1.8 - Electrochemistry

Why does Zinc become aqueous in the following reaction:

• Zn(s) + CuSO₄(aq) → Cu(s) + ZnSO₄(aq)

Zinc becomes an aqueous ion more easily as less energy is needed for atomisation + ionisation

What dictates if a Metals ability to displace another metal in a reaction?

The metal that displaces the other is a greater reducing agent, but weaker oxidising agent

How do you know if a Metal will reduce or oxidise another based off the Electrochemical series?

Based on the “order” of reduction/oxidising power:

• Rule of thumb = Anything below a reducing/oxidising agent will be affected in the same way

What is a Half Cell?

One half of an electrochemical cell, formed from a metal electrode dipped in its ions, or a platinum electrode with 2 aqueous ions

When is a Platinum electrode used? (2)

1. When a solution has 2 or more aqueous ions

2. When the element may be in gaseous form

Why use a Platinum electrode? (2)

1. Chemically inert

2. Electrically conductive

Why do we separate Half Cells?

So that the less reducing metal half cell doesn’t gain electrons + ensure electron flow to be usable → electricity 

What is an Electrochemical cell?

Cell made of 2 half cells joined by a wire, voltmeter + salt bridge

What is a Voltmeter used for?

Measures PD between 2 half cells

How do electrons flow in an Electrochemical cell?

Electrons flow from a more reactive metal to a less reactive metal

Where are the +ve and -ve electrodes placed in an Electrochemical cell?

• Why

+ve = Right

-ve = Left

• Since electrons flow from negative to positive charges

Where does equillibrium lie to in:

• +ve elecrode

• -ve electrode

• Positive = More reducing → Equilibrium lies left (more ions)

• Negative = More oxidising → Equilibrium lies right (more atoms)

What is the Acronym for identifying which electrode is Oxidising + Reducing?

NOR PRoblem:

• N.O.R = Negative → Oxidises → (Reverse when combining half cell equations)

• P.R = Positive → Reduces
  

Describe the distribution of charge or electrons in an Electrochemical cell

• Negative = Since the (more) reducing gains electron charge builds up on its electrode. We connect a wire to the more oxidising electrode + electrons flow from the negative electrode to the other. The number of electrons in the reducing agent decreases to maintain equilibrium.

• Positive = The more oxidising electrode (more negative electrode) atoms revert to aqueous ions. The number of electrons increase for the more reducing electrode (more positive electrode), therefore to oppose the change, equilibrium shifts right so more ions in solution convert to atoms.

What is the EMF of an Electrochemical cell?

E_cell = E_reduced - E_oxidised

What is the following answers to the Electrochemical series: (why)

• Most reducing agent

• Most oxidising agent

• Most reducing = F₂ (most electronegative)

• Most oxidising = Sr (least electronegative)

What is the Standard Hydrogen Electrode?

• Why use Hydrogen gas? (2)

• Why use that specific electrode? (2)

Where gaseous Hydrogen is bubbled into a half cell containing 1 moldm^-3 of H+ ions in solution, a platinum electrode = 0.00V

• Hydrogen is abundant + cheap

• Inert + Conducts electricity

What are the 3 Standard Conditions of S.H.E?

• Diprotic acid

1. 298K

2. 100kPa

3. 1 moldm^-3 of H+ ions -> 1 moldm^-3 of HCl or 0.5 moldm^-3 of H₂SO₄ (since its diprotic)

How does equilibrium shift change EMF?

Same Le Chateliers principle changes:

• Makes EMF more negative = More electrons

• Makes EMF more positive = Less electrons 

Cells

Cells

Give the following:

• Electrode polarity

• What occurs as the electrode

• What type of agent is it

• Negative electrode → Undergoes Oxidation → Becomes a Reducing agent

• Positive electrode → Undergoes Reduction → Becomes an Oxidising agent

Explain how dynamic equilibrium in a Cell is met

• Include mention of movement of electrons 

• Solution charges

• The more reducing agent (negative electrode) loses electrons as electrons flow between electrodes via a wire, decreasing the number of electrons causing an equilibrium shift to the left to increase the number of +ve ions → solution of the negative electrode is more +ve 

• The more oxidising agent (positive electrode) gains electrons from the negative electrode via a wire, increasing the number of electrons and therfore reacting with ions to form more atoms → decreasing the concentration of +ve ions decreases the charge of the solution making the solution more -ve

What causes an EMF to drop?

