Chapter 23 Flashcards by Ishir Sharma

Reduction

Gain of electrons or decrease in oxidation number

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Oxidation

Loss of electrons or increase in oxidation number

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What are likely products of redox reactions?

H2O, H+ or OH-

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Two common redox titrations?

Potassium manganate (VII) under acidic conditions
Sodium thiosulfate Na2S2O3 for determination of iodine (I2)

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Manganate VII titrations

These ions are reduced and so the other chemical must be the reducing agent that is oxidised - a standard solution is added to the burette and using a pipette a measured volume of the solution being analysed is added to the conical flask with an excess of dilute Sulfuric Acid to provide H+ ions required for reduction of MnO4- (this is self indicating)
During the titration the manganate solution reacts and is decolourised as it is added - END POINT = first permanent pink colour - repeat until obtain concordant titres

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What to note about the meniscus?

KMnO4 is a deep purple colour so it is very difficult to see the bottom of the meniscus through the intense colour ; thus burette readings are read from the top rather than the bottom - titre is still the same provided it is read the same way from burette

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Different reducing agents for potassium permanganate?

Iron (ii) ions
Ethanedioic acid
Reduce from MnO4- to Mn2+

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Iodine thiosulfate titrations?

Thiosulfate ions are oxidised
Iodine is reduced

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Oxidation of thiosulfate

2S2O32- -> S4O62- + 2e-

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Reduction

I2 + 2e- -> 2I-

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Overall redox reaction

2S2O32- + I2 -> 2I- + S4O62-

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How to carry out analysis of oxidising agent using iodine thiosulfate titration?

Na2S2O3 to burette
Solution of oxidising agent is added to the conical flask with excess of potassium iodide - oxidising agent reacts with iodide to produce iodine which turns the solution a yellow brown colour
Titrations this solution with Sodium thiosulfate in burette and iodine is reduced back to I- ions and brown colour fades quite gradually with no set end point - by using a starch indicator we can fix this as at the end point all iodine will have just reacted and blue black colour disappears (all iodine reduced to I-)

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What two oxidising agents can be analysed using iodine thiosulfate titrations?

Chlorate (I) ions, ClO-
Copper (ii) ions, Cu2+

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Voltaic cell?

Converts chemical energy into electrical energy - electrical energy results from movement of electrons so you need chemical reactions that transfer electrons from one species to another (these are redox reactions)

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Half cell

Contains chemical species present in a redox half-equation - voltaic cell can be made by connecting together two different half cells which allows electrons to flow

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What if chemicals in two half cells are allowed to mix?

Electrons would flow in an uncontrolled way and heat energy would be released rather than electrical energy

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Simples half cell

Metal rod dipped into a solution of its aqueous metal ion with the vertical line representing a phase boundary between aqueous solution and the metal - (Zn2+IZn)

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What happens at the phase boundary?

Equilibrium is set up and by convention FORWARD REACTION SHOWS REDUCTION and reverse shows oxidation - when two half cells are connected, direction of electron flow depends upon relative tendency of each electrode to release electrons

Ion/ion half cells?

Contain ions of the same element in different oxidation states - in this type of half cell there is no metal to transport electrons either into or out of the half-cell so an inert metal electrode made out of platinum is used

In an operating cell?

Electrode more reactive metal loses electrons more readily and is oxidised - AT THE ANODE
Electrode with less reactive metal that gains electrons is reduced - at the CATHODE

Standard electrode potential

The tendency to be reduced and gain electrons is measured as a standard electrode potential - standard is a half cell containing hydrogen gas and a solution of H+ ions ; inert platinum electrode is used to allow electrons into and out of half cell

Standard conditions used are

Solutions have a concentration of 1 mol dm^-3
Temperature is 298 K
Pressure is 100kPa

Standard electrode potential?

E.m.f of a half cell connected to a standard hydrogen half cell under standard conditions and by definition the standard electrode potential of a standard hydrogen electrode is exactly 0V

How to measure standard electrode potential?

The half cell is connected by a wire to allow a controlled flow of electrons - these are usually connected with a salt bridge which allows ions to flow - a salt bridge typically contains a concentrated solution of an electrolyte that does not react with either solution such as KNO3

The more negative the standard electrode potential?

Greater tendency to lose electrons and undergo oxidation - ANODE

The more positive the standard electrode potentials?

The greater the tendency to gain electrons and undergo reduction - CATHODE

More negative the Electrode potential

Greater the reactivity of a metal in losing electrons

More positive the standard electrode potential

Greater the reactivity of a non-metal in gaining electrons

Standard electrode potentials?

Electrode potential of positive electrode - electrode potential of negative electrode

Oxidising and reducing agent?

Oxidising agent is reduced thus on the left Reducing agent is oxidised thus on the right

Feasibility of a redox system rule?

Reaction should take place between oxidising agent and reducing agent provided that the redox system of the oxidising agent has a more positive standard electrode potential than the redox system of the reducing agent

The more positive the electrode standard potential?

Greater tendency to be reduced - so it’s oxidising agent should react with reducing agents on the right in redox systems with a lower electrode potential

Redox system with the more positive electrode potential?

Will gain electrons (reduction)

Redox system with the more negative electrode potential?

Will react from right to left losing electrons

Limitations of predictions using standard electrode potentials?

Reaction rate - reactions that have a very large activation energy resulting in a very low rate ; they may indicate thermodynamic feasibility of a reaction but give no indication on the rate of reaction Concentration - if concentration of solution is not 1 mol dm^3 then value will be different - if concentration greater than 1, equilibrium shifts to right removing electrons from system and electrode potential = less negative If concentration of solution less than 1 mol dm^3 equilibrium shifts to the left increasing the electrons in the system and making the electrode potential more negative

What other factors may affect the value of the electrode potential?

Actual conditions used for the reaction may be different from the standard conditions - standard electrode potential apply to aqueous equilibrium and many reactions take place that are not aqueous

Types of cells

Primary Secondary Fuel

Primary cells

Non-rechargeable and designed to be used once only - the reactions cannot be reversed and eventually the chemicals will be used up, the voltage will fall and the battery goes flat Still present in low current long storage devices like clocks

Secondary cells?

They are rechargeable - cell reaction producing electrical energy can be reversed during recharge if ; chemicals in the cells are regenerated and the cell can be used again Lead acid batters in cars and lithium ion polymer cells used in our modern appliances laptop/tablets etc

Fuel cells

Use energy from the reaction of a fuel with oxygen to create a voltage - fuel and oxygen flow into the fuel cell and products flow out with the electrolyte remaining in the cell ; fuel cells can operate continuously provided that the fuel and oxygen are supplied

Do fuel cells have to be recharged?

Hydrogen fuel cells?

Produce no carbon dioxide during combustion with water being the only combustion product - they can either have an acid of alkaline electrolyte (same electrode potential regardless)

Hydrogen fuel cell alkali

Anode - oxidation - H2 + 2OH- -> 2H2O + 2e- Electrons travel to cathode Reduction - 1/2O2 + H2O + 2e- -> 2OH-

Acid hydrogen fuel cell?

Anode : H2 -> 2H+ + 2e- - oxidation Travels to cathode Reduction - Cathode - 1/2O2 + 2H+ + 2e- -> H2O