Chemistry - Electrochemistry Exam

Electrolytic Cell

It uses electrical energy to produce chemical change

Electrolysis

It involves forcing a current through a cell to produce a chemical change. It is used to obtain active elements like sodium and chlorine by the electrolysis of fused (molten) salts. The cation is reduced, and the anion is oxidized.

Ex. Electrolysis of NaCl

In the electrolysis of molten NaCl, sodium chloride is melted so that it forms free-moving Na⁺ and Cl⁻ ions. A direct current is supplied by an external power source, which forces electrons to move through the circuit. At the cathode (negative electrode), Na⁺ ions are attracted and gain electrons, forming sodium metal. At the anode (positive electrode), Cl⁻ ions are attracted and lose electrons, forming chlorine gas. The power supply continuously removes and supplies electrons to maintain this electron flow. As a result, molten NaCl is decomposed into sodium and chlorine because electricity forces a non-spontaneous redox reaction to occur.

How to write the net ionic equation electrolytic equation?

• Make sure they are balanced

• Remove spectator ions

• Keep phases of matter

How do you identify what goes through oxidation or reduction in electrolytic cell?

• Identify the ions

• Decide what gets reduced or oxidized

Metal Plating

is when a thin layer of one metal is coated onto another object using electrolysis.This helps protect metals that readily corrode when applied as a coating.

• You use electrolysis in a solution containing ions of the coating metal.

Anode (positive)

• Made of the coating metal

• It dissolves into solution

• Metal atoms lose electrons → form ions

Cathode (negative)

• The object being plated (e.g. spoon, ring, key)

• Metal ions in solution gain electrons

• They become solid metal and stick onto the object

Solution: Aqueous solution containing ions of the metal being plated

Corrosion (Oxidation)

is the gradual destruction of a metal due to oxidation (loss of electrons), usually when it reacts with oxygen and water.

Preventing corrosion - rusting of iron

The corrosion of iron is affected by contact with another metal. It rusts more quickly in contact with a less active metal and more slowly in contact with a more active metal.

• Attach a more active metal to it (cathodic protection) - sacrificial anode

Redox Reaction

Electrons lost = electrons gained. Electrons cannot be created or destroyed — every electron one species loses must be gained by another species.

Electrochemistry

The study of the relationship between chemical change and electrical work.

What is the difference between the types of electrochemical cells?

• One’s overall overall redox reaction is spontaneous

• The other one’s overall redox reaction is non-spontaneous

Voltaic Cell (Galvanic Cell)

It uses a spontaneous reaction, that is they convert chemical energy to electrical energy. All batteries contain voltaic cells.

Electrolytic Cell

It uses electrical energy to drive a non-spontaneous reaction. They convert electrical energy into chemical energy.

Breakdown of Voltaic Cell:

• In any voltaic cell, the components of each half-reaction are placed in a separate container or half-cell.

• Each half cell contains an electrode that conducts electricity, which can be measured using a voltmeter

• The electrode is dipped into an electrolyte that contains the cation (positively charged ion) of the same metal as the electrode

• A salt bridge is required for the proper functioning of the cell

Anode

Is a electrode where oxidation occurs

Cathode

Is a electrode where reduction occurs

What happens in the half cell during oxidation?

Zinc solid dissolves into solution → forms Zn²⁺ ions

Concentration of Zn²⁺ increases in the electrolyte

The mass of the zinc electrode decreases (because solid zinc is being used up)

Electrons are released and flow through the wire to the other half-cell

What happens in the half cell during reduction?

• Copper ions (Cu²⁺) in solution gain electrons

• They turn into solid copper metal

• The solid copper plating builds up on the electrode

• The mass of the cathode increases

• The concentration of Cu²⁺ in solution decreases (because ions are being removed)

How do electrons flow in a voltaic cell?

They flow from ANODE to CATHODE. They go through the wire which creates a electric current.

Why do electrons flow at all? This because their is a difference in tendency to gain or lose electrons.

Salt Bridge

A salt bridge keeps both half-cells electrically neutral, so the reaction can keep going.

🔴 At the anode (oxidation)

• Metal atoms lose electrons and form positive ions in solution

• So the solution becomes too positive

• To balance this, negative ions (anions) from the salt bridge move in

🔵 At the cathode (reduction)

• Positive ions in solution gain electrons and form solid metal

• So positive ions are removed, leaving the solution too negative

• To balance this, positive ions (cations) from the salt bridge move in

**Without it, the buildup of charge would stop electron flow meaning the redox reaction would stop

How do you choose whats in a salt bridge?

It should contain an inert electrolyte, meaning:

• Does NOT react with metals or ions in the cell

• Dissociates completely into ions

• Allows ions to move to balance charge

1. Must be chemically unreactive

• It should not form a precipitate

• It should not be oxidized or reduced

• It should not interfere with the redox reaction

3. Must be soluble ionic compounds

So it can easily release ions in solution.

4. Must provide spectator ions

What is a half reaction?

Half-Reactions are equations that describe the changes in only the compound that is oxidized or teh compound that is reduced.

Oxidation Half-Reaction

X (s) yields X^+ (aq) and e^-

• Atom to Ion

Reduction Half-Reaction

e^- + Y^+(aq) yields Y(s)

• Ion to Atom

What is the net ionic equation for Reducation and Oxidation Reactions?

X (s) + Y^+ (aq) yields X^+(aq) + Y (S)

Voltage

IT comes form how strongly the reaction wants to happen.

Oxidation-reduction/Redox reactions

These reactions involve the transfer of electrons from one reactant to another. The movement occurs from the reactant with less attraction for electrons to the reactant with more attraction for electrons.

Oxidation

A process where a substance loses electrons, causing the oxidation number to increase

Reduction

A process where a substance gains electrons, causing the oxidation number to decrease

Why are oxidation numbers important?

They help you determine whether electrons were transferred during a reaction (redox). Helped identify substances that were oxidized vs reduced.

Oxidation Rule for Atom in Elemental Form (Element by itself)

Oxidation Number = 0

Oxidation Rule for Monoatomic Ion

The Oxidation Number = ion charge (with the sign in front of it)

Oxidation Number values for all the atoms in a molecule or formula unit of a compound (neutral)

Oxidation number =0

The sum of O.N values for the atoms in a polyatomic ion

= the ion’s charge

Group 1A (1) and 2A (2) Oxidation Numbers

(1) O.N = +1 in all compounds

(2) O. N = +2 in all compounds

Hydrogen Oxidation Numbers

+1 = in combination with nonmetals

-1 = in combination with metals and boron

Oxygen Oxidation Numbers

-1 = peroxides

-2 = in all other compounds (except with F)

F = +2 (positive and it is more electronegative)

Group 7A (17)

O.N = -1 in combinaiton with metals, nonmetals (except O), and other halogens lower in group.

Oxidizing agent

• Gets reduced

• Causes another substance to be oxidized

• It accepts electrons

Reducing agent

• Gets oxidized

• Causes another substance to be reduced

• It gives away electrons

General Rule: Oil Rig

Oxidation is Loss of Electrons, Reduction is gain of Electrons

Oxidation + Activity Series

Any metal that is above another metal on the activity series is more reactive and more readily oxidized.

• Just like how in single replacement reaction, a more reactive metal will replace a less reactive metal.

Redox vs Non-redox Reactions

Redox: Combustion, Decomposition, and Single Replacement.

• Some combustion and decomposition reactions can be non-redox reactions

NonRedox: Double-replacement (neutralization) + Double-replacement (precipitation)