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Electrochemistry = The study of electron movement in an oxidation or reduction reaction at an electrode surface.

Two main ways in which they interact:

1. Echem -> Eelectr. Some chemical reactions can generate electricity.

2. Eelectric -> Echem. Electricity can also cause a chemical reaction to happen that would not occur spontaneously.

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ELECTRODE CHARGES IN GALVANIC VS ELECTROLYTIC CELLS

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ELECTRODE CHARGES IN GALVANIC VS ELECTROLYTIC CELLS

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Electrolytic Cells:

In an electrolytic cell, an external power source drives a non-spontaneous chemical reaction. Here, the anode is the positive electrode, and the cathode is the negative electrode.

Anode: Positive electrode where oxidation occurs. It attracts anions (negative ions) from the electrolyte because opposite charges attract.

Cathode: Negative electrode where reduction occurs. It attracts cations (positive ions).

Voltaic (Galvanic) Cells:

In a voltaic or galvanic cell, a spontaneous redox reaction generates an electric current. Here, the roles of the anode and cathode are determined by the direction of the spontaneous chemical reaction.

Anode: Negative electrode where oxidation occurs. In a galvanic cell, it is negative because it releases electrons into the external circuit. Anions in the electrolyte do not move towards the anode in this case. Instead, the movement of ions in the electrolyte is more about maintaining electrical neutrality rather than being attracted to the electrodes.

Cathode: Positive electrode where reduction occurs. It attracts cations from the electrolyte.

Key Differences:

Electrode Charge: In an electrolytic cell, the anode is positive and attracts anions, while in a galvanic cell, the anode is negative.

Direction of Current: In an electrolytic cell, the current is driven by an external power source, whereas in a galvanic cell, the current is generated by the chemical reaction.

In summary, in a galvanic cell, the anode does not attract anions because it is negatively charged, and the movement of ions is primarily driven by the need to balance charge during the redox reaction, rather than attraction to the electrodes.

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Recap of previous section on electrochemistry.

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REDOX Reactions Tutorial — The Crash Chemistry Academy

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Metals (and non-metals) can be arranged in a reactivity series based on the reactions of the metals and their compounds.

The reactivity series lists elements (mostly metals) in order of decreasing reactivity.

Understanding the reactivity series is fundamental to chemistry.

This is the order of METALS in decreasing reactivity, below:

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An original Song by Adam Goddard Anna Nash Laurajoy Newman Disha Naik We do not own the rights to the Star wars theme tune.

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A COMPARISON - using Group 1 metals (alkali metals).

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Spontaneous reaction between Zn and CuSO4 solution.

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Spontaneous reaction between Zn and CuSO4 solution.

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Reaction is SPONTANEOUS when the stronger reducing agent react with the stronger oxidizing agent.

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The Reaction of Zinc Metal and Aqueous Copper(II)

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An investigation of the reaction between magnesium and copper oxide 

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An investigation of the reaction between magnesium and copper oxide 

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Thoroughly mix small amounts of magnesium powder and copper oxide powder.

Place the mixture on a bottle top.

Place the bottle top on a pipe-clay triangle resting on a tri-pod.

Place a medium Bunsen flame below the bottle top and stand back.

Observations:
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6. The reaction between aluminium and iron oxide is used to join lengths of railway track.
It is called the THERMIT REACTION:
Fe ₂ O ₃(s)   +   2 Al(s)  -->  Al ₂ O ₃(s)   +   2 Fe(l)
a. Why does aluminium react with iron oxide?
b. What does (l) after the Fe in the equation mean?
c. Suggest why the thermit reaction can be used to join lengths of railway track.

7. Metal X displaces metal Y AND Z from solutions of their compounds. Metal Z displaces metal Y from solutions of its compounds, but not metal X from solutions of its compounds. Place X, Y and Z in order of reactivity.

8. Carbon will reduce lead (II) oxide to lead and zinc oxide to zinc. a. Write the chemical equations for the two reactions. b. Which of the two oxides is reduced more easily? Give a reason.

