# data[].attributes.body
1. movement of electrons always occurs in the external circuit

2. less reactive metal, harder to oxidise, easier to reduce, poor reducing agent

3. more reactive metal, easier to oxidise, harder to reduce, good reducing agent

4. half equation has to be in reduction form

5. flipping the equation changes the sign of the E standard

6. electric potential cannot be measured directly, instead we can compare the difference between a metal/metal ions system and another system (normally standard hydrogen electrode which has potential difference of 0.00V)

7. electrical double layer is a double layer between the charges and the electrode that will cause a potential difference

8. standard hydrogen electrode (must know how to draw the set up)

9. when potential difference is + then go to the right, when - then go to the left

10. salt bridge is a filter paper soaked in a saturated solution of potassium nitrate, it allows movement of ions

11. standard electrode potential for a half-cell is the voltage measured under standard conditions with a standard hydrogen electrode as the other half-cell

12. measuring standard electrode potentials (metal/metal half cell, non-metal/non-metal half cell, ion/ion half cell)

# data[].attributes.body
learnt that more positive an electrode potential the more ions are reduced, the more negative the more ions oxidised.

the standard electrode potential is
-1mol/dm^3
-25°C
-1 atm

salt bridge is a filter paper dipped in potassium nitrate for transferring ions

need to revise more on the parts of electrolysis and the steps for it including flow of electrons and ions

# data[].attributes.body
The more positive the voltage will be, the easier it is to reduce the ions on the left

Salt Bridge (a strip of filter paper soaked in potassium nitrate) is needed to be inserted in both beakers

Learned more about the three types of half cell: metal - metal ion half cell,

non metal - non metal ions and ion - ion half cell

# data[].attributes.body
From this lecture I revised about OILRIG which directly stands for oxidation is losing, reduction is gaining. This means that when you lose an electron, it is oxidation and the opposite would mean reduction. I also learnt about the ability of metals to lose electrons is directly linked to their reactivity. There are also 3 types of half cells that we learnt which are metal/metal, non-metal/non-metal and ion/ion. We also learnt about the standard hydrogen electrode.

# data[].attributes.body
Redox equilibrium = Harder to lose electron, equilibrium at right, easier to lose electron, equilibrium at left
Electrode potential = measurer volt to see how easy it is to reduce substance, more positive volt, easier to reduce, less positive volt harder to reduce. Later used to calculate volt difference between cell.
E standard (298K,1mol/dm3,1atm)
Salt bridge
Half cell
- metal/metal ion
- non metal/non metal ion
-ion/ion

# data[].attributes.body
What I learnt:
How electrolysis works (redox, electrode potential)
How an electrochemical cell works (negative ions go to anode, positive to cathodes, flow of electrons and ions, structure of different types of half cells, etc)
How to understand and calculate electrode potential and potential difference for cells.

# data[].attributes.body
OIL RIG:
Oxidation is loss of electrons
Reduction is gain of electrons

Electrode Potential:
More positive is reduced
More negative is oxidized

Salt bridge:
Filter paper dipped in potassium nitrate used to transfer ions

Half cells:
Metal/Metal ion; Non-metal/Non-metal ion; Ion/Ion

Conditions:
Concentration of ions is 1.00 mol dm-3
Temperature is 25 degrees Celcius / 298 Kelvin
Pressure of 1 atm

# data[].attributes.body
From the video I've learnt :
- electrode potentials tell us which element is reduced and which element is oxidised inthe redox reaction with electrochemical reactions, such that in electrolysis.
- the more positive the electrode potential, the more likely it's ions will be reduced
- the more negative the electrode potential, the more likely it's atoms will be oxidised
- electrode potentials can shift the position of the redox equilibrium

Electrochemical cells = modern day batteries

Electrode potentials using the standard hydrogen electrode potential

I've learnt how to draw
- metal/metal ion half cells
- non metal/metal ion half cells
- ion/ion half cellscell

# data[].attributes.body
Recalled:

OIL RIG (Oxidation is loss of electrons, reduction is gain of electrons.)

use standard conditions (298K, 1.00 mol/dm3, 1 atm)

more positive electrode potential - metal more unreactive (harder to be oxidised - oxidising agent); more negative electrode potential - metal more reactive (easily oxidised - reducing agent)

electrode potential cannot be measured directly, can only be compred

Cells:

salt bridges - made from porous materials, e.g. filter paper, soaked in KNO3 or NaNO3; connects circuits through separation of electrolytes

three half-cells:

metal-metal ion

metal-non-metal ion

solid

gas

ion-ion ion

This chapter is quite nice. I'll have to remember the half-equations better and revise content from chapter seven.

