fuel cells as sources of energy

Question 1

redox reactions

Answer 1

reactions that involve the transfer of one or more electrons between chemical species

Question 2

reduction

Answer 2

a gain of electrons; a decrease in the oxidation number - electrons with reactants in half-equation

Question 3

oxidation

Answer 3

a loss of electrons; an increase in the oxidation number - electrons with products in half-equation

Question 4

oxidising agents

Answer 4

electron acceptors

Question 5

reducing agents

Answer 5

electron donors

Question 6

half-equations

Answer 6

an equation that gives one half of a redox reaction, showing the movement of electrons in either an oxidation or a reduction reaction

Question 7

oxidation numbers

Answer 7

numbers used to find an oxidising agent and a reducing agent by a change in perceived valency

Question 8

spontaneous reactions

Answer 8

reactions that proceed on their own, without the need for any external supply of energy

Question 9

salt bridge

Answer 9

a component that provides a supply of mobile ions that carry the charge through the solution of a galvanic cell during a reaction - often simply filter paper soaked in an ionic salt solution such as KNO3 or NaCl
Purposes:
- It connects the circuit, allowing for the flow of charged particles throughout the circuit
- It physically separates the two half cells, preventing the reactants from coming into direct contact with each other
- It provides cations and anions that can migrate into the solutions to balance the charges in each half cell. Cations migrate to the cathode and anions migrate to the anode

Question 10

electrolytes

Answer 10

liquids that can conduct electricity

Question 11

internal circuit

Answer 11

a circuit within a solution; anions flow to the anode and cations flow to the cathode

Question 12

half-cell

Answer 12

one half of a galvanic cell containing an electrode immersed in an electrolyte that may be the oxidising agent or the reducing agent, depending on the oxidising strength of the other cell to which it is connected
3 types:
- metal ion-metal half-cell
- solution half-cell
- gas-non-metal ion half-cell

Question 13

electrodes

Answer 13

a solid used to conduct electricity in a galvanic half-cell
Design:
- Electrodes used are porous to increase their surface area allowing for an increased surface for oxidation and reduction to occur
- Electrodes are often coated with a catalyst (Ni or Pt) to allow the reaction to take place a lower temps while maintaining an increased rate of reaction
- Highly conductive to allow redox reactions to occur at the surface
- Often two layers to withstand high temps

Question 14

external circuit

Answer 14

a circuit composed of all the connected components within an electrolytic or a galvanic cell to achieve desired conditions

Question 15

anode

Answer 15

the electrode at which oxidation occurs (AnOx); in a galvanic cell it is the negative electrode, since it is the source of negative electrons for the circuit; if the reducing agent is a metal, it is used as the electrode material

Question 16

cathode

Answer 16

the electrode at which reduction occurs (RedCat); in a galvanic cell it is the positive electrode, because the negative electrons are drawn towards it and then consumed by the oxidising agent, which is present in the electrolyte

Question 17

electrochemical cell

Answer 17

a cell that generates electrical energy from chemical reactions

Question 18

Daniell cell

Answer 18

one of the first electrochemical cells to produce a reliable source of electricity; it uses the redox reactions between zinc metal and copper ions to produce electricity

Question 19

voltmeter

Answer 19

a device used for measuring the potential difference between two points in a circuit

Question 20

cell potential differences

Answer 20

the difference between the reduction potentials of two half-cells

Question 21

electrical potential

Answer 21

the ability of a galvanic cell to produce an electric current

Question 22

standard electrode potential

Answer 22

the voltage or potential difference due to the difference in charge on the electrode and electrolyte compared to the hydrogen half-cell

Question 23

reduction potential

Answer 23

a measure of the tendency of an oxidising agent to accept electrons and so undergo reduction

Question 24

standard cell potential difference

Answer 24

the measured cell potential difference, under standard conditions, when the concentration of each species in solution is 1 M, the pressure of a gas (where applicable) is 100 kPa and the temperature is 25 °C (298 K)

