redox and electrode potentials Flashcards
1
Q
what are the two definitions for reduction
A
- gain of electrons
- decrease in oxidation number
2
Q
what are the two definitions of oxidation
A
- loss of electrons
- increase in oxidation number
3
Q
what is an oxidising agent
A
- takes electrons from species being oxidised
- contains species that is reduced
4
Q
what is a reducing agent
A
- add electrons to the species being reduced
- contains species that is oxidised
5
Q
what occurs in a redox reaction
A
- there will always be oxidising agent and a reducing agent
6
Q
what are two common redox titrations
A
- potassium manganate (VII) (KMn)4(aq)) under acidic conditions
- sodium thiosulfate (Na2S2O3(aq)) for determination of iodine (I2(aq))
7
Q
describe the proceduree of manganate (VII) titrations
A
- a standard solution of potassium manganate (VII) KMNO4 is added to the burette
- using a pipette add a measured volume of the solution being analysed to the conical flask. an excess of dilute sulfuric acid is also added to provide tthe H+(aq) ions required for the reduction of MnO4-(aq) ions you do not need to add an indicator as the reaction is self indicating
- during the titration the manganate (VII) solution reacts and is decolourised as it is being added. the end point of the titration is judged by the first permanent pink colour indicating when there is an excess of MnO4- ions present. in titrations this end point is one of the easiest to judge
- repeat the titration until you obtain concordant titres (two titre that agree within 0.10 cm3)
- read the top of meniscus
8
Q
what are two examples of manganate (VII) titrations
A
- used for the analysis of many different reducing agents
- iron (II) ions Fe2+ (aq)
- ethanedioic acid (COOH)2(aq)
9
Q
describe the process how you analyse the percentage purity of an iron(II) compound
A
- prepare a 250cm3 solution of impure FeSO4 .7H2O in a volumetric flask
- using a pipette measure 25cm3 of this solution into a concical flask then add 10cm3 of 1 moldm-3 H2SO4(aq) (an excess)
- using a burette titrate this solution using a standard 0.02moldm-3 solution of potassium manganate (VII) KMnO4 (aq)
- finally analyse your result to detemermine the percentage purity
10
Q
how can redox titrations be used to anaysis of different substances
A
- manganate (VII) titrations can be used to analyse reducing agents that reduce MNO4- to Mn2+
- KMnO4 can be replaced with other oxidising agents the comonest used being acidified dichromate (VI) H+/Cr2O72-
*
11
Q
what happens in iodine thiosulfate titrations
A
- thiosulfate ions S2O32-9aq) are oxidised and iodine I2 is reduced
12
Q
what is the equation for the iodine/thiosulfate titrations
A
2S2O32-(aq) +I2(aq) ->2I-(aq) +S4O62-(aq)
13
Q
what is determined by titration with a standard solution of sodium thiosulfate
A
concentration of aqueous iodine
14
Q
what can iodine/thiosulfate titrations used to be determine
A
- ClO- content in household bleach
- Cu2+ content in copper (II) compounds
- Cu content in copper alloys
15
Q
describe the process of oxidising agent
A
- add a standard solution of Na2S2O3 to the burette
- prepare a solution of the oxidising agent to be analysed.using a pipette add this solution to a conical flask. then add an excess of potasium iodide. the oxidising agent reacts with iodide ions to produce iodine which turns the solution a yellow-brown colour
- titrate this solution with the Na2S2O3 (aq). during the titration the iodine is reduced to back to I-ions and the brown colour fades quite gradually making it difficult to decide on an end point
- this problem is solved by using a starch indicator . when the end point is being approached the iodine colour has faded enough to become a pale straw colour
- the fading of the yellow brown iodine colour as aqueous sodium thiosulfate is added is shown
16
Q
how do you use starch as an end point with iodine thiosulfate titrations
A
- when the end point is being approached and then the iodine colour has faded enough to become a pale straw colour a small amount of starch indicator is added
- a deep blue black colour forms to assist with the identification of the end point
- as more sodium thiosulfate is added to the blue black colour fades
- at the end point all the iodine will have just reacted and the blue black colour dissappears
17
Q
what is the ingredients in household bleach
A
