Chapter 23 quizlet pt 2

• what is an oxidising agent ? What is a reducing agent?

  oxidising agent- is the species that is reduced> it takes electrons from the species being oxidised

  reducing agent- the species that is oxidised> adds electrons to the species being reduced
• how do you write a

  half equation?1)assign oxidation numbersand determine change in oxidation number2)balancetheelectrons3) balance anyremaining atomsand predict further species> balance using H₂O, H⁺ and OH⁻ etc.
• how do you write ahalf equation:

  Example using MnO₄⁻ reduced to Mn²⁺ ions

  1) assign oxidation numbers and determine change in oxidation number MnO₄⁻ + H⁺ → Mn²⁺+7+2 decreased by 5

  2) balance the electrons MnO₄⁻ + H⁺ → Mn²⁺decrease by 5 means there must be 5 electrons on the left MnO₄⁻ + H⁺ +5e⁻→ Mn²⁺

  3) balance any remaining atoms and predict further species MnO₄⁻ + H⁺ + 5e⁻→ Mn²⁺there is 1 H and 4O on the left, none on right MnO₄⁻ +8H⁺ + 5e⁻→ Mn²⁺ +4H₂O✓
• how do you write a redox reaction

  from 2 half equations(of the reducing and oxidising agent)?1)balance electronsin each half equation2)addandcancel electrons3)cancel any speciesthat are on both sides of the equation
• how do you write a redox reaction

  from 2 half equations:example using hydrogen peroxide and Cr³⁺ metal1) balance electrons in each half equationH₂O₂ + 2e⁻ → 2OH⁻Cr³⁺ + 8OH⁻ → CrO₄²⁻ + 4H₂O + 3e⁻↓x33H₂O₂ + 6e⁻ → 6OH⁻x22Cr³⁺ + 16OH⁻ → 2CrO₄²⁻ + 8H₂O + 6e⁻2) add and cancel electronscancel the 6 electrons from each side3H₂O₂ + 2Cr³⁺ + 16OH⁻ → 2CrO₄²⁻ + 8H₂O + 6OH⁻3) cancel any species that are on both sides of the equation3H₂O₂ + 2Cr³⁺ +16OH⁻→ 2CrO₄²⁻ + 8H₂O +6OH⁻there are 16OH on left and 6 on right so cancel 6 of the 163H₂O₂ + 2Cr³⁺ +10OH⁻→ 2CrO₄²⁻ + 8H₂O
• turn for q

  
• turn for ans

  
• what is the

  overall redox equationbetween manganate ions with iron ions?MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 5Fe³⁺ +4H₂O
• how do you carry out a

  titrationfor the reduction ofmanganate ions?> prepare astandard solutionof KMnO₄> using a burette, add a measured volume of solution being analysed toconical flask>excess sulfuric acid, so produce excess H⁺ ions for reduction> do not need an indicator, it isself indicating> obtainconcordanttitres(0.1cm3)

  
• what is the

  end pointof themanganatetitration?thefirst permanent pink colour, this is when there is an excess of MnO₄⁻ ions present.

  
• how is the

  reading from the meniscustaken in themanganatetitration?burette readings aretaken from the top,rather than the bottom because of the deep purple colour it ishard to see.

  > this will have no effect on the results, as the titre is thedifferencebetween the 2 readings.
• question

  Metal M can be identified following the steps below. 1. The amount, in mol, of excess HCl(aq) that remains after the reaction of M with HCl(aq).2. The amount, in mol, of HCl(aq) that reacted with M.3. The identity of metal M.

  Analyse the results toidentify metal M. [6 marks]

  
• what is theoverall redox equation between iodine ions and thiosulfate ions?

  2S₂O₃²⁻ + I₂ → 2I⁻ + S₄O₆²⁻
• what are iodine/thiosulfate titrations

  used to determineand why?→ClO⁻concentration in bleach→Cu²⁺content in copper alloys

  because both react with iodide (I⁻) in a redox reactionwhere I₂ is formed↑ the I₂ used is titrated with thiosulfate
• how do you carry out a

  titration using iodine/thiosulfateto analyse content of oxidising agents?1)standard solution of Na₂S₂O₃(thiosulfate ions) to burette

  1. prepare solution ofoxidising agentto be analysed in conical flask, addEXCESS potassium iodide, this forms iodine, which forms a yellow/brown colour

  3)titrate with thiosulfate ions, iodine is reduced back to I⁻, causing the brown colour to fade.
• with theiosulfate/iodine reaction, how is end point determined?

