Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry - A-level > Physical: Electrode Potential and Electrochemical Cells > Flashcards

Physical: Electrode Potential and Electrochemical Cells Flashcards

1
Q

What is the basis of an electrochemical cell

A

When a rod of metal is dipped/placed in a solution of its own metal ions an equilibrium is set up
There is a tendency for the metal to form +ve ions in the solution
Or gain electrons and form a metal atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What happens if the equilibrium lies to the left

A

Metal acquires -ve charge due to the build up of e- on the metal
The electrode has a -ve charge

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What happens if the equilibrium lies to the right

A

+ve charge builds up in the metal as the electrons have been used up to form the metal ions
Electrode has a +ve potential

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What determines wither a metal wants to become an atom or ion

A

Reactivity and position of equilibrium

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Do more reactive metals tend to form atoms or ions

A

Ions - eq shifts left as easier to lose an electron

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Do less reactive metals tend to form atoms or ions

A

Metals - eq shift right easier to gain

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What is a half-cell / electrode

A

Metal in a solution of its ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What are the 3 types of electrode

A

Metal - metal surrounded by solution of its ions
Gas- gas and solution of ions - inert metal is actual electrode to allow flow of electrons
Redox-two ions of same element with inert metal electrode

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Give an example of a half-cell

A

For example:
Zn (s) + CuSO4 (aq)→ Cu (s) + ZnSO4 (aq)

Instead of electrons being transferred directly from the zinc to the copper ions, a cell is built which separates the two redox processes
For example:
Zn (s) ⇌ Zn2+ (aq) + 2e–

If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up
Each part of the cell is called a half cell

The Zn atoms on the rod can deposit two electrons on the rod and move into solution as Zn2+ ions:
Zn(s) ⇌ Zn2+(aq) + 2e–

This process would result in an accumulation of negative charge on the zinc rod
Alternatively, the Zn2+ ions in solution could accept two electrons from the rod and move onto the rod to become Zn atoms:
Zn2+(aq) + 2e– ⇌ Zn(s)

This process would result in an accumulation of positive charge on the zinc rod

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

How do u measure the potential of a half cell

A

Measure the potential difference between the 2 half cells

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What factors does the p.d depend on

A

the nature of the ions in solution
the concentration of the ions in solution
the type of electrode used
the temperature

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What is the electrode potential

A

Show how easily a substance is reduced
Half equations must be wrote with the electrons on the left

The more positive (or less negative) an electrode potential, the more likely it is for that species to undergo reduction
The equilibrium position lies more to the right

The more negative (or less positive) the electrode potential, the less likely it is that reduction of that species will occur
The equilibrium position lies more to the left

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What is the standard hydrogen electrode

A

Potential that is used to measure against the potentials produced so that they can be measured
This is assigned to potential of O
And is known as the primary standard

The standard hydrogen electrode is a half-cell used as a reference electrode and consists of:
Hydrogen gas in equilibrium with H+ ions of concentration 1.00 mol dm-3 (at 100 kPa)
2H+ (aq) + 2e- ⇌ H2 (g)

An inert platinum electrode that is in contact with the hydrogen gas and H+ ions
When the standard hydrogen electrode is connected to another half-cell, the standard electrode potential of that half-cell can be read off a high resistance voltmeter

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is an electrochemical cell

A

Two half cells
- two metals joined by a wire
-solutions joined with a salt bridge - ions flow through and go into each solution
Voltmeter used to collect he p.d

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is the salt bridge

A

-either
-a piece of filter paper soaked in a solution if unreactive irons
-tube containing unreceptive ions in an Agar gel
-compound s such as KNO3 are used as ions re unreactive

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What are the standard conditions of an electrode potentials

A

Ion concentration of 1.00 mol dm-3
A temperature of 298 K
A pressure of 100 kPa
It is written as E dot
Needs to be standard conditions as a change in conditions could effect the equilibrium

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

How is the EMF calculated

A

standard cell potential (Ecellꝋ) can be calculated by subtracting the less positive Eꝋ from the more positive Eꝋ value
The half-cell with the more positive Eꝋ value will be the positive pole
By convention this is shown on the right hand side in a conventional cell diagram, so is termed Erightꝋ
The half-cell with the less positive Eꝋ value will be the negative pole
By convention this is shown on the left hand side in a conventional cell diagram, so is termed Eleftꝋ
Ecellꝋ = Erightꝋ - Eleftꝋ

Since oxidation is always on the left and reduction on the right, you can also use this version
Ecellꝋ = Ereductionꝋ - Eoxidation

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Example of calculating the EMF/ standard potential of a cell

Calculate the standard cell potential for the electrochemical cell below and explain why the Cu2+ / Cu half-cell is the positive pole. The half-equations are as follows:

Cu2+(aq) + 2e- ⇌ Cu(s) Eꝋ = +0.34 V

Zn2+(aq) + 2e- ⇌ Zn(s) Eꝋ = −0.76 V

A

Step 1: Calculate the standard cell potential. The copper is more positive so must be the right hand side.