- active ingredient is household bleach is chlorate (I) ions CLO- (aq) commonly known as hypochlorite
- many bleach as NaClO content which supplies hypochlorie (CLO-) ions
18
Q
describe the process for analysis of household bleach
A
- using a pipette add 10cm3 of the bleach into 250cm3 volumetric flask and add water to prepare 250cm3 of solution
- using a pipette measure 25 cm3 of this solution into a conical flask then add 10cm3 of 1moldm-3 potassium iodide (KI) followed by sufficient 1 moldm-3 HCL(aq) to acidify the solution HCl provide to H+ ions for the reaction
- using a burette titrate this solution using a standard 0.05moldm-3 aolution of sodium thiosulfate (Na2S2O3)(aq)
- repeat the titration to obtain concordant results
- finally analyse your results to determine the concentration of chlorate (I) ions in the bleach
19
Q
describe the analysis of copper
A
- iodine / thiosulfate titrations can be used to determine the copper content of coppe (II) salts or alloys
- for copper (II) salts Cu2+(aq) ions are produced simply by dissolving the compound inwater
- insoluble copper (II) compounds cna be reacted with acids to form Cu2+ (aq) ions
20
Q
describe the analyis of copper in copper alloys
A
- cooper alloys such as brass or bronze the alloy is reacted and dissolved in concentrated nitric acid followed by neutralisation to form Cu2+(aq) ions
- Cu(s)->Cu2+(aq)
- in this analysis:
- Cu2+9aq) ions react with I-(aq) to form a solution of iodine I2(aq) and a white precipirate of copper (I) iodide CUI(s)
- mixture appears as a brown colour
- 2Cu2+(aq) +4I-(aq) ->2CUI(s)+ I2(aq)
- the iodine in the brown mixture is then titrated with a standard solution of sodium thiosulfate
- 2S2O32-(aq) +I2(aq)->2I-(aq) +S4O62-(aq)
21
Q
what is an electrochemical cell
A
- voltaic cell
- converts chemical energy into electrical energy
22
Q
where does the conversion of chemical energy into electrical energy takes place
A
- modern cells
- bateries that power devices such as mobile phpnes
23
Q
what is the redox reactions happening in electrochemical cells
A
- electrical energy results from the movement of electrons
- do not need chemical reactions that transfer electrons from one species to another
24
Q
what is a half cell
A
- contains the chemical species present in a redox half equation
- a voltaic cell can be made by connecting together two different haldd cells which then allows electrons to flow
- in the cell the chemicals in the two half cells must be kept apart if allowed to mix electrons would flow in an uncontrolled way and heat energy would be released rather than electrical energy
25
describe the metal/metal ion half cells
* metal rod dippeed into a solution of its aqueous metal ion
* where metal is in contact with its ions an equilibrium will be set up
* in an isolated half cell there is no net transfer of electrons either into or out of the metal
* when two half cells are conncetred the direction of electron flow depends upon the relative tendency of each electrode to release electrobs
26
what is ion / ion half cells
* an ion/ion half cell contains ions of the same elment in different oxidation states
* no metal to transport electrons either into or out of the half cell so an inert metal electrode made out of platinum is used
27
give two examples of a metal/ metal ion half cells
* Zn2+(aq)/Zn(s)
* Cu2+(aq)/Cu(s)
28
give an example of ion/ion hald cells
* aquous iron (II) and iron(III) ion
* redox equlibrium
* Fe3+ +e- reversible sign Fe2+ (aq)
29
describe how you know what electrode has a greater tendency to gain or lose electrons
* in a cell with two metal/metal ion half cells connected the more reactive metal releases electrons more readily and is oxidised
* in an operating cell:
* the electrode with more reactive metal loses electrons and is oxidised - the negative electrode
* the electrode with the less reactive metal gains electrons and is reduced - positive electrode
30
what is the standard electrode potential
* tendency to be reduced and gain electrons
31
what is the standard chosed for standard electrode potential
* is a half cell containg hydrogen gas H2(g) and a solution containing H+(aq) ions
32
what is an inert plantinum electrode used for in a standard half cell
allow elecctrons into and out of the half cell
33
what are standard conditions for the standard electrode potential
* solutions have a concentration of extacly 1 moldm-3
* the temeprature is 298k
* the prezsure is 100kpa
