Chapter 23 Redox and Electrode potentials

define reducing agent

transfers electrons to the species being reduced

is oxidised

define oxidising agent

accepts electrons from the species being reduced

is reduced

manganate (VII) titration method

1) standard solution of KMnO₄ is added to the burette

2) pipette a measured volume of the solution being analysed to the conical flask. add an excess of dilute H₂SO₄ to provide the H⁺ ions needed for the reduction of MnO₄^- (aq) ions

3) end point judged by first permanent pink colour = indicating excess of MnO₄^- ions present

4) read the meniscus from the from the top = KMnO₄ (aq) is a deep purple colour, very difficult to see bottom of meniscus

what can manganate (VII) titrations be used for the analysis of

Fe²⁺(aq)

(COOH)₂ (aq)

KMnO₄ under acidic conditions

oxidiSIng agent reduced to form Mn²⁺

what are iodine/thiosulfate redox titrations used for

determining content of different oxidising reagents (ClO^-, Cu²⁺)

equation for the reduction of iodine

I₂(aq) + 2e^- → 2I^- (aq)

equation for the oxidation of thiosulfate ions

2S₂O₃^2- → S₄O₆^2- (aq) + 2e^-

overall equation for iodine-thiosulfate titrations

2S₂O₃^2-(aq) +I₂(aq) → 2I (aq) + S₄O₆^2- (aq)

analysis of oxidising agents

1) add a standard solution of Na₂S₂O₃ to the burette

2) prepare a solution of the oxidising agent to be analysed. pipette this solution to a conical flask. add an excess a KI. oxidising agent reacts with I^- ions to produce iodine

3) titrate this solution with Na₂S₂O₃ (aq). iodine is reduced back to I^- ions, brown colour fades gradually making it difficult to decide on an end point = add starch near to the end point = colour change from pale yellow to blue-black = as more Na₂S₂O₃ is added, blue-black colour fades

define voltaic cell

a type of electrochemical cell which converts chemical energy into electrical energy

can be made by connecting together two different half-cells

define half-cell

contains the chemical species present in a redox equation

describe metal/metal ion half-cells

half-cell made by dipping a metal rod into a solution of its metal ion (aq)

represented using a vertical line for the phase boundary between the (aq) solution and the metal

at the phase boundary, an equilibrium will be set up = forward reaction shows reduction, reverse reaction shows oxidation

the direction of electron flow depends on the relative tendency of each electrode to release electrons

ion/ion half-cells

contains ions of the same element in different oxidation states

inert Pt(s) electrode used

define standard electrode potential

the emf of a half-cell compared with a standard hydrogen half-cell measured at 298K with solution concentration of 1moldm^-3 and a gas pressure of 100kPa

what is the standard electrode potential of a standard hydrogen electrode

0V

diagram of a standard hydrogen electrode

what does a salt bridge contain

a concentrated electrolyte that does not react with either solution

e.g. a strip of filter paper soaked in KNO₃ (aq)

the more negative the standard electrode potential…

the greater the tendency to lose electrons and undergo oxidation

the greater the reactivity of a metal in losing electrons

the more positive the standard electrode potential…

the greater the tendency to gain electrons and undergo reduction

the greater tendency of a non-metal in gaining electrons

steps of measuring standard cell potentials

1) prepare 2 standard half-cells

2) connect the metal electrodes of the half-cells to a voltmeter using wires

3) prepare a salt bridge by soaking a strip of filter paper in a saturated solution of KNO₃(aq)

4) connect the two solutions of the half-cells with a salt bridge

5) record the standard cell potential from the voltmeter

formula for calculating standard cell potential

standard electrode potential (positive electrode) - standard electrode potential (negative electrode)

the more negative redox system has the greatest tendency to…

be oxidised and lose electrons

reverse reaction

the more positive redox system has the greatest tendency to…

be reduced and gain electrons

forward reaction

how to predict feasibility of a redox reaction

redox system of the oxidising agent has a more positive standard electrode potential value than the redox system of the reducing agent

3 limitations of predictions using standard electrode potential values

1) reaction rate = electrode potentials give no indication on the rate of reaction

2) concentration = if the conc of the sol is not 1moldm^-3 the value of the electrode potential will be different from the standard value

3) other factors = standard conditions required, or reactions are not aqueous

what assumption are cells and batteries based on

the two electrodes have different electrode potentials

3 types of cells

primary, secondary, fuel

describe primary cells

non-rechargeable, designed to be used once

provide electrical energy until all the chemicals have reacted

used for low-current, long-storage devices (wall clocks, smoke detectors)

describe secondary cells

rechargeable

when recharging the reactions of the cells can be reversed

used in lead-acid batteries (cars), lithium polymer cells (phones)

describe fuel cells

uses energy from the reaction of a fuel with O₂ to create a voltage

operate continuously provided that the fuel and O₂ are supplied into the cell

hydrogen fuel cells most common

describe hydrogen fuel cells

produce no CO₂ during combustion, H₂O only combustion product

a hydrogen fuel cell can have either an alkali or acid electrolyte