  Using starch> when the solution no longer turns blue-black in the presence of starch, this means iodine is no longer present and has been reduced
• what is the

  overall redox equationbetweenClO⁻ions andI⁻ions?ClO⁻ + 2I⁻ + 2H⁺ → Cl⁻ + I₂ + H₂O
• what is the

  overall redox equationbetweenClO⁻ions andCu²⁺ions?2Cu²⁺ + 4I⁻ → 2CuI (s) + I₂
• how do you determine theconcentration of ClO⁻ ions in bleach? : worked example

  1) calculate mol of thiosulfate ions reacted

  2) find mol of I₂ reacted2 mol S₂O₃²⁻ = 1 mol I₂ divide by 2↑ this is the mol of I₂ formed in reaction with ClO⁻

  3) find mol of ClO⁻ in 25cm³1 mol I₂ = 1 mol ClO⁻

  4) scale up find mol in 250cm³

  5) find concentration of ClO⁻ ions mol/vol
• question

  Determine thepercentage, by mass, of chromiumin the ore. Give your answer to one decimal place. [6 marks]

  
• what is a

  voltaic cell?a cell thatconverts chemical energy into electrical energy, made from connecting 2 different half cells.

  
• what is a

  half cell?a part of a voltaic cell thatcontains the chemical species presentin a redox half-equation.

  
• why must the chemicals in the

  2 half-cells be kept apart?to allow the electrons to flow.if they are allowed to mix, electrons would flow in anuncontrolled wayand heat energy would be released instead than electrical energy
• what is a

  metal half celland how are theyrepresented?a metal rod dipped into a solution of its aqueous metal ion.→ represented with a vertical line for the phase boundary between the solution and metal.

  e.g. Zn²⁺(aq)|Zn(s)

  
• when is an

  equilibriumset up in a voltaic cell and how is this redox equilibriumwritten?equilibrium is set up at the phase boundary where themetal is in contact with its ions.

  it is written so the forward reaction shows reduction and backward reaction is oxidation
• what is an

  ion/ion half cell?contains ions of thesameelement indifferent oxidation states

  1. Fe³⁺(aq) + e⁻ → Fe²⁺ (aq)

  there is no metal to transport electrons, so a an inert metal electrode made of platinum is used.
• State the

  charge carriersthat transfer currenta) through the wireb) through the solution[1 mark]

  
• how do you know which electrode has a

  greater tendency to gain or lose electron?in a cell with 2 metal/metal ion half cells connected,the more reactive metal releases electrons more readilyand is oxidised.

  →this is thenegative electrodeas it is loses electrons.

  you can find the tendency to be reduced or oxidised with thestandard electrode potential.
• what is thestandard electrode potentialE⁰?

  The tendency to be reduced and gain electrons. Temperature of 298K→ pressure of 100kPa. The voltage measured under standard conditions when the half cell is connected to a standard hydrogen electrode

  
• Define the term

  standard electrode potential.Include all standard conditions in your answer. [2 marks]

  
• how is

  standard electrode potentialmeasured?using a salt bridge,thesalt bridgeallows ions to flow and contains a concentrated solution of electrolyte that does not react.

  
• A standard cell is set up in the laboratory with the cell reaction shown below. Ni(s) + I2(aq) Ni2+(aq) + 2I–(aq)

  (i) Draw alabelled diagramto show how this cell could be set up to measure its standard cell potential. Include details of apparatus, solutions and the standard conditions required. [4 marks]

  
• if the E⁰ value is

  more negativethan the otherE⁰ value, will the chemical beoxidised or reduced?the more negative theE⁰ value:→ thegreater tendencytoloseelectrons and undergooxidation→ theless the tendencytogainelectrons and undergoreduction

  metals tend to have negative E⁰values whereas non-metals tend to have positive E⁰ values
• how do you work out the

  E⁰ of a cell?E⁰ cell =E⁰ (positive electrode) -E⁰ (negative electrode)
• The reactions at each electrode are shown below.

  Al (s) + 4OH⁻ → (aq) Al(OH)₄ (g) + 3e⁻O₂ + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)

  (i) The standard electrode potential for the O₂/OH⁻ redox system is +0.40 V.

  The standard cell potential of an aluminium–oxygen cell is 2.71 V.

  What is thestandard electrode potentialof thealuminiumredox system in this cell? [1 mark]

  
• question

  An electrochemical cell can be made based on redox systems 1 and 2.

  Write down thestandard cell potentialof this cell. [1 mark]

  
• predictions from electrode potentials - look at notes in my folder because hard to write flashcards on
• how do you tell if a reaction is

  feasiblebased onelectrode potentials?if q was.. can Fe oxidise Ag to Ag+look at electrode potential values:Fe²⁺ + 2e⁻ ⇌ Fe =-0.44Ag⁺ + e⁻ ⇌ Ag =+0.8as Ag has agreater electrode potentialthan Fe, Ag will have a greater tendency to be reduced and gain electrons, so reaction isnot feasible.
• question

  Zinc reacts with acidified Cr2O7 2– ions to form Cr2+ ions in two stages.

  Explain why this happens interms of electrode potentials and equilibria.

  Includeoverall equationsfor the reactions which occur. [4 marks]

  
• Hydrogen is oxidised in a electrode potential reaction.

  Explain, in terms of electrode potentials and equilibrium, why the pH of the solution in thehydrogen half-cell decreasesas this cell delivers current. [2 marks]