# data[].attributes.body
I learnt that more reactive metals have equilibrium lying closer to the left and unrective lies closer to the right. We can measure the the potential difference from the two systems but cannot measure the voltage directly.

Standard electrode potential
conditions
-1mol/dm^3
-25°C
-1 atm

If the voltage is more posititve than metals on the left are  poor reducing agent and if the voltage is negative, metals on the left are good reducing agents.

When reversing an equation, the voltage value is muliplied by -1

Positive terminal is the electrode that receives electrons.
Negative terinals is the electrode that loses electrons.

# data[].attributes.body
1. I revised about OIL RIG, (oxidation is loss of electron, reduction is gain of electron)

2. Equilibrium lies on the left for reactive metals and right for unreactive metals.

3. Faraday is electric charge carried by 2 mol of electron or charged ion

4. Electrode potential (for reduction):
- equilibrium goes right for more positive voltage
- equilibrium goes left for more negative voltage

5.  How to draw the half cell diagrams, electrochemical cells

6. Salt bridge is the movement of ions between 2 fall cells.

# data[].attributes.body
From this lecture, I learnt that

Metals are more reactive when they can easily lose electron, good reducing agent, bad oxidising agent.

Half equation must be in form of reduction.

For reactive metals, the equilibrium lies further to the left (lose electron, oxidised). For unreactive metals, the equilibrium lies further to the right (gain electron, reduced).

We cannot measure directly the electric potential of the metal but we can measure the difference in the potential between metal/metal ions system or other system.

Standard hydrogen electrode : E = 0.00V, 2H+ + 2e ⇌ 2H2.

Standard Electrode Potential : concentration of ions 1.00 mol dm^-3, 298K, gas must be in 101kPa, use Standard hydrogen electrode.

Electrolysis use battery, power supply, to move the ions. Electrochemical cell use the potential difference to move the electron and ion to move.

Salt bridge is for ions.

The E value is not affected in magnitude when an equation is being balanced, only the sign when switching sides.

1. Switch the equation to the LHS/RHS to cancel the electrons, this will change the sign of E value. 2. Add the E values together to get the total E value.

Silver nitrate cannot be used when ZnCl2 is used because the Ag+ will react with Cl- to form AgCl ppt that will block the salt bridge. Other nitrates are allowed.

3 types of half cell : metal/metal ion, non metal/non metal ion, ion/ion half cell.

Need to be careful for half cells like MnO4-/Mn2+. MnO4-(aq) + 8H+(aq) + 5e ⇌ Mn2+(aq)+ 4H2O(l). H+ ion need to be included in the solution.

Positive pole is the half equation that has the more positive/ less negative E value.

# data[].attributes.body
I learnt about how to find the electric potential of an electrochemical cell and learnt how it indicates whether the substance is easier to reduce or not. The more positive the voltage is, the easier it is to reduce the ion on the left while the more negative it is, the harder it is to reduce the ions on the left. I also pearnt that excluding the hydrogen half-cell, there are metal/metal ion half cell, non-metal/non-metal ion half cell and lastly ion/ion half cell. Lastly I learnt that the salt bridge is usually made of KNO3 in order to prevent inefficiency.

# data[].attributes.body
Learnt about
-OIL RIG
Oxi is loss
Reduc is gain

-Electrode Potential where
- more positive, more right (reduced)
- more negative, more left (oxidized)

Salt bridge
-Which can be made from a strip of filter paper (or other inert porous material) soaked in a saturated solution of potassium nitrate.and allows movement of ions between 2 half cells

Half cells
- Metal/Metal ion
- Non-metal/Non-metal ion
- Ion/Ion
Need to memorize and revise more on the half-cell diagrams ahd reactions

Lesson easy to understand

# data[].attributes.body
I learned about Redox reaction which consists of both oxidation as well as reduction.

Or in other words OILRIG, Oxidation Is Loss of electrons Reduction Is Gain of electrons.

I also learned about electrode potential. To determine the equilibrium of the equation, we have to look at the electrode potential whether it is positive or negative. The higher the value, the easier it is reduced. The electrons travel through a "Salt bridge" to move. Drawing the circuits will also be different for metal and metal ion, ion and ion half cell, as well as non-metal and non-metal half cell.

# data[].attributes.body
-Electron lost is oxidation while electron gain is reduction

-if electrode potential is positive, the equation moves more towards the right.

