13 — electrochemistry

Electrolysis is the process of using electricity to break down or decompose a compound (usually an ionic compound in the molten or aqueous state）

IE: - electrodes that do not undergo chemical changes n do not tk part in the electrolysis reaction - eg graphite, platinum RE: -metal anodes undergo oxidation during electrolysis - eg copper, silver

A molten binary ionic compound is typically a salt containing only one cation and one anion in the liquid state

Sodium ions gain electrons and is reduced to form sodium atom. Grey globules (Liquid Metal) of sodium obtained at cathode. Chloride ions loses electrons and is oxidised to form chlorine molecules. Yellow green chlorine gas obtained at anode. (Bromide: dark reddish-brown gas) Cathode: Na+ discharged Na+ (l) + e- -> Na (l) Na+ ions gains electrons and is reduced to form sodium atom. Grey globules of sodium obtained at cathode. Anode: Cl- discharged 2Cl- (l) -> Cl2 (g) + 2e- Cl- ions loses e- and is oxidised to form chlorine molecules. Yellow green chlorine gas obtained at anode. Overall equation: 2NaCl (l) -> 2Na (l) + Cl2 (g)

Advantages - high melting point - will not melt when used in the electrolysis of molten binary ionic compounds Disadvantages - graphite will react w oxygen gas under high temperatures to produce CO2 - graphite anodes might hv to be periodically replaced

Advantages - is an INERT ELECTRICAL CONDUCTOR and remains chemically unchanged hence does not tk part in the electrolysis reaction Disadvantages - lower melting point than graphite - might melt when used in the electrolysis of molten binary ionic compounds - mainly used in the electrolysis of aqueous electrolytes

4 OH- (Aq) -> O2 (G) + 2H2O (l) + 4e-

2H+ (aq) + 2e- ->H2 (g)

The ratio of H2 to Cl2 produced is 1:1 The solution becomes alkaline as there is a net discharge of H+ ions. (Concentrated hence Cl- gets discharged at the anode) The remaining Na+ and OH- ions form NaOH, an alkaline solution. When a few drops of Universal Indicator is added to the electrolyte, the UI changes from green to violet/purple. Cathode: H+ (aq) discharged 2H+ (aq) + 2e- -> H2 (g) Anode: Cl- (aq) discharged Since conc of Cl- ions > OH- ions, Cl- ions r SD to form yellow green chlorine gas. OE: 2H+ (aq) + 2Cl- (aq) -> H2 (g) + Cl2 (g) Th ratio of H2 to Cl2 produced is 1:1 Solution becomes alkaline as there is a net discharge of H+ ions. Remaining Na+ and OH- ions form NaOH, an alkaline solution. UI changes from green to violet.

After Cu^2+ and OH- ions r discharged, H+ (aq) and SO4 ^2- (aq). Ions remains in the solution. Hence, the resulting electrolyte becomes increasingly acidic as the concentration of H+ is greater than the concentration of OH-. Concentration of Cu^2+ ions decreases thus the blue colour of the electrolyte gradually fades and eventually turns colourless.

Electroplating allows us to coat a thin layer of metal onto an object Cathode is coated w a layer of copper metal and increases in mass Anode dissolves to form Cu^2+ ions and decrease in mass

The concentration of the electrolyte remains the same as 1 mol of Ag oxidises to form 1 mol of Ag+ ions at the anode and 1 mol of Ag+ ions reduce to form 1 mol of Ag at the cathode. Thus, there is no net change in concentration of Ag+ ions.

A simple cell is a device that converts chemical energy into electrical energy

More reactive metal: - acts as anode - oxidises forming cations that enter the electrolyte - releases electrons that flow thru external circuit Less reactive metal: - acts as cathode - causes cations from the electrolyte to gain e- and be reduced The further apart the 2 metals r in the reactivity series, the greater the voltage produced

Source: SC: electrical energy produced thru chemical reactions EC:electrical energy supplied by an external source Electron movement: SC: electrons move from the anode to the cathode EC: electrons move from the battery to the cathode, thru the electrolyte and into anode Polarity SC: Anode — negative Cathode: positive EC: Anode — positive Cathode — negative

Voltmeter is used SI: V

A fuel (hydrogen) is continuously added at the anode An oxidiser (oxygen) is continuously added at the cathode Cathode: O2 (g) + 2H2O(l) + 4e- -> 4OH- (aq) Anode: H2 (g) + 2OH- (aq) -> 2H2O (l) + 2e- Overall equation: 2H2(g) + O2(g) -> 2H2O (l)

Hydrogen is a renewable fuel and can be obtained via the electrolysis of water or from the cracking of hydrocarbon Hydrogen fuel cells produce only water as a by-product It is moe efficient than fuel-burning electricity sources. A larger percentage of the chemical energy stored in the fuel ends up as useful electricity in a fuel cell

It is difficult to store and transport hydrogen safely as hydrogen is a highly flammable gas at room temperature and pressure. Hydrogen is ofte transported in high-pressure cylinders which can be dangerous to handle. Large amount of energy is needed to produce hydrogen from electrolysis. However, this can be mitigated by using renewable sources like solar and wind energy to power electrolysis

• Connect the tin metal to the positive terminal of the cell to make it an anode. • Connect the iron can to the negative terminal of the cell to make it a cathode. • Immerse both electrodes in the electrolyte which is aqueous tin (II) nitrate solution.

When there is a difference in reactivity between 2 metals, a voltage is produced. The further apart the 2 metals r in the reactivity series, the greater the difference in reactivity between the metals, the greater the voltage produced.

The overall reaction is 2H2O (l) -> 2H2 (g) + O2 (g) 2 moles of H2O is broken down into 2 moles of H2 and 1 mole of O2. Since mole ratio is equal to volume ratio for gases, volume of H2:O2 theoretically evolved is 2:1.

As O2 is more soluble in water than H2, less O2 will be collected, causing the volume ratio to bebgreater than 2:1

As the electrolysis runs, lesser oxygen is dissolved in water as the solution becomes saturated w oxygen. Hence, more O2 is collected, bringing the ratio closer to 2:1