Chapter 9: Redox Processes Flashcards
Define redox reaction
When oxidation and reduction occur simultaneously
Define oxidation:
- ___ of oxygen
- ___ of hydrogen
- ___ of electrons
- ___ in oxidation number
- Gain of oxygen
- Loss of hydrogen
- Loss of electrons
- Increase in oxidation number
Define reduction:
- ___ of oxygen
- ___ of hydrogen
- ___ of electrons
- ___ in oxidation number
- Loss of oxygen
- Gain of hydrogen
- Gain in electrons
- Decrease in oxidation number
What is a disproportionation reaction?
When the same reactant is both oxidised and reduced
How do you assign oxidation numbers to:
- Elements
- Elements in an ionic salt
- Elements in a covalent compound
- Elements in an ionic compound
- Elements: O.N always 0
- Elements in an ionic salt: O.N = Ionic charge
- Elements in a covalent compound: Assign O.N as if ionic (most electronegative element is assigned negative charge)
- Elements in an ionic compound: Sum of O.N of elements = Total charge of compound
What are the exceptions to rules for O.N?
- Transition metals: Can have multiple O.N, and are named using oxidation state (roman numerals)
- O.N of H in metal hydrides (Eg. NaH): -1
- O.N of O in OF2: +2 (because F is more electronegative)
- O.N of O in hydrogen peroxide (H2O2): -1
Define oxidising and reducing agents
- Oxidising agent: A species that oxidises another species (takes electrons), and is themselves reduced in the process (gain electrons)
- Reducing agent: A species that reduces another species (donates electrons), and is themselves oxidised in the process (lose electrons)
What are half equations?
Equations that show oxidation and reduction separately
How are half-equations balanced and re-combined to form an overall redox equation?
- Balance number of (non H and O) atoms
- Balance O by adding H2O to the O-deficient side
- Balance H by adding H+ to the H-deficient side
- Balance charges by adding electrons to the electron-deficient side (more positive side)
- Recombine half-equations by multiplying entire equation to balance number of electrons (no. of e- lost during oxidation = no. of e- gained during reduction)
- Subtract any duplications
Define the activity series
Arranging metals according to how easily they are oxidised (how readily they lose electrons)
More active = more easily oxidised = loses electrons more readily = stronger reducing agent
*Link to chemical bonding: first ionisation energy
How can the activity series be used to predict the feasibility of redox reactions?
More reactive metals will displace less reactive ones to create displacement redox reactions, as they are stronger reducing agents and will reduce the less reactive metal while losing electrons more readily to be oxidised
When reacting with an acid (H+), metals more reactive than H will reduce H+ to H2 gas (*Link to Acids and Bases: acid/reactive metal reaction)
Give an example of a redox titration
Winkler method
S₂O₃ : O₂
4 : 1
Define Biochemical Oxygen Demand (BOD)
Amount of oxygen used by aerobic microorganisms to decompose the organic matter in water over 5 days at 20°C
BOD (mg dm⁻³/ppm) = Initial concentration of dissolved oxygen - Concentration of dissolved oxygen after 5 days
(mol dm⁻³ x mR = g dm⁻³)
(g dm⁻³ x 1000 = mg dm⁻³ = ppm)
Example question (Winkler method):
MnSO₄ , KI and H₂SO₄ were added to 100cm³ of water. Iodine formed was titrated against 16.00 cm³ of 5.00x10⁻³ mol dm⁻³ of Na₂S₂O₃. What is the BOD of the water sample?
- Find ₙ[Na₂S₂O₃]
- Find ₙ[O₂]
- Find [O₂]
- Find mass of O₂
- Convert from g dm⁻³ to ppm
Describe the energy conversion in a voltaic cell
Conversion of chemical energy (redox reactions) to electrical energy (voltage) through displacement reactions (electrons lost are transferred through the external circuit)