-if electrode potential is negative, the equation moves more towards the left

-Salt bridge is used to transfer the ions

-When calculating we do not multiply the V when balancing the equations.

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I learnt how to...

Electrode Potential ++ (half) Reactions:
- Always write in reduction form
- Always use standard conditions! (298K, pressure of 1 atm, concentration of 1 mol dm^-1)
- Differences in EP can only be compared, not measured directly
- The more positive it is, the more unreactive the metal is, harder to be oxidised——good oxidising agent.
- The more negative it is, the more receive it is, more easily oxidised—— good reducing agent.
- To put it simply, the electrode with the more positive EP will be the positive terminal.
- To put it simply, the electrode with the more negative EP will be the negative terminal, bc it loses electrons.

Electrochemical Cell:
- Does not need power supply
- Salt bridge (filter paper soaked in solution of KNO3 or NANO3): Separate ions and connects the circuit
*Electrolysis is different because it is a redox reaction

Voltage:
1. Get your reactions in reduction forms, compare their EPs
2. Rearrange the half-reaction with the more negative EP (just reverse it). Then also multiple the EP by -1
3. Add up the EP (no multiplying or dividing the EP)
4. balance the half-equation into one reaction

Three types of Half-cells:
- Metal/Metal ion
- Non Metal/Non Metal
- Ion/Ion

I feel like I need to memorise and revise more on the half-cell diagrams? As well as really understand why certain solutions cannot be used.

The lesson was great, there was also YT videos of the lectures for personal revisions as well.

# data[].attributes.body
- OIL RIG (oxidation: loss of electrons, reduction: gain of electrons)

- for reactive metals, equilibrium lies further to the left and they are good reducing agents

- for unreactive metals, equilibrium lies to the right and they are poor reducing agents

- we cannot measure directly the electric potential of the metal and metal ions in a solution

- electrolysis needs a power supply to move ions, electrochemical cell no need

- salt bridge separates the electrolytes and allows movement of ions

- salt bridge is filter paper soaked in saturated potassium nitrate solution

Measuring standard electrode potential

- metal/metal ion half cell

- non-metal/non-metal ion half cell

- ion/ion half cell

# data[].attributes.body
Redox Reaction
OIL RIG
oxidation - loss of electrons
reduction - gain of electrons

Electrode Potential

position of equilibrium differs for different combinations of metals placed in solutions of their ions

the electric potential of the metal and metals ions in solution cannot be measured directly, however, we can measure the difference in potential between the metal/metal ions system and another system (standard hydrogen electrode, E = 0.00V)

half-reaction is always written in reduction form

Standard Electrode Potential (E°)

The standard electrode potential for a half-cell is the voltage measured under standard conditions with a standard hydrogen electrode as the other half-cell.

Conditions:

Concentration of ions at 1.00 mol dm-3

Temperature of 25 °C (298K)

Pressure of 1 atm (101kPa)

Electrode Potential and Redox Reactions

Electrode potential values give us an indication of how easy it is to reduce a substance

if the E value is more positive, more right (reduced, good oxidizing agent)

if the E value is more negative, more left (oxidised, good reducing agent)

How to calculate the voltage

Use the table to get values in which afterwards, you have to switch signs and just add them. (voltage does not need to be balanced)

Salt bridge

can be made from a strip of filter paper (or other inert porous material) soaked in a saturated solution of potassium nitrate.

movement of ions between 2 half cells

Types of half cells (also learnt how to draw them):

metal/ metal ion

non-metal/ non-metal ion

ion/ ion

Overall, I learnt a lot from today's lecture and I think I need to revise more on drawing the half cells. Other than that, the teaching was clear and easy to understand as usual! :D

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In this lecture I learnt :
- More reactive metals are harder to reduce
- Less reactive metals are easier to reduce
- When calculating the standard cell potential, we don’t need to multiply the electrode potential to the numbers used to eliminate the electrons.
-Three types of half cells : 1. Metal/Metal 2. Non metal/ Non metal 3. Ion/ion
- The standard conditions of standard hydrogen electrode cells are : 1. Conc. Of ions at 1.00 mol/dm^3 2. Temperature at 298K 3. Gases at pressure of 1 atm 4. The value of the electrode potential of the half cell is measured relative to the standard hydrogen electrode.

