Chemistry Unit 3 Topic 2 (Redox Reactions)

Redox reactions

transfer of electrons

What is more likely to be a reducing agent

• fewer valence electrons

• low ionization energy

• low electronegativity

• Metals

What is more likely to be an oxidizing agent

• more valence electrons

• high ionization energy

• high electronegativity

• non-metals

more reactive a metal is, the more likely it is to

oxidise

Identify metals with set oxidation state and their exception

• fluorine = -1 (most electronegative)

• hydrogen = +1 (except metal hydrides…metal + hydrogen)

• oxygen = -2 (h2O2)

element oxidation state by itself

0

oxidation sates in monatomic ions

equals to their charge

sum of oxidation states in polyatomic ions

=polyatomic ion charge

the most electronegative atom in a molecule, their oxidation state is always

negative

when molecule/atom is oxidized, oxidation state

increases

how to determine oxidation number from transition metal

represented by their pronumeral

when molecule/atom is reduced, oxidation state

Decreases

when molecule/atom is reduced, oxidation state

increases

Type of reactions that can be redox

how to balance redox reaction under acidic conditions

1. oxidation number for every atom

2. determine which is the reducing and oxidising agent

3. Split into unbalanced half equations

• balance all atoms through coefficient except O and H first

• Balance O by adding H2O to one side

• balance H by adding it to H+ ion

• balance charge by adding electrons to the side with more positive charge

1. combining electrons

• ensure electron lost through oxidation = electron gained in reduction which can be down by multiplying to a common factor

• simplify by cancelling out like terms on the reactant and product side

Galvanic/voltaic cell is 

spontanous: a reaction that can occur without external energy input

anode in galvanic/voltaic cell is

negative

anode in galvanic/voltaic cell is

positive

electrolytic cell is 

non-spontaneous: requires a continuous supple of external energy to proceed

anode in electrolyic cell is

positive

cathode in electrolyic cell is

negative

similarities between galvanic and electrolytic

• oxidation occurs at anode

• reduction occurs at cathode

• e flow from anode to cathode

• anions more towards anode

• cations move towards cathode

• anode and cathode are electrodes

a more positive standard electrode potential means a

stronger oxidizing agent(left) and weaker reducing agent (right)

how to represent galvanic cell in cell diagram

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

• Left side = anode (oxidation)

• Right side = cathode (reduction)

• “|” separates phases, “||” = salt bridge

what is salt bridge in galvanic cell

A tube filled with a neutral ionic solution used to balance the charge and keep the reactions going. 

If they are multiple ions which types of ions gets oxidized at the anode 

the stronger reducing agent

If they are multiple ions which types of ions gets reduced at the cathode 

the stronger oxidizing agent

If the difference in E is small for oxidizing agents, how do you determine which one gets oxidized at the anode

The one with the higher concentration

standard Hydrogen Electrode

The Standard Hydrogen Electrode (SHE) is the reference electrode against which all other electrode potentials are measured.

E is the tendency of the reactant to be reduced, with a more positive indicating its more likely to be reduced?

Identify the limitations associated with standard electrode (reduction) potentials, E°.

standard electrode (reduction) potential, E°

The standard electrode potential (E°) is the voltage measured when a half-cell is connected to the standard hydrogen electrode, under standard conditions:

Calculate cell potential, E°cell

E°cell​=E°reduction(cathode)​−E°oxidation(anode)​

Predicting Spontaneity using E° of cell

• If E°cell > 0, the reaction is spontaneous.

• If E°cell < 0, the reaction is non-spontaneous.

Describe electrolysis plating

• impure metal attached to anode.

• A stronger reducing agent than the metal dissolve at the anode, but dont reduce at cathode

• Less reactive metals than metal dont oxidise so they fall off the electrode.

• THIS IS HOW METAL PLATING WORK BECAUSE EVEN THOUGH STRONGER REDUCING IMPURITY WILL OXIDISE AT ANODE FIRST IT WILL NOT BE REDUCED BECAUSE IT WILL BE A WEAKER OXIDISING AGENT.

• Strong reducing agent ↔ weak oxidising agent (and vice versa).

• Metal plating relies on this:

  • More reactive metals (strong reducers like Zn, Ni) dissolve at the anode but do not plate out.

  • Less reactive metals (weak reducers, strong oxidisers like Cu²⁺, Ag⁺) are the ones that get reduced and deposit as pure metal.

Hydrogen Fuel Cells

a type of electrochemical cell which combines hydrogen and oxygen to form water

In Acidic Hydrogen fuel cells what is oxidized at anode and reduced at cathode

• Hydrogen at anode

• Oxygen at cathode

What happens at the cathode in Acidic Hydrogen fuel cell

oxygen reacts with positive hydrogen that came from the anode through the center and electron producing water

What is the electrolyte in acidic hydrogen fuel cell (the ions that moves through the center)

Positive hydrogen aka proton

In Alkaline Hydrogen fuel cells what species react at the anode

In Alkaline Hydrogen fuel cells what species react at the anode

What is the electrolyte in alkaline hydrogen fuel cell (the ions that moves through the center)

Hydroxide ions

In electrolysis plating is the impure metal attached to anode or cathode

anode