Electrochemistry Flashcards
What is oxidation number a measure of?
- The hypothetical charge an atom would have if all bonds to atoms of different elements were completely ionic
- It measures degrees of oxidation and reduction and is related to electronegativity
Oxidation number is represented by +(n) or -(n), or as a roman numeral after the species, e.g. iron (II) oxide
How can the oxidation number of monoatomic ions be determined?
- For monoatomic ions, it denotes the number of electrons lost or gained, so is equal to the charge?
- For example, a chloride ion has an oxidation number of -1
Which element in a compound is given the negative oxidation number?
- The more electronegative element
- For example, each phosphorus atom in P₂O₅ has an oxidation number of +5 as each oxygen atom has an oxidation number of -2
What is the oxidation number of uncombined elements?
- 0
- For example, each oxygen atom in O₂ has an oxidation number of 0
What are the oxidation numbers of group 1, group 2 and group 17 elements in compounds?
- Group 1: +1
- Group 2: +2
- Group 17: -1
There are very few exceptions
What does the oxidation number of hydrogen in compounds tend to be?
- +1
- However, it can vary depending on the electronegativity of the atom it is bonded to
- For example, the hydrogen in NaH has an oxidations state of -1 as sodium is electropositive and the hydrogen in F₂H has an oxidation state of +2 as fluorine is very electronegative
What does the oxidation number of oxygen in compounds tend to be?
- -2
- However, it often adopts an oxidation number of -1 when bonded to hydrogen, like in H₂O₂
- It has a positive oxidation number when bonded to fluorine, the only atom more electronegative than oxygen
What is the sum of oxidation numbers in a compound?
0
What is the sum of oxidation numbers in a polyatomic ion?
The charge on the ion
What are redox reactions?
Reactions where oxidation and reduction take place
Oxidation numbers change accordingly
What are disproportionation reactions?
- Reactions where atoms of an element are both oxidised and reduced to form atoms of the same element with higher and lower oxidation numbers
- For example, in the reaction 2H₂O₂ → 2H₂O + O₂, oxygen goes from a -1 oxidation state to a -2 and 0 oxidation state
- The opposite is comproportionation
- Species do not necessarily need to lose electrons, but the electronegativity of the atoms they are bonded to does
What is the balanced version of this half-equation?
14H⁺ + Cr₂O₇²⁻ + 6e⁻ → 2Cr³⁺ + 7H₂O
Assuming water is formed from oxygen and balancing with hydrogen is conventional for aqueous reactions and ensures both sides have equal charge
What is the balanced version of this half-equation?
VO²⁺ + H₂O → VO₂⁺ + 2H⁺ + e⁻
What is the result of combining and simplifying these two half-equations?
- FeO₄²⁻ + 8H⁺ + 3e⁻ + H₂O₂ → Fe³⁺ + 4H₂O 2H⁺ + O₂ + 2e⁻
- 2FeO₄² + 16H⁺ + 6e⁻ + 3H₂O₂ → 2Fe³⁺ + 8H₂O + 6H⁺ + 3O₂ + 6e⁻
- 2FeO₄² + 10H⁺ + 3H₂O₂ → 2Fe³⁺ + 8H₂O + 3O₂
- The result is the full ionic equation for this redox reaction, not just the half equation
- Ensure you understand why each step is taken
What is the difference between an oxidising agent and a reducing agent?
- An oxidising agent is a substance that oxidises another atom or ion by causing it to lose electrons and is reduced itself
- A reducing agent is a substance that reduces another atom or ion by causing it to gain electrons and gets oxidised itself
- The reducing or oxidising agent will be the entire species, not just an atom within a compound
- For some species, the context of the reaction determines whether it acts as an oxidising or reducing agent
- The transfer of electrons isn’t necessary, only an alteration in oxidation state