34
what is the definition of a standard electrode potential
* the emf od a half cell connected to a standard hydrogen half cell under standard contitions at 298k , solution concentrations of 1moldm-3 and a pressure of 100kpa
35
what is the standard electrode potential of a standard hydrogen electrode
0v
36
describe wht the sign of a standard electrode potential shows
* sign of the half cell connnected to a standard hydrogen electrode and shows the relative tendency to gain electrons compared with the hydrogen half cell
37
how do you measure a standard electrode potential
* half cell connected to a standard hydrogen electrode
* the two electrodes are connected by a wire to allow a controlled flow or electrons
* the two solutions are connected with a salt bridge which allows ions to floww the salt bridge typically contains a concentrated solution of an electrolyte which does not react with either solution
* an example of a salt bridge is a strip of fliter paper soaked in aqueous potassium nitrate KNo3(aq)
38
describe wht the more negative electrode potential value shows
* greater tendency to lose electrons and undergo oxidtion
* less tendency to gain electrons and undergo reduction
39
describe what a more positive electrode value shows
* greater tendency to gain electrons and undergo reduction
* less tendency to lose electrons and undergo oxidation
40
describe the electrode potential values of metals and non metals
* metals - negative electrode values so lose electrons
* non metals - positive electrode values o gain electrons
41
describe the general trend of electrode potentials for metals and non metals
* the more negative the electrode potential the greater reactivity of a metal in losing electrons
* the more positive the electrode value the greater the reactivity of a non metal in gaining electrons
42
describe the process to set up cells to measure standard cell potentials
* prepare two standard half cells:
* for a metal/metal ion half cell the metal ion must have a concentration of 1moldm-3
* for an ion/ion halff cell both metal ions present in the solution must have the same concentration there must be an inert electrode usually platinum
* for a hald cell containing gase (hydrogen half cell) the gas must be at 100kpa pressure in cpntact with a solution wiht an ionic concentration of 1moldm-3 there must be an inert electrode usually platinum
* for all half cells temeprautre must be 298k
* connect the metal electrode of the half cells to a voltmeter using wires
* prepare a salt bridge by soaking a strip of filter paper in a saturated aqueous solution of potassium nitrate KNO3
* connect the two solutions of the half cells with a salt bridge
* record the standard cell potential from the voltmeter
43
describe what happens in a zinc-copper cell
* the copper half - cell has a more positive electrode potential and a greater tendecy to undergo reduction and to gain electrons
* the zinc hald cell has the more negative electrode potential and a hreater tendency to undergo oxidation and lose electrons
* electrons flow along wire from more negative zinc hald cell to the less negative copper half cell
* the zinc electrode negative and copper electrode is positive
44
write the equation for a zinc-copper half cell explain how you reached this conclusion
* the copper equilibrium has the more positive electrode potential undergoing reduction gaining electrons, the reduction half equation is the same way round as the equilibrium with the electrons on the left hand side
* the zinc equillibrium as the more negative electrode potential and undergoe oxidation losing electrons so the equilibrium i reversed to give the oxidation half equation with the electeons on the right hand side
* Zn(s)+Cu2+(aq)->Zn2+(aq) +Cu(s)
45
describe how to calculate a standard cell potential from standard electrode potentials
* electrode potential of the cell = electrode potential (positive electrode- electrode potential (negative electrode)
46
describ how to predict the feasilibility of any potential redox reactions using standard electrode potentials
* most negative system has the greatest tendency to be oxidised and lose electrons
* the most positive system has the greatest tedency to be reduced and gain electrons
* an oxidising agent takes electrons away from species being oxidised
* oxidising agents are reduced and are on the left
* a reducing agent adds electrons to the species being reduced