The lecture made the chapter easier to understand. Very good teaching style 👍

# data[].attributes.body
Electrode potential

Position of equilibrium differs for different combination of metals placed in solutions of their ions

Cannot measure directly the electric potential of the metal and metal ions in solution

Can measure the difference in potential between metal/metal ions system by comparing two standard hydrogen electrode

What is an electrical double layer?           it is an interface between an electrode and an electrolyte that forms a layer of positive ions that is covered by negative ions due to attraction

Electrode potential and redox reaction

More positive, to the right

Most reactive metal can lose electrons easily(metal)

more negative, to the left

Salt bridge allows the movement of ions

Measuring standard electrode potential and how to draw

metal/metal ion half cell

Non-metals and non-metal ions

ion/ion cell half cell/

# data[].attributes.body
Electrolysis is a redox reaction. Oxidation is loss of electron. Reduction is gain of electron. Redox equilibrium is established when the rate of electron gain equals to the rate of electron loss. Voltage cannot be measured directly but the potential difference between metals can be measured. Salt bridge is used to connect the circuit and allowing ions to move. Salt bridge is usually potassium nitrate or sodium nitrate. Three type of half cells: metal/metal ion half cell, non metal/non metal ion half cell, ion/ion cell half cell.

# data[].attributes.body
More positive value for E means easier to be reduced.
Electrons move around at the external circuit, ions move around at salt bridge.
Voltage no need to multiply by any number when there is more mole of ions.
The way to draw different electrode.

# data[].attributes.body
Electrode potential

Position of equilibrium differs for different combination of metals placed in solutions of their ions

Cannot measure directly the electric potential of the metal and metal ions in solution

Can measure the difference in potential between metal/metal ions system and another system (Cu and Mg with salt bridge thingy)

Comparing with standard hydrogen electrode

What is an electrical double layer?

Due to electrical double layer, we cannot measure the absolute electrical potentials by the formation of the double layer.

Electrode potential and redox reaction

More positive, go to the right

Most reactive means can lose electron easily

Metal on the right is relatively unreactive and a relatively poor reducing agent

Positive : poor reducing agent (metals on the right)

Negative : good reducing agent (metals on the right)

Salt bridge can be made from a strip of filter paper (or other iner

Measuring standard electrode potential

metal/metal ion half cell

Non-metals and non-metal ions

ion/ion cell half cell/

# data[].attributes.body
1. Electrode potential, more + goes to the right, more - goes to the left. More + gains electrons, more - lose electrons
2. Ions move around salt bridge
3. To calculate the Voltage, no need to multiply by a number to balance the electrons.
4. 3 types of half cell : metal and metal ions, non metal and non metal ions, Ions and Ions

# data[].attributes.body
Electrical Potential
- Voltage more positive (+), shift to the right. Best oxidizing agent
- Voltage more negative (-), shift to the left. Best reducing agent

Salt Bridge
- Filter paper
- maintains ionic balance
- Movement of ions between 2 half cells
- Completes the circuit

Half cells
- Metal/Metal ion
- Non-metal/Non-metal ion
- Ion/Ion

# data[].attributes.body
I learnt about the conditions for reduction and oxidation, the half equation will always be reduction, and I also learnt about electrode potentials, and how to form equations for the anode and cathode.

# data[].attributes.body
i learnt that the more reactive, the position of the equilibrium will be shifted to the left. more positive = equation will go to the right. more negative= equation will go to the left. most negative means it goes through oxidation and is the best reducing agent. more negative means it loses electrons and is hardest to reduce.

# data[].attributes.body
I learnt that
- half equation is always written in reduction
- we use standard hydrogen electrode to compare the potential difference
- standard electrode potential: conc. of ions at 1moldm-³, temp. of 298K, gases at 1atm(101kPa)
- E value more positive = reaction will go right
- metals that gain electrons tend to be unreactive and a poor reducing agent
- ions are moved around salt bridge
- to count E value, add the respective E value from each reaction
- E value should not be multiplied
- salt bridge ions should not react with the ions from the substances
- positive terminal is anything that gains electrons
- non metal/non metal half cell
1. for gaseous atoms, platinum cell fully in contact with the gas
2. half of the platinum in contact with the solid and half with the aqueous solution.

I need to practice drawing the half cell and the diagrams of half cell for measuring between metal/metal ion, non metal/non metal ion and ion/ion cell.

# data[].attributes.body
During this lecture, I learnt about electrode potential, redox reactions, half cells, and salt bridge. We can use the mnemonic 'OIL RIG' to remember redox; oxidation is loss of electrons and reduction is gain of electrons. The cathode attracts cations (+) and the anode attracts anions (-). When electrode potential is positive, there is a greater tendency for the reaction to proceed to the right (reduction). Likewise, when electrode potential is negative, there is a greater tendency for the reaction to proceed to the left (oxidation). There are three types of half cells; metal/metal, non-metal/non-metal, and ion-ion. To complete the circuit, a salt bridge, where ions move between two half cells, is needed.