Ecellꝋ = Erightꝋ - Eleftꝋ

Ecellꝋ = (+0.34) - (-0.76)

= +1.10 V

The voltmeter will therefore give a value of +1.10 V

Step 2: Determine the positive and negative poles

The Cu2+ / Cu half-cell is the positive pole as its Eꝋ is more positive than the Eꝋ value of the Zn2+ / Zn half-cell

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What is the conventional notation of electrochemical cells

A

A solid vertical (or slanted) line shows a phase boundary, that is an interface between a solid and a solution
A double vertical line (sometimes shown as dashed vertical lines) represents a salt bridge

The substance with the highest oxidation state in each half cell is drawn next to the salt bridge
The cell potential difference is shown with the polarity of the right hand electrode

The cell convention for the zinc and copper cell would be
Zn (s)∣Zn2+ (aq) ∥Cu2+ (aq)∣Cu (s) E cell = +1.10 V

This tells us the copper half cell is more positive than the zinc half cell, so that electrons would flow from the zinc to the copper
The same cell can be written as:
Cu (s)∣Cu2+ (aq) ∥Zn2+ (aq)∣Zn (s) E cell = -1.10 V

The polarity of the right hand half cell is negative, so we can still tell that electrons flow from the zinc to the copper half cell

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What are secondary standards

A

The standard hydrogen electrode potential is difficult to use so other standards are used to calibrate against the SHE

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

In electrochemical cells which one is the anode and which is the cathode

A

Anode - metal that is oxidised
Cathode - metal that is reduced

22
Q

In electrolytic cells which one is the anode and the cathode

A

Cathode = -ve
Anode = +ve

23
Q

what are hydrogen-oxygen fuel cells

A

The fuel cell consists of
-a reaction chamber with separate inlets for hydrogen and oxygen gas
-an outlet for the product - water
-an electrolyte of aqueous sodium hydroxide
-a semi-permeable membrane that separates the hydrogen and oxygen gases

  1. Hydrogen enters at the negative electrode and releases electrons. Hydrogen is oxidised by hydroxide ions to water
  2. The electrons flow through the external circuit. This is where they can do work and transfer energy. The electrons flow to the positive electrode where oxygen enters.
  3. The electrons are accepted and releases hydroxide ions. Electrons are accepted by the oxygen entering at the positive electrode
  4. The hydroxide ions travel to the negative electrode. The hydroxide ions travel through the semi permeable membrane to the negative electrode where they react with the hydrogen atoms and oxidation occurs. The cycle repeats itself.
24
Q

what are fuel cells

A

A fuel cell is an electrochemical cell in which a fuel donates electrons at one electrode and oxygen gains electrons at the other electrode

These cells are becoming more common in the automotive industry to replace petrol or diesel engines

As the fuel enters the cell it becomes oxidised which sets up a potential difference or voltage within the cell

Different electrolytes and fuels can be used to set up different types of fuel cells

25
Q

what is the oxidation equation for a hydrogen-oxygen fuel cell

A

2H2 (g) + 4OH– (aq) → 4H2O (l) + 4e– Eꝋ = -0.83 V

26
Q

what is the reduction equation for the hydrogen-oxygen fuel cell

A

O2 (g) + 2H2O + 4e– → 4OH– (aq) Eꝋ = +0.40 V

27
Q

what is the overall equation of the hydrogen-oxygen cell

A

2H2 (g) + O2 (g) → 2H2O (l)
Eꝋ = +1.23 V

28
Q

what are the benefits of a hydrogen-oxygen cell

A

Water is the only reaction product, so fuel cells present obvious environmental advantages over other types of cells

The reaction is the same as hydrogen combusting in oxygen, but since the reaction takes place at room temperature without combustion, all the bond energy is converted into electrical energy instead of heat and light

There are no harmful oxides of nitrogen produced, which are usually formed in high temperature combustion reactions where air is present

29
Q

what are the risks of a hydrogen-oxygen cell

A

Hydrogen is a highly flammable gas and the production and storage of hydrogen carries safety hazards

Very thick walled cylinders and pipes are needed to store hydrogen which has economic impacts

The production of hydrogen is a by-product of the crude oil industry, which means it relies on a non-renewable, finite resource

Until a cheap way is found to make hydrogen, its widespread use in fuel cells will be limited

Hydrogen has high energy density, that is, the amount of energy contained in 1g of the fuel is high compared to other fuels, but because it is a gas, its energy density per unit volume is low which means larger containers are needed compared to liquid fuels

