Chapter 23: Redox Flashcards
What does an oxidising agent do, and what other feature does it have in redox reactions?
- It takes electrons from the species being oxidised
- It contains the species being reduced
What does an reducing agent do, and what other feature does it have in redox reactions?
- It gives electrons to the species being reduced
- It contains the species being oxidised
What needs to be remembered when naming the oxidising and reducing agent?
- Oxidising/ reducing agents are not just the element that is reduced/ oxidised, but the whole molecule/ ion
What has to happen for a reaction to be classed as a redox reaction?
- There has to be a change in oxidation number
How can redox reactions be constructed?
- By using half equations
- By using oxidation numbers
What needs to be done when using half equations to construct redox reactions?
- Write the half-equations for oxidation and reduction, making sure both are balanced
- The number of electrons in both equations needs to be the same so the equations can be combined
- This means one or both need to be scaled up
How can oxidation numbers be used to write redox equations?
- You find which species had a change in oxidation number, and by how much
- You scale up the other species (wherever it appears in the reaction- watch out for any that are diatomic) so that the overall change in oxidation number sums to 0
What can be done to balance redox reactions that take place in acidic, aqueous conditions?
- Start by balancing elements other than hydrogen and oxygen (especially any with a change in oxidation number)
- Balance oxygen by adding water (take note of any oxygens already present)
- Balance hydrogens by adding protons
- Add electrons to balance out charges if necessary (equations also have to be balanced in terms of charge)
What can be done to balance redox reactions that take place in alkaline, aqueous conditions?
- Add water molecules to balance oxygens
- Add protons to balance hydrogens
- Add the same number of protons in hydroxide ions to both sides of the equation
- Add protons and hydroxide ions together to form water
- Add electrons to balance the equation
If a redox reaction is missing a product, how can this product be found?
- It is usually either water, a proton or a hydroxide ion
How can the concentrations of reactants in a redox reaction be calculated?
- Through redox titrations
Give 2 examples of redox titrations used for analysis.
- Fe2+/MnO4–
- I2/S2O32−
Describe how to carry out a titration between manganate (VII) ions and iron (II) ions.
- Standard solution of potassium manganate (VII) is added to the burette.
- A pipette is used to add a measured volume of the reducing agent, such as a solution containing iron (II), to the conical flask with an excess of dilute sulfuric acid (as it is needed for the reduction of MnO4-).
- No indicator is needed as the reaction is self-indicating; manganate (VII) is purple while manganese (II) is colourless.
- The end point is judged by the first permanent pale pink colour, as it shows an excess of manganate (VII) is present.
- Burette readings are taken from the top of the meniscus, as potassium manganate (VII) has a deep purple colour that makes it hard to take readings from the bottom of the meniscus.
- Repeat until you obtain concordant results.
Give the equation for the reaction of manganate (VII) and iron (II). How can it be derived, and why is it important for their titration?
- Manganate (VII) is reduced to manganese (II), and iron (II) is therefore oxidised to iron (III)
- The reactions for both equations can be written, balanced and combined (iron’s would need to be scaled up by a factor of 5)
- This tells you they react in a 1:5 ratio
- MnO4- + 5Fe2+ + 8H+ -> Mn2+ + 5Fe3+ + 4H2O
Give the equations for oxidation, reduction and the overall equation in the reaction between iodine and thiosulfate ions.
- Oxidation: 2S2O3^2- -> S4O6^2- + 2e-
- Reduction: I2 + 2e- -> 2I-
- Overall: 2S2O3^2- + I2 -> 2I- + S4O6^2-
How do you carry out an iodine/ thiosulfate redox titration?
- Add a standard solution of sodium thiosulfate to the burette.
- Prepare a solution of an oxidising agent that you want to analyse, and add it to a conical flask along with an excess of potassium iodide (to supply iodide ions). This will make the solution yellow-brown as iodide is oxidised to iodine.
- Add the sodium thiosulfate which reduces iodine to iodide, making the yellow-brown colour fade.
- When the solution is pale yellow, add starch to make the end point easier to identify. This turns the solution blue-black.
- The end point is reached when the blue-black colour disappears.
How can the concentration of chlorate ions in bleach be found?
- Make the chlorate ions into a standard solution
- Pour the solution into a conical flask with potassium iodide and anything else needed for the production of iodine