Need to revise more on: Drawing electrochemical diagrams.

Overall, the lecturer managed to explain the topic in a detailed manner, including clearing up doubts. The lecturer went through the topic at a fine pace, and the explanations, including presentations, notes, diagrams, and voice were easy to follow and understand. Lecture archives are also available for personal revision.

# data[].attributes.body
Always remember: OIL RIG

Equilibrium lies on:
- further to right for unreactive metals
- further to left for reactive metals

Electrode potential values -> how easy it is to gain electron
A trick (for equations in reduction form):
- more positive voltage, equilibrium goes right
- more negative voltage, equilibrium goes left / behind

Reducing agent reduces others but gets itself oxidised, while Oxidation agent oxidises others but gets itself reduced.

# data[].attributes.body
oxidation - loss of electron
reduction - gain of electron

Faraday - electriv charge carried by 1 mol of electron or charged ion

standard hydrogen electrode
one of the several types of half cell that can be used as reference electrode.

# data[].attributes.body
Redox
Reduction - gain of electrons
Oxidation - lose electrons

Electrode potential positive/ higher, easier to reduce
The ion with less negative/ positive electrode potential is oxidizing agent
Cathode attracts cations and anode attracts anions
Reduction occurs in cathode, oxidation occurs in anode

# data[].attributes.body
Redox
Oxidation -> loss of electrons
Reduction -> gain of electrons
(OIL RIG)

Electrode potential
- more positive, more right (reduced)
- more negative, more left (oxidised)

To calculate V generated :
Change the sign of E
No need to multiply if for example 2X+ needed

Salt bridge - movement of ions between 2 half cells

Types of half cell :
- metal metal
- nonmetal nonmetal
- ion ion

Need to revise more on :
Calculating electrode potential
Drawing diagrams

# data[].attributes.body
electrode potential (more positive move right, more negative move left)

electrochemical cell don’t need to use electricity

salt bridge soaked in solution of potassium nitrate

3 types of half cells

need to revise more on: drawing half cells

# data[].attributes.body
Learnt about :
- electrode potential ( more positive goes to right, more negative goes to left )
- salt bridge : functions and how to use
-  electrochemical cell Vs electrolysis : no need power supply while electrolysis needs power supply
- 3 types of half cells :
1. metal and metal ions
2. non metal and non metal ions
3. ion and ion
- electrode potential equation should not be multiplied ( voltage only )
- signs interchanged when equilibrium is to the left

need to revise more on :
drawing the electrochemical cells and half cell diagrams

teaching :
very detailed and easy to understand

# data[].attributes.body
Redox
oxidation is gain reduction is lose electron
(more positive is reduction)
(more negative is oxidation)

Standard hydrogen electrode diagram and use to determine electrode potential along with half cell diagram

Electrochemical cell with metal or non metal ion diagram

Salt bridge cannot use solution which can form precipitate

Electrode potential equation does not have to be multiplied or divided

# data[].attributes.body
Equation for potential difference needs to be in reduction form. The more unreactive one is the more positive one which is easier to reduce. The one that is more reactive is more negative which is harder to reduce. The positive one gains electrons and the negative one loses electrons.

To get voltage we need to use the table and switch the signs in the equation.

Salt bridge is the movement of ions between the two half cells.

3 types of half-cells
- Metal/Metal ion (Zn or Zn2+)
- Non Metal/Non Metal Ion (Cl2 or Cl-)
- Ion/Ion (Fe2+ or Fe3+)

This lesson was pretty good and I understood a lot. The teaching is great and it is very clear and thorough.

# data[].attributes.body
Today, i learnt about Electrode potential.

The voltage shows which metal is more reactive and unreactive. The more unreactive one is more positive and is easier to reduce. The more reactive metal is more negative and is harder to reduce. The one with the more positive voltage is the one that is gaining electrons while the more negative one is the one giving electrons. We can use this to determine which one is the positive side and negative side in the circuit.

Calculating the voltage is simple. Use the table to get the values and add them normally after switching signs. The voltage does not scale and does not need to be balanced along with the electrons.

I also learnt about the 3 types of half-cells and how to draw them.
- Metal/Metal ion
-Non-Metal/Non-Metal ion
-ion/ion
A salt bridge is needed to complete the circuit

I learnt alot from the lesson and i will need to revise on the detaisl and conditions for all these reactions to apply. The lesson was good and easy to understand.