30
Q

what is the history of a hydrogen-oxygen fuel cell and what are they used for

A

-invented in late 1930s
-used by NASA in Apollo missions from 1968 and 1972
-used to produce potable water, heat and electricity
-Fuel cells have been used on space craft, where the product can be used as drinking water for astronauts

31
Q

what are lithium cells

A

-used to power laptops and mobile devices
-Lithium is used because it has a very low density and relatively high electrode potential

The cell consists of:
-a positive lithium cobalt oxide electrode
-a negative carbon electrode
-a porous polymer membrane electrolyte - this cannot leak

The cell consists of a sandwich of different layers of lithium cobalt oxide and carbon

When the cell is charged and discharged the lithium ions flow between the negative and the positive through the solid electrolyte

32
Q

what is the oxidation reaction for a lithium cell

A

Li (s) → Li+ (s) + e– Eꝋ = -3 V

33
Q

what is the reduction equation for a lithium cell

A

Li+ (s) + CoO2 (s) + e– → Li + (CoO2)(s)
Eꝋ = +1 V

34
Q

what is the overall equation for a lithium celll

A

Li (s) + CoO2 (s) → Li + (CoO2) – (s) Eꝋcell ~ +3.5

35
Q

what are some of the problems of lithium cells

A

A global shortage of lithium is likely to make lithium ion cells unsustainable as the current demand for lithium exceeds the supply

If cells are not recycled but thrown away in landfills, then a huge amount of lithium becomes lost to future generations

Reports of lithium ion cell fires have raised concern about the safety of these batteries in electronic devices; it is a reminder to us that lithium is a very reactive element in Group 1 of the periodic table, which is why it has a high electrode potential

36
Q

what is the history and uses of a lithium cell

A

-invented in 1970s but fist used in 1991
-rechargable battery

37
Q

what are lead-acid batteries

A

Lead-acid batteries consist of six cells joined together in series

The cells use lead metal as the negative electrode and and lead(IV) oxide as the positive electrode

The electrolyte is sulfuric acid

In a commercial car battery, the six cells in series give a combined voltage of about 12 V

When the car is in motion, the generator provides a push of electrons that reverses the reaction and regenerates lead and lead(IV) oxide

38
Q

what is the oxidation equation for a lead-acid battery

A

Pb (s) + SO42- (aq) → PbSO4 (s) + 2e- Eꝋ = -0.36 V

39
Q

what is the reduction equation for a lead-acid battery

A

PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e- → PbSO4 (s) + 2H2O (l)
Eꝋ = +1.70 V

40
Q

what is the overal equation for a lead-acid battery

A

PbO2 (s) + 4H+ (aq) + 2SO42- (aq) + Pb (s) → 2PbSO4 (s) + 2H2O (l)
Eꝋcell = +2.06 V

41
Q

what is the history and use of lead acid batteries

A

-first type of rechargable battery - 1859
-Lead-acid batteries are designed to produce a high current for a short period of time, hence their use in powering a starter motor in car engines

42
Q

problems of lead aicd batteries

A

-They are very heavy
-They contain toxic materials: lead and lead(IV) oxide
-The sulfuric acid electrolyte is very corrosive

This presents challenges of disposal when lead-acid batteries come to the end of their useful life

43
Q

what are zinc-carbon cells

A

a zinc casing which acts as the negative electrode

a paste of ammonium chloride which acts as an electrolyte as well as the positive electrode

a carbon rod which acts as an electron carrier in the cell

44
Q

what is the oxidation equation of a zinc-carbon cell

A

Zn (s) → Zn2+ (aq) + 2e-
Eꝋ = -0.76 V

45
Q

what is the reduction equation for the zinc-carbon cell

A

2NH4+ (aq) + 2e- → 2NH3 (g) + H2 (g) Eꝋ = +0.74 V

46
Q

overall equation for a zinc-carbon cell

A

2NH4+ (aq) + Zn (s) → 2NH3 (g) + H2 (g) + Zn2+ (aq) Eꝋcell = +1.50 V

47
Q

use and history of a zinc-carbon cell

A

-used in low drain applications - remote controls, clocks and torches
-invented in 1886 but first used in 1920

48
Q

benefits of zinc-carbon cell

A

The cell provides a small current and is relative cheap compared to other cells

Extra long life cells have a similar chemistry, but supply a higher current

49
Q

negatives of zinc-carbon cells

A

-As the cell discharges, the zinc casing eventually wears away and the corrosive contents of the electrolyte paste can leak out, which is an obvious disadvantage of zinc-carbon cells

50
Q

which cells are non-rechargable

A

-zinc-carbon
-hydrogen-oxygen

51
Q

what are the rechargable cells

A

-lead
-lithium

pass charge of current through the battery to regenerate the current/electron flow

Chemistry - A-level (15 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions