> For the complete documentation index, see [llms.txt](https://jamesbrayy.gitbook.io/atar/llms.txt). Markdown versions of documentation pages are available by appending `.md` to page URLs; this page is available as [Markdown](https://jamesbrayy.gitbook.io/atar/chemistry/redox.md).

# redox

***

## **oxidation and reduction:**

* oxidation is the loss of electrons or gain of $$\mathrm{H}^{+}$$
* reduction is the gain of electrons or loss of $$\mathrm{H}^{+}$$ **redox rules:**
* *rule 1:*
  * substances in the elemental state **but not in a compound** have an oxidation number of 0
    * *e.g.* $$\mathrm{Cl\_2} = 0, \ \mathrm{C} = 0, \ \mathrm{P\_4} = 0$$
* *rule 2:*
  * monoatomic ions have an oxidation number equal to the charge of the ion
    * *e.g.* $$\mathrm{Na^{+}} = +1, \ \mathrm{Cl^{-}} = -1, \ \mathrm{Al^{3+}} = +3$$
* *rule 3:*
  * oxygen in compounds has an oxidation number of -2 **except** in peroxides
    * *e.g.* in $$\mathrm{H\_2O\_2},$$ $$\mathrm{H}$$ is +1 so $$\mathrm{O}$$ must be -1
* *rule 4:*
  * hydrogen has an oxidation number of +1 **except** in metal hydrides
    * *e.g.* in $$\mathrm{NaH},$$ $$\mathrm{Na}$$ is +1 so $$\mathrm{H}$$ must be -1
* *rule 5:*
  * the oxidation number of all atoms in a neutral molecule must sum to make the charge present on the ion 0
  * atoms that don't appear in the rules can have different oxidation number in different compounds

## **making equations:**

#### *in neutral conditions:*

1. write down the oxidant/reductant and what product it forms

* $$\mathrm{I\_2} \rightarrow \mathrm{I^-}$$

2. balance the atoms undergoing reduction/oxidation

* $$\mathrm{I\_2} \rightarrow 2\mathrm{I^-}$$

3. balance charges by adding $$e^-$$ where required

* $$\mathrm{I\_2} + 2e^- \rightarrow 2\mathrm{I^-}$$

4. add state symbols to the half equation

* $$\mathrm{I\_2 (s)} + 2e^- \rightarrow 2\mathrm{I^-} (aq)$$

#### *in acidic conditions:*

1. balance the atoms given first

* $$\mathrm{Cr\_2O\_7^{2-}} \rightarrow 2\mathrm{Cr^{3+}}$$

2. add in water molecules to balance the number of oxygen atoms missing

* $$\mathrm{Cr\_2O\_7^{2-}} \rightarrow 2\mathrm{Cr^{3+}} + 7\mathrm{H\_2O}$$
* $$\mathrm{HNO\_2} + \mathrm{H\_2O} \rightarrow \mathrm{HNO\_3^-}$$

3. add in $$\mathrm{H^+}$$ to balance the number of hydrogen atoms

* $$\mathrm{Cr\_2O\_7^{2-}} + 14\mathrm{H^+} \rightarrow 2\mathrm{Cr^{3+}} + 7\mathrm{H\_2O}$$
* $$\mathrm{HNO\_2} + \mathrm{H\_2O} \rightarrow \mathrm{NO\_3^-} + 3\mathrm{H^+}$$

4. balance charges by adding electrons to the most positive side of each half equation

* $$\mathrm{Cr\_2O\_7^{2-}} + 14\mathrm{H^+} + 6e^- \rightarrow 2\mathrm{Cr^{3+}} + 7\mathrm{H\_2O}$$
* $$\mathrm{HNO\_2} + \mathrm{H\_2O} \rightarrow \mathrm{NO\_3^-} + 3\mathrm{H^+} + 2e^-$$

5. to make an overall equation, multiply by a factor to make the number of electrons equal and combine

* $$\mathrm{Cr\_2O\_7^{2-}} + 14\mathrm{H^+} + 6e^- \rightarrow 2\mathrm{Cr^{3+}} + 7\mathrm{H\_2O}$$
* $$3\mathrm{HNO\_2} + 3\mathrm{H\_2O} \rightarrow 3\mathrm{NO\_3^-} + 9\mathrm{H^+} + 6e^-$$
* $$\therefore \mathrm{Cr\_2O\_7^{2-}} + 5\mathrm{H^+} + 3\mathrm{HNO\_2} \rightarrow 2\mathrm{Cr^{3+}} + 4\mathrm{H\_2O} + 3\mathrm{NO\_3^-}$$