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NEW CONCEPT INTRO: TWO TYPES OF ELECTROCHEMICAL CELLS:

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NEW CONCEPT INTRO: TWO TYPES OF ELECTROCHEMICAL CELLS:

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1. Galvanic cell
2. Electrolytic cell

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First: We will study Galvanic cells.
Second: We will study Electrolytic cells.

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1. GALVANIC CELL:                                               

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1. GALVANIC CELL:                                               

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The flow of electrons between species in a redox reaction can be harnessed through a wire connection, as shown below.

This is the setup for a Galvanic cell.

It converts chemical energy into electrical energy.

This makes a simple battery.

We call it a GALVANIC CELL

It is also called a VOLTAIC CELL

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THE ORIGIN OF THE GALVANIC CELL - Luigi Galvani and animal electricity PART 1

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THE ORIGIN OF THE GALVANIC CELL - Luigi Galvani and animal electricity PART 1

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Galvani VS. Volta - who was right? (#2 of 2)

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VOLTAIC (GALVANIC) CELLS VS. ELECTROLYTIC CELLS

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VOLTAIC (GALVANIC) CELLS VS. ELECTROLYTIC CELLS

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The CURRENT that an electrochemical cell can deliver is increased by:

INCREASING THE RATE OF THE FORWARD REACTION, WHICH RELEASES MORE ELECTRONS AT THE ANODE AND ACCEPTED AT THE CATHODE.

increasing the surface area (size) of the electrodes

increasing the diameter of the salt bridge and decreasing its length  (making it wider and shorter)

increasing the concentration of the salt solution in the salt bridge.

The EMF (cell Voltage) is increased by:

Increasing the concentration of REACTANTS.
Decreasing the concentration of PRODUCTS.
This favours the forward reaction, which transfers more electrical energy in the form of more electrons moving from anode to cathode.
HOWEVER - EMF IS NOT AFFECTED BY:
* Increased electrode surface area
* Wider, shorter, more conductive salt bridge             * Because those changes do not change the
  concentration of species in solution.
* Therefore, it does not affect the equilibrium
  position to favour either fwd or reverse
  reaction in the net ionic reaction.

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KAC32.7 - Electrochemistry: Effect of Concentration on Cell Potentials

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Standard hydrogen electrode

The electrode potential of the hydrogen electrode at 25℃ is estimated to be 4,4 V.

However, in order to use this as a reference electrode this value is set to zero at all temperatures so that it can be compared with other electrodes

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Reduction Potential Table Voltage Values

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Reduction Potential Table Voltage Values

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the Voltage value shows how easily the element is reduced.

A negative Voltage reduction potential value means:

negative value = DIFFICULT to be reduced.

the other element or compound has a greater tendency to REDUCED.

A positive value means:

positive value = EASY to be reduced.

The other element or compound has a greater tendency to be OXIDISED.

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2. ELECTROLYTIC CELL:                                                    E electrical  —> E chemical.                                       (Using electricity to make a chemical reaction occur)

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2. ELECTROLYTIC CELL:                                                    E electrical  —> E chemical.                                       (Using electricity to make a chemical reaction occur)

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Electrolysis rule:

At the negative electrode, the least reactive species (from the possible cations in solution) will be reduced into atoms from ions.

Because hydrogen is less reactive than sodium, the hydrogen is discharged as H atoms that then bind to form H₂-molecules.

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Electrochemical cell vs. Electrolytic cell

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How Does Electroplating Work | Reactions | Chemistry | FuseSchool

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Electrolysis is used to electroplate objects (coat them with a thin layer of metal). This is useful for coating a cheaper metal with a more expensive one, such as copper or silver.

The negative electrode should be the object to be electroplated.

The positive electrode should be the metal that you want to coat the object with.

The electrolyte should be a solution of the coating metal, such as its metal nitrate or sulfate.

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USES OF ELECTROLYTIC CELLS - ELECTROPLATING

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USES OF ELECTROLYTIC CELLS - ELECTROPLATING

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Electroplating = a method of using an electric current to deposit a thin layer of metal over another from a solution containing a dissolved salt of the selected metal.

NOTE:
The anode is connected to the positive terminal:

The anode is generally composed of the material that we want to deposit.

E.g. Copper in the case of copper plating, nickel in the case of nickel plating.