# data[].attributes.body
- Electrode potential
1. Redox (OIL , RIG)
2. Unreactive metal : Further to the right - easier to reduce
3. Reactice metal : Further to the left - difficult to reduce
4. Half eq *in terms of reduction - electrons on the left*
5. Standard hydrogen electrode : E=0.00V
6. Need to know how to draw : Standard condition (298K)
*Platinum wire
*Glass bell
*Supply of hydrogen
*Platinum electrode
*H+ , 1.00 mol dm^-3
*Have hole (inside water)

- Electrode potential and redox reactions
1. Must be in reduction form (gain of e-)
2. V= Positive (go to right)- unreactive , V= Negative (go to left) - reactive

- Electrochemical cell
1. Do not need power supply
2. Salt bridge : Separate ions and connect the circuit
* Filter paper - soak in saturated solution of KNO3 or NaNO3
* AgNO3 -> Ag+ + NO3- , Zn2+ + Cl- , AgCl (participate)

- Measuring standard eletrode potentials
1. Metal / metal ions half cell
* loses electrons : Negative terminal
* Gain electrons : Positive terminal
2. Non- metal / non metal ion half cell
3. Ion/ ion cell half cell
* MnO4- (aq) + 8H+ (aq) + 5e- <=> Mn2+ (aq) + 4H2O (l)

- Using E⁰
1. Standard cell potential can be calculated for two half cells
2. Don't (times/multiply) voltage , only change sign for voltage

# data[].attributes.body
I learnt about electrode potential, salt bridge

The equation of potential difference is always in reduction form. The E value of any equation is the same even if the products are multiplied.

Salt bridge is the movement of ions between 2 fall cells.

# data[].attributes.body
During this lecture i learnt about electrode potential. I learnt that for potential difference, the equation must be in reduction form. The one that is more positive goes to the right and the more negative goes to the left.
The V is always the same even when halfed. (Cl2 and 1/2Cl2)

To calculate the voltage generated, reverse the more negative equation and change the voltage to positive (if negative) or vice versa.
Then add up the voltage.

I also learnt about salt bridge , which is the movement of ions between two half cells. (Prepare with potassium/sodium nitrate, so it wont form any precipitate)

Finally, i learnt further about the 3 types of half cells.
Metal/Metal ion (Zn/Zn2+)
Non metal/non metal ion (Cl2/Cl-)
Ion/ion (Fe3+/Fe2+)

# data[].attributes.body
1. Oxidation and Reduction can be remembered as OIL RIG (ox: loss of e, red: gain of e)

2. Redox Equilibrium: Rate of e gain = e loss

3. Reactive metals: Eq lies to left (ion)

4. Unreactive metals: Eq lies to right (metal)

5. EP cannot be measured directly, only the PD can be measured. The relative system is a hydrogen fuel cell.

6. Black platinum electrode is in contact with both H+ and H2.

7. Undergo under standard conditions (1 atm, 298K, 1 mol/dm3)

8. EP of HFC is 0

9. The hole in HFC is placed below the solution level to prevent H2 from escaping without reacting.

10. Eq must be in reduction form. When EP (+), the equation shifts to the right.

11. Reducing agent: Reduces others, but it itself will get oxidised.

12. Electrochemical cell has two metals with different reactivities so no power supply is necessary.

13. Salt bridge id a strip of filter paper that completes the equation, maintains ionic balance and allows ionic flow.

14. Electrons will only flow in the external circuit (the wires).

15. Electrolysis is a redox reaction.

16. When the equation is reversed, the EP symbol is switched.

17. When AgNO3 is used in the presence of Cl-, AgCl, a white precipitate will be formed and the reaction flow will be blocked.

18. EP can be measured with three setups M/M, NM/NM, I/I.

19. H2 cell releases e and Cu electrode accepts it. So H2 (-), Cu (+) terminal.

20. H2 cell loses e, Cl2 accepts so H2 (-), Cl2 (+)

21. I2 is a solid, Pt must be in contact with both I2 and I-.

22. Must memorise or be able to derive MnO4- (aq) + 8H+ (aq) + 5e <—> Mn2+ (aq) + 4H2O, but can be found in Data Booklet.

23. Two beakers: MnO4-, Mn2+, H+ in B1 and H+ in B2.

24. EP (Ag+ red) = EP (2Ag+ red)

25. Those who lose e- will become the (-) electrode.