6. to check, ensure oxidation numbers are equal on both sides

## **electrochemical cells:**

#### *galvanic cells:*

* ![](/files/10PTNeQcJuBTAu4pkedD)
* through the salt bridge:
  * anions move into oxidation half-cell
  * cations move into reduction half-cell
* electrons move through the wire from the anode to the cathode
  * the oxidation half-cell is the anode (-, an ox)
  * the reduction half-cell is the cathode (+, red cat)
* impacts on cell potential:
  * cell potential will vary depending on the chemistry of the half cells being used, the concentration of dissolved substances, pressure, and temperature
  * at STP, the cell potential becomes $$E^\circ$$

#### *standard hydrogen electrode:*

* ![](/files/xEba55HnQaPjFHy0TKzn)
* at STP, the standard reduction potential for the reduction of hydrogen ions is zero volts
  * $$2\mathrm{H^+}*{(aq)} + 2e^- \rightarrow \mathrm{H\_2}*{(g)}$$
* this 'reference' electrode enables standard reduction potentials to be determined for a range of substances

## **electrolysis:**

#### *electrorefining:*

* unlike electrolysis, electrorefining involves the oxidation of the anode itself
* in electrolysis, the anode is inert, but in this process it is involved in the reaction
  * the anode in electrorefining is made out of the material that is intended to be refined and the cathode can be made out of any highly-conductive substance like graphite and gold
* ![](/files/M8Nn90s8RKyNoXpgSAoI)

#### *electroplating:*

* involves depositing a thin layer of one metal over another
* thin metal layer may act as protection for underlying metal or look better
* ![](/files/GICAHrX6hvQwBjQqyqvK)

## **commercial cells:**

* a primary cell can only be discharged once and can never be recharged with electricity
  * they are easy and cheap to manufacture and use, and an example would be a zinc-carbon cell
* a secondary cell can be discharged and recharged repeatedly using electricity
  * an example would be a lead-acid cell
  * secondary cells are easily reusable as they can be recharged with electricity
* a tertiary cell or fuel cell requires a continuous source of fuel and oxygen in order to sustain a chemical reaction to produce electricity
  * an example would be a hydrogen fuel cell
  * tertiary cells can produce electricity continuously as long as fuel and oxygen are supplied

#### *lead-acid batteries:*

* the lead-acid battery consists of numerous banks of lead and lead dioxide plates in a sulfuric acid electrolyte
* they are the oldest example of a rechargeable battery (secondary cells)
* as they are rechargeable, there are two states:
  * *discharging (producing voltage)*
    * ![](/files/BxeXYKrX3jol7bXdZLog)
    * *at the anode:* $$\mathrm{Pb\_{(aq)}} + \mathrm{SO\_4^{2-}}*{(aq)} \rightarrow \mathrm{PbSO\_4}*{(s)} + 2e^-$$
    * *at the cathode:* $$\mathrm{PbO}*{2,(s)} + 4\mathrm{H^+}*{(aq)} + \mathrm{SO\_4^{2-}}*{(aq)} + 2e^- \rightarrow \mathrm{PbSO\_4}*{(s)} + 2\mathrm{H\_2O}\_{(l)}$$
    * *overall:* $$\mathrm{Pb}*{(s)} + \mathrm{PbO}*{2,(s)} + 4\mathrm{H^+}*{(aq)} + \mathrm{SO\_4^{2-}}*{(aq)} \rightarrow 2\mathrm{PbSO}*{4,(s)} + 2\mathrm{H}*{2}\mathrm{O}\_{(l)}$$

#### *hydrogen fuel cells:*

* the hydrogen fuel cell is an electrochemical cell that can operate perpetually because the reactants required for the redox reaction can be constantly replaced
* the reactions can take place in either an acidic or alkaline electrolyte
* *in acid electrolyte:* $$\mathrm{H\_{2,(g)}} \rightarrow 2\mathrm{H^+*{(aq)}} + 2e^- \quad \left( E^0 = 0.00\mathrm{V} \right)$$ $$\mathrm{O*{2,(g)}} + 4\mathrm{H^+*{(aq)}} + 4e^- \rightarrow 2\mathrm{H\_2O*{(l)}} \quad \left( E^0 = +1.23\mathrm{V} \right)$$
* *in alkaline electrolyte:* $$\mathrm{H\_{2,(g)}} + 2\mathrm{OH^-*{(aq)}} \rightarrow 2\mathrm{H\_2O*{(l)}} + 2e^- \quad \left( E^0 = -0.83\mathrm{V} \right)$$ $$\mathrm{O\_{2,(g)}} + 2\mathrm{H\_2O\_{(l)}} + 4e^- \rightarrow 4\mathrm{OH^-\_{(aq)}} \quad \left( E^0 = +0.40\mathrm{V} \right)$$