* so reducing agents are oxidised and are on the right
* the strongest reducing agent has a more negative electrode potential
* the strongest oxidising agnt is the more positve electrode potential
47
describe how you would write the overall equations of half equations with electrode potentials
* the redox ystem with the more positive electrode value will react from left to right andd gain lectrons
* the redox system with the more negative electrode value will react from right to left and lost electrodes
48
describe the effect of predicting electrode values on reaction rate
* one limitation of predictations for feasibilut based on gibbs free energy lies with reactions that have a very large activation energy result in a very slow rate
* so electrode potentials may indicate the thermodynaic feasiblity of a reaction but no indication on a rate of a rection
49
describe how concentration has limitations of predictions using electrode potential values
* standard electrode potentials measured using concentrations of 1moldm-3
* many reactions take place using concentrated or dilut solutions
* if concentrations of a solution is not 1moldm-3 then the value of the electroed potential will be different from the standard value
* if concentration greater on left the equilibrium will shift to the right removing electrons from the systm and making electrode potential less negative
* in concentrations less on the left the equilibrium will shift to the left increasing electrons in the system and making the electrode potential more negative
* any change in the electrode potential will affect the value of the overall cell potential
50
describe other factors which limit the predictions of using electrode potentials
* reaction rates
* concentrations
* the actual conditions used for reaction different from standard conditions
* affect value of the electrode potential
* standard electron potentials apply to aquous equilibria many reactions take place that are not aqueous
51
what are the three types of cells
* primary
* secondary
* fuel cells
52
what are primary cells
* non rechargeable and are designed to be used once only
* when in use the electrical energy produced by oxidation and reduction at the electrodes
* reactions cannot be reversed
* eventually chemicals will be used up voltage will fall the battery will go flat and cell discarded or recycled
53
where are primary cells found
* low current storage devices such as wall clocks and smoke detectors
54
what are most modern primary cells based on and give an equation
* alkaline based on zinc and manganese oxide zn/MnO2 and potassium hydroxide alkaline electrolyte
* Zn(s)+2MnO2->ZnO(s)+Mn2O3(s)
55
what are secondary cells
* rechargeable
* cell reaction producing electrical energy can be reversed during recharging
* chemicals in the cell are then regenerated and cell can be used again
56
what are three examples of secondary cells
* lead-acid batteries used in car batteris
* nickel - cadmium ,NiCd, cells and nickel-metal hydride, NiMH, The cyclindrival batteries used in radios, torches and so on
* lithium-ion and lithium ion polymer cells used in our moderm appliances - laptops, tablets, cameras, mobile phones alnd also being developed for cars
57
what are fuel cells
* uses energy from the reaction of a fuel with oxygen to create a voltage
* fuel and oxygen flow into the fuel cell and the products flow out the electrolyte remains i the cell
* fuel cells can operate continuously provided that the fuel and oxygen are supplied into the cell
* fuel cells don't have to be recharged
* fuel cells using many other hydrogen rich fuels such as methanol are also being developed
58
describe hydrogen fuel cells
* most common fuel cells
* produce no carbon dioxide during combustion
* with water being the only combustion produce
59
what are the cell voltages of hydrogen fuel cells in an alkali or acid electrolyte
1.23V
60
what are the redox systems for alklai hydrogen fuel cell and the overall equation
* 2H2O(l)+2e- reversible sign H2(g)+2OH-(aq)
* 1/2O2(g)+H2O(l)+2e- reversible sign 2OH-(aq)
* H2(g)+1/2O2(g)->H2O(l)
61
what are the redox systems for acid hydrogen fuel cells and the overall equation
* 2H=(aq)+2e- reversible sign H29g)
* 1/2O2(g)+2H+(aq)+2e- reversible sign H2O(l)
* H2(g)+1/2O2(g)->H2O(l)
62
what are the two types of hydrogen fuel cells
* either have an alkali or acid electrolyte