Lack of ions in solution prevent the flow of electrons, causing charge build up → dropping PD

How do we prevent the build up of charge (decrease PD)

• How does it work?

Adding anions (-ve ions) to the +ve solution (the negative electrodes solution) and cations (+ve ions) to the -ve solution (the positive electrodes solution) to allow the flow of ions to prevent charge build up

• Prevention = Salt bridge

What is a Salt Bridge?

Filter paper soaked in saturated ionic solution that contains (cat/an)ions, to allow for ion flow to balance charges

What are 3 causes that lead to PD drop?

1. All of an electrode has been turned to aqueous ionised

2. All of aqueous ions form atoms on the electrode surface

3. When the ions of a salt bridge are depleted

What is a commonly used solution for Salt bridges?

• Why (2)

KNO₃: 

1. It contains (cat/an)ions

2. Ions dont react with metal ions in their solutions

What 4 ions aren’t used for Salt Bridges?

• Why

1. Halides

2. Hydroxides (OH^-)

3. S^2-

4. CO₃^-

• All form precipitates which coat an electrode → slowing the electron transfer + drops EMF

What is the Layout for Cell Notation

• Layout

• Symbol meanings

• Order

• Rule (use example if needed)

• Reduced Form | Oxidised Form || Oxidised Form | Reduced Form

• " | " = State difference

• " || " = Salt bridge

• " , " = Separating DIFFERENT ions in the SAME

• Oxidised form = Ions -> Higher charge = Closer to centre/Salt bridge

When do we know if EMF is feasible?

+ve PD value

Non-Rechargeable cells + Commercial applications

Non-Rechargeable cells + Commercial applications

What is the rule when combining half cell equations?

When combining equations, use LCM to get the same number of electrons, then cancelling them out

***USE CHEM NOTES 1 + BLURT 2 FOR DIAGRAMS AND OTHER BATTERY TYPES***

***USE CHEM NOTES 1 + BLURT 2 FOR DIAGRAMS AND OTHER BATTERY TYPES***

What is a Separator?

A porous material that separates +ve and -ve solutions to prevent an instantaneous redox → causing an explosion, allowing for the flow of ions between solutions to prevent charge build up

What is a Dry cell?

• Advantage

A cell where reactants have moisture, but not entirely, to allow the flow of ions but decrease mass and size:

• Prevents loss of reactants via spillage

What is an Electrolyte?

• 3 examples

Solution or paste with free moving ions allowing flow of electron in a circuit:

• ZnSO₄

• Fe₂SO₄

• MgSO₄

How do we recharge a cell?

• 2 reasons it doesnt work with non-rechargeable cells

Apply strong electrical force to reverse electron motion + reverse to allow flow of electrons to allow for power output:

• Doesn’t work for Non-rechargeable cells:

1. Forms irreversible products

2. Reversible reactions producing gaseous products → increasing Pa causing explosions

Give an advantage of rechargeable batteries:

• Why

Dont have to be thrown away/recycled:

• Prevents harmful electrolytes + metals decomposing into the atmosphere

Describe a Lead - Acid Battery

• Electrodes

• Electrolytes

• -ve electrode = Pb 

• +ve electrode = PbO₂ → oxygen is lost to form water molecules

• Electrolyte = H₂SO₄ → Hydrogen ions forms water with oxygen from +ve electrode

Give the half equation that occurs at the Cathode (+ve)

• What type of reaction is it?

• What happens when the electrode is “used up”

• PbO₂(s) + 4H⁺ + 4e^- → Pb(s) + 2H₂O(l)

• Reduction

• Lead reacts to form PbSO₄ like the anode does

Give the half equation that occurs at the Anode (-ve)

• What type of reaction is it?

• Pb(s) + HSO₄^- → PbSO₄(s) + H⁺ + 2e^-

• Oxidation

Give the net overall ionic equation for a Lead-Acid battery

Pb(s) + PbO₂(s) + 4H⁺ + SO₄^2-(aq) → 2PbSO₄(s) + 2H₂O(l)

What happens to a Lead-Acid battery when used up?

Electrodes are all coated in Lead Sulfate → since H₂SO₄ forms ions which react with the electrodes

How do electrons flow in Lead-Acid battery?