Question 25

standard hydrogen half-cells

Answer 25

a standard reference electrode; it is assigned 0.00 volts - consists of hydrogen gas bubbling around an inert platinum electrode in a solution of hydrogen ions - The reaction that occurs at the electrode surface is: 2H+(aq) + 2e- <=> H2(g) E^0 = 0.00V

Question 26

electrochemical series

Answer 26

a series of chemical half-equations arranged in order of their standard electrode potentials

Question 27

primary cells

Answer 27

a type of galvanic cell that store the reactants and products of the redox reaction taking place - an electrolytic cell in which the cell reaction is not reversible (single use - cannot be recharged) - During the operation of primary cells electrons build up at the anode and are drawn to the cathode by travelling through a connecting wire (external circuit) to balance the charge of the cell - As the redox reaction continues to take place reactants are used up, this will lead to the cell eventually stopping producing electrons and the battery will ‘die’ The most common types of primary cells are: the dry cell, the alkaline zinc/manganese dioxide cell and lithium cells

Question 28

lithium cells

Answer 28

cells that use lithium anodes and can produce a high voltage - more expensive than other batteries but shelf life of ten years The most common lithium battery is the lithium-manganese cell, consisting of a lithium anode, a manganese dioxide cathode and a non-aqueous electrolyte such as propylene carbonate C4H6O3

Question 29

dry cell

Answer 29

an electrochemical cell in which the electrolyte is a paste, rather than a liquid; also called a Leclanché cell - It consists of a zinc container filled with electrolyte paste (a mixture of MnO2, ZnCl2, NH4Cl and H2O) A carbon rod is pushed through the paste and forms the cathode The zinc container is the anode No salt bridge is required as the paste prevents cell contents mixing Slightly warming or using intermittently will increase the life of the cell and prevent products forming around the electrodes Cell equations: - Oxidation (Anode): Zn(s) → Zn2+(aq) + 2e- - Reduction (Cathode): 2MnO2(s) + 2NH4+(aq) + 2e- → Mn2O3(s) + 2NH3(aq) + H2O(l) - Overall: 2MnO2(s) + 2NH4+(aq) + Zn(s) → Mn2O3(s) + 2NH3(aq) + Zn2+(aq) + H2O(l) The oxidising agents and reducing agents used in such cells should: - be far enough apart in the electrochemical series to produce a useful voltage from the cell - not react with water in the electrolyte too quickly, or they will discharge early (therefore, highly reactive metals such as sodium, potassium and calcium are not found in such batteries) - be inexpensive.

Question 30

fuel cell

Answer 30

a type of galvanic cell / an electrochemical cell that produces electrical energy directly from a fuel - By combining fuels such as hydrogen and oxygen in the presence of an electrolyte, the products of a fuel cell are electricity, heat and water. - Fuel passes over the anode (and oxygen over the cathode) where it is split into ions and electrons. The electrons go through an external circuit while the ions move through the electrolyte towards the oppositely charged electrode. At this electrode, ions combine to create by-products. Depending on the input fuel and electrolyte, different chemical reactions occur.

Question 31

proton exchange membrane fuel cell (PEMFC)

Answer 31

-a fuel cell being developed for transport applications, as well as for both stationary and portable fuel cell applications -Hydrogen from the fuel gas stream is consumed at the anode, producing electrons that flow to the cathode via the electric load (external circuit) and hydrogen ions that enter the electrolyte - At the cathode, oxygen combines with electrons from the anode and hydrogen ions from the electrolyte to produce water - The PEMFC operates around 80oC so the water does not dissolve in the electrolyte, instead it is collected at the cathode and carried out of the fuel cell by excess oxidising agent (O2) flow Cell equations: - Reduction (Cathode): O2(g) + 4H+(aq) + 4e- → 2H2O(l) - Oxidation (Anode): H2(g) → 2H+(aq) + 2e- - Overall: 2H2(g) + O2(g) → 2H2O(g)

Question 32

fuel reformer

Answer 32

a device or system that converts a fuel source — typically hydrocarbons or alcohols — into a hydrogen-rich gas mixture

Question 33

direct methanol fuel cell (DMFC)

Answer 33

- a new technology that is powered by liquid methanol -The liquid methanol is more energy dense than hydrogen and easier to transport Cell equations: Reduction (Cathode): O2(g) + 4H+(aq) + 4e- → 2H2O(l) Oxidation (Anode): 2CH3OH(aq) + 2H2O(l) → 2CO2(g) + 12H+(aq) + 12e- Overall: 2CH3OH(aq) + 3O2(g) → 2CO2(g) + 4H2O(l)

Question 34

alkaline fuel cell (AFC)

Answer 34

- a fuel cell that converts oxygen (from the air) and hydrogen (from a supply) into electrical energy and heat - uses KOH as the electrolyte - The alkaline environment of the AFC allows for the use of non-precious metal catalysts such as iron, cobalt, silver and graphite, which lowers the cost of these fuel cells Cell equations: Reduction (Cathode): O2(g) + 2H2O(l) + 4e- → 4OH-(aq) Oxidation (Anode): H2(g) + 2OH-(aq) → 2H2O(l) + 2e- Overall: 2H2(g) + O2(g) → 2H2O(l)

Question 35

atom economy

Answer 35

a measurement of the efficiency of a reaction that considers the amount of waste produced, by calculating the percentage of the molar mass of the desired product compared to the molar mass of all reactants

Question 36

thermochemical splitting

Answer 36

refers to when very high temperatures are used to decompose molecules by breaking chemical bonds

Question 37

photodecomposition

Answer 37

the use of light (photons) to break down molecules

Question 38

electolysis

Answer 38

the process in which a non-spontaneous chemical reaction occurs by passing an electric current through a substance in solution or molten state

Question 39

cellulosic fermentation

Answer 39

the use of enzymes to obtain glucose from cellulose to make alcohol

Question 40

carboxylic acid

Answer 40

the homologous series containing the —COOH functional group

Question 41

electrolysis

Answer 41

the process in which a non-spontaneous chemical reaction occurs by passing an electric current through a substance in solution or molten state

Question 42

electrolytic cells

Answer 42

an electric cell in which a non-spontaneous redox reaction is made to occur by the application of an external potential difference across the electrodes; also known as an electrolysis cell

Question 43

Faraday's First Law

Answer 43

states that the amount of current passed through an electrode is directly proportional to the amount of material released from it

Question 44

Faraday constant

Answer 44

represents the amount of electric charge carried by 1 mole of electrons

Question 45

Faraday's Second Law

Answer 45

states that when the same quantity of electricity is passed through several electrolytes, the mass of the substances deposited are proportional to the stoichiometric coefficients in the balanced half-equations of the respective electrochemical reactions

Question 46

Does the oxidant or reductant come first in a conjugate pair?

Answer 46

The oxidant

Question 47

How to balance half-equations in acidic environments

Answer 47

K - Key elements O - oxygen by adding water H - Hydrogen by adding H+ E - balance electrons S - assign states

Question 48

How to balance half equations in alkaline environments

Answer 48

K - Key elements O - Oxygen by adding H2O H - Hydrogen by adding H+ OH - Add OH- on BOTH sides to balance H+ H2O - Combine the H+ and OH- and cancel out water on both sides E - balance electrons S - assign states

Question 49

galvanic cell

Answer 49

A type of electrochemical cell (converts chemical energy to electrical energy) by utilising spontaneous reactions between oxidising agents and reducing agents Galvanic cells consist of: - Two half cells, containing two electrodes (anode and cathode) and two electrolytes - Conducting wire (external circuit) - Salt bridge, containing another electrolyte (internal circuit) - An electrolyte is a solution containing ions that can conduct electricity - An electrode is a conductor through which electrons enter or leave a galvanic cell - The anode is the electrode where oxidation occurs - The cathode is the electrode where reduction occurs - Electrons flow from the anode to the cathode

Question 50

metal ion-metal half cell

Answer 50

consist of a metal rod in a solution of its ions, usually from the sulfate salt. The sulfate ion is unreactive.

Question 51

solution half cell

Answer 51

use an inert (unreactive) electrode in the reactive solution containing the oxidant or reductant. Inert electrodes are often graphite or platinum)

Question 52

gas-non-metal ion half-cell

Answer 52

cell involves an inert electrode that has a gas bubbled over it. Its conjugate redox non-metal ion is in solution.

Question 53

Calculating standard cell potential difference

Answer 53

E^0 cell = E^0 oxidising agent - E^0 reducing agent

Question 54

How do you read the electrochemical series

Answer 54

The strongest oxidants are on the top left (reduction equations, read forwards) and they have weak conjugate reductants The strongest reductants are on the bottom right (oxidation equation, read backwards) and they have weak conjugate oxidants

Question 55

Alkaline zinc/manganese dioxide cell

Answer 55

similar to the dry cell in terms of their components, however have added features that all it have a longer shelf live and generate a higher current output (work faster and longer) Even though these batteries may last five times longer than a dry cell, they are difficult to make and more expensive. Both dry cell and alkaline cell batteries are bulky making them unsuitable for smaller devices like watches and calculators The alkaline cell has a powdered zinc anode in an electrolyte paste of KOH The cathode is a compressed mixture of manganese dioxide and graphite A separator, porous fibre soaked in electrolyte, prevents the mixing of anode and cathode components The cell is contained within a steel shell Cell equations: - Oxidation (Anode): Zn(s) + 2OH-(aq) → Zn(OH)2(s) + 2e- - Reduction (Cathode): 2MnO2(s) + 2H2O(l) + 2e- → 2MnO(OH)(s) + 2OH-(aq) - Overall: 2MnO2(s) + 2H2O(l) + Zn(s) → 2MnO(OH)(s) + Zn(OH)2(s)

Question 56

How do fuel cells differ from primary cells?

Answer 56

- Fuel and oxygen are supplied externally. - Unreacted fuel and products are removed from the cell. - Fuel cells don’t go flat — electricity is generated for as long as reactants are supplied.

Question 57

solid oxide fuel cell

Answer 57

- The SOFC uses a ceramic (solid oxide) electrolyte that conducts O2- ions at high temperatures - Oxygen gas is reduced at the cathode to produce oxide ions that travel through the ceramic electrolyte to the anode, where they react with hydrogen ions from the oxidised fuel and make water Cell equations: - Reduction (Cathode): O2(g) + 4e- → 2O2-(in ceramic) - Oxidation (Anode): H2(g) + O2-(in ceramic) → H2O(g) + 2e- Overall: 2H2(g) + O2(g) → 2H2O(g)

Question 58

advantages of fuel cells

Answer 58

- Clean energy – no CO2 produced (H2O is a greenhouse gas, so can not say NO greenhouse gases are produced) - Electricity will be continuously produced, as long as a fuel is supplied, don’t go flat - Hydrogen gas can be produced from biomass - Doesn’t require main grid electricity - Quieter

Question 59

disadvantages of fuel cells

Answer 59

- Storage of hydrogen – flammable, ‘light’ gas = tend to escape - Produces DC but AC used in homes, therefore requires an inverter - Expensive

Question 60

How to calculate electric charge:

Answer 60

Q = It where: - Q = electrical charge in coulombs (C) - I = current in amperes (A) - t = time in seconds (s)

Question 61

n(e-) = Q/F

Answer 61

where: - n = moles - Q = electrical charge in coulombs (C) - F = Faraday constant, 96,500 C/mol