It is also possible to plate using “inert” anodes like carbon rods in certain cases.

The cathode connected to the negative terminal:

The cathode is the component on which we intend to deposit the metal.

In a plating workshop this is commonly referred to as the “work-piece”.

The English physicist Michael Faraday (1791 – 1867) provided us with the important rules that govern electro-deposition:

He showed that the amount of metal deposited at the cathode and the amount dissolved at the anode are directly proportional to the quantity of electricity passed, in other words, to the current and time period of the electro-deposition.

Thanks to Faraday's efforts and Faraday's Constant (F = 96 500 C), we can calculate the deposit mass of any metal that we are plating based on its atomic mass and the number of electrons involved in the electrochemical reaction. The formula use is:

n = q/F  (where n=number of moles of e⁻ transferred between electrodes, q=charge, F=96500 C)

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NO3 - ions would be reduced in preference over Cu2+ ions in the electrolyte solution.

So, to electroplate a medal with Copper, do not use copper(II) nitrate as electrolyte.

Use CuSO4 (aq)

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ELECTRONS ARE TRANSFERRED FROM ANODE TO CATHODE IN ELECTROCHEMICAL CELLS

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ELECTRONS ARE TRANSFERRED FROM ANODE TO CATHODE IN ELECTROCHEMICAL CELLS

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Can we work out the mass of metal produced at the cathode, if we know how many moles of electrons have moved across?

YES!

We call this QUANTITATIVE ELECTROCHEM.

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Counting electrons moving across and calculating mass of metal deposited

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USE FARADAY'S CONSTANT TO CALCULATE MASS OF METAL DEPOSITED AT CATHODE

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USE FARADAY'S CONSTANT TO CALCULATE MASS OF METAL DEPOSITED AT CATHODE

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Use:

Q = It

n = q/F

m = n.Mr

See below:

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Quantitative electrolysis | Redox reactions and electrochemistry | Chemistry | Khan Academy

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THE GALVANIC CELL

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Q: Aluminium extraction:
1. What energy conversion takes place in the type of cell used in aluminium extraction?
2. Write the half reaction occurring at the anode.
3. Write the half reaction occurring at cathode.
4. Explain what happens with the graphite anodes during operation, and write the reaction to state this.
Also write the overall cell reaction down.
5. What is the advantage to using cryolyte here?
6. Give a disadvantage of using cryolite in this process.
7. Al2O3 is insoluble in water. You decide to extract aluminium from aqueous AlCl3 and electrolyse the solution.
Would you be successful in extraction Aluminium? w
Would aluminum ions or water be reduced at the cathode? Explain.

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There are three major types of cells used in the electrolysis of brine:

the mercury cell

the diaphragm cell

the membrane cell

The PURPOSE of the chlor alkali industry is to make:

Chlorine

Hydrogen gas

Sodium hydroxide (caustic soda)

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When a current is run through a brine solution (a high-concentration of NaCl and H ₂ O solution) in an electrolysis process, we can extract these products:

chlorine gas

hydrogen gas

sodium hydroxide (caustic soda)

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BRINE SOLUTION - IONS AND MOLECULES PRESENT

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BRINE SOLUTION - IONS AND MOLECULES PRESENT

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BRINE = AN ELECTROLYTE
ELECTROLYTE = Solution containing free ions and which can conduct electricity

Why do you think brine can conduct electricity?

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ANIMATION OF EACH CELL IN THE CHLOR-ALKALI INDUSTRY

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ANIMATION OF EACH CELL IN THE CHLOR-ALKALI INDUSTRY

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The Chlor-alkali Industry: by Animations

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Electrolysis of brine - general reaction

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Electrolysis of brine - general reaction

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Electrolysis of Brine | Reactions | Chemistry | FuseSchool

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Three cells used in industry to electrolyse brine:

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Three cells used in industry to electrolyse brine:

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Mercury Cell

Diaphragm Cell

Membrane Cell

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IC#16 Industrial production of sodium hydroxide

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IC#18 Industrial production of sodium hydroxide membrane process

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Animation

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Exercise

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What Is The Electrochemical Series | Reactions | Chemistry | FuseSchool