The anode (-ve) forms positive lead ion (2+), increasing the number of electrons so they flow to the cathode to form lead atoms

Describe the diagram of a Lithium-Ion battery

• Electrodes

• Electrolytes

• Cathode (+ve) = CoO₂

• Anode (-ve) = Lithium with graphite powder

• Electrolyte = LiPF₆

What is the role of Graphite in the Lithium-Ion batteries?

• Why?

Graphite powder acts as a support medium:

• Allows dissolved Li ions to travel through the electrolyte to the cathode to form LiCoO₂

Explain the movement of Li ions to and from electrodes (usage + charging)

Li ions absorbed by graphite allowing them to move easily to the cathode → Li ions inserted in the Cobalt layers to form LiCoO₂

When charging they return to the cathode and forms their ions of Li⁺ and CoO₂^-, due to the structure of LiCoO₂ allows the ions to move through the electrolyte

Give the half equations of a Lithium Ion battery:

• Anode

• Cathode

• Anode (-ve) = Li → Li⁺ + e^-

• Cathode (+ve) = CoO₂ + e^- → CoO₂^-

What is the state of the electrolyte? (why)

Solid:

• Since Li⁺ and H₂O is an explosive reaction in solution/moisture

Give the 4 advantages of Lithium Ion batteries

1. No water which makes it lighter + prevents rust damage

2. Per atom Pb produces 2 times the amount of electrons as what is used → makes Li ion battery more efficient

3. Lithium is a light element

4. Lithium is a very strong reducing agent → produces higher PD

What is a Fuel Cell?

Battery which is continuously fed reactants so that it supplies PD without being used up over time

Describe the Hydrogen Fuel cell

• Reactants

• Electrolytes

• Hydrogen gas in (Anode -ve)

• Oxygen gas in (Cathode +ve)

• Water leaves

• Electrolyte = KOH/NaOH

Describe the process that Hydrogen goes through the Hydrogen Fuel cell (Alkaline)

• Equation

• Which electrode

Hydrogen is oxidised in the anode to form H⁺ which enter the electrolyte to react with OH to form water and electrons (which go to the cathode):

• H₂ + 2OH^- → 2H₂O + 2e^-

• Anode

Describe the process that Oxygen goes through the Hydrogen Fuel cell (Alkaline)

• Equation

• Which electrode

Oxygen reduces to Oxide ions (O^2-) reacts with H₂O to form the OH^- ions for the electrolyte:

• O₂ + 2H₂O + 4e^- → 4OH^-

• Cathode

Give the overall equation for an Alkaline Hydrogen fuel cell

• 2H₂ + 2O₂ → 2H₂O

What are the 2 equations that occur at a Hydrogen fuel cells?

• Anode (Oxidation):
  H₂ → 2H⁺ + 2e^-

• Cathode (Reduction):
  O₂ + 4H⁺ + 4e^- → 2H₂O

What is the electrolyte in an Acidic Hydrogen Fuel cell?

H₃PO₄

Give the summary for equations at the Anodes + Cathodes of:

• Acidic fuel cell

• Alkaline fuel cell

• Acidic Fuel Cell:
  > Anode (Oxidation)  - - -  H₂ → 2H⁺ + 2e^-
  > Cathode (Reduction)  - - - O₂ + 4H⁺ + 4e^- → 2H₂O
  
  Anode makes Ions / Cathode makes water
  

• Alkaline Fuel Cell:
  > Anode (Oxidation)  - - -  H₂ + 2OH^- → 2H₂O + 2e^-
  > Cathode (Reduction) - - -  O₂ + 2H₂O + 4e^- → 4OH^-
  
  Anode makes water/ Cathode makes ions
  
  Both Anodes involve Hydrogens, Cathodes involve Oxygen

Give the 3 advantages of Hydrogen fuel cells

1. Only byproduct is water

2. Refuelling is quicker than recharging

3. Recharging involves electrolysis means no carbon emissions

Give the 3 disadvantages of Hydrogen fuel cells

1. Production of cells use machines + electricity which release emissions

2. Space consuming

3. H₂ is flammable so risky to be in heat

Rank the efficiency of:

• Battery

• Petrol engines

• Hydrogen Fuel cell

Battery > Hydrogen Fuel cell > Petrol engine

How can we increase the efficiency of petrol engine efficiency?

• 2 advantages

Hydrogen is combusted in petrol engines, increasing efficiency:

• Water byproduct

• Increased efficiency → better than normal petrol (ONLY)