Electrochemistry Unit

Faraday’s Law

amount of substance formed at an electrode during electrolysis is proportional to time and electricity

electrical charge formula

• q = IΔt

• q = electrical charge

• I = electrical current

• Δt = change in time (seconds)

Faraday’s constant

9.65 x 10⁴ C/mol e-

Faraday’s stoichiometry

(molar mass(g) / 1 mol) x (1 mol / mol of electron) x (1 mol of electron / 9.65 × 10⁴ c) x (couloumbs / 1 s) x (s / 1)

REDOX stoichiometry

(mol/1L) x (L/1) x (mol ratio/mol ratio) x (mol ratio/mol ratio) x (1/L)

voltaic cell

• generate electricity

• spontaneous REDOX

• electrolyte, two electrodes, salt bridge, external circut

electrolytic cell

• consume electricity (forced reaction)

• non-spontaneous REDOX

• purify/extract metals and produce non-metals

• electrolyte, two electrodes, power supply

anode

• lose electrons to become ions

• detach from electrode

• voltaic - negative

• electrolytic - positive

• oxidation

cathode

• gain electrons to become atoms

• attach to electrode

• voltaic - positive

• electrolytic - negative

• reduction

cathode half reaction

reduction

anode half reaction

oxidation

shorthand cell notation

• Anode  |   Half-Cell Species  ||   Half-Cell Species  |   Cathode

• solid line - different state

• double line - salt bridge

Brine Solution

• NaCl solution

• H₂O is the SOA

• Cl^- is the SRA

cation

• positively charged ion

• move towards cathode

anion

• negatively charged ion

• move towards anode

standard electrode potential

shows how easily a substance is reduced

positive standard electrode potential

• more easily reduced

• equilibrium position lies more to the right

negative standard electrode potential

• less likely to be reduced

• equilibrium position lies more to the left

electrode potential depends on

• temperature

• pressure

• concentration

reduction potentials

stay the same regardless of reference

Salt Bridge / Porous Cup

• maintain electrical neutrality in internal circut

• allow flow of ion

• complete internal circut

external circut

• pathway outside cell

• electrons flow from anode to cathode

• conducting wire that measures electrical energy produced

electrolyte solution

• substance that contains free ions

• conduct electricity

• aqueous solution, molten salt, solid with ionic conductivity

• facilitate flow of ions from one electrode to another

anion

• negatively charged ion

• move towards anode

• formed when atom gains electrons

cation

• positively charged ion

• move towards cathode

• formed when atom loses electrons

voltaic cell

• generate electricity

• spontaneous REDOX

• electrolyte, two electrodes, salt bridge, external circut

electrolytic cell

• consume electricity (forced reaction)

• non-spontaneous REDOX

• purify/extract metals and produce non-metals

• electrolyte, two electrodes, power supply

cathode voltaic cell

• gain electrons to become atoms

• attach to electrode

• positive

• reduction

anode electrolytic cell

positive

cathode electrolytic cell

negative

electrolytic cell / voltaic cell similarities

• REDOX reaction

• electrodes

• electrolyte

• circut completion

steps voltaic cell

1. identify SOA - cathode

2. identify SRA - anode

3. electrons leave anode to cathode - label flow

4. ions from salt move to half cells (anion → anode, cation → cathode)

5. identify evidence

shorthand cell notation

• Anode  |   Half-Cell Species  ||   Half-Cell Species  |   Cathode

• solid line - different state

• double line - salt bridge

inert electrode

do not participate in REDOX reaction

Brine Solution

• NaCl solution

• H₂O is the SOA

• Cl^- is the SRA

redox titration

• oxidizing agent titrated against reducing agent

• transition metal ions change color when changing oxidation state

• no indicator

titrant

• strong oxidizing agent

• significant color change when undergoing reduction

burette

• contains titrant

• known concentration

flask

• contains analyte

• unknown concentration

molar ratio ionic compounds

molar ratio from the half reaction used in unit analysis

Redox titration steps

• identify which species are OA or RA

• find SRA and SOA

• write the reaction for SRA and SOA (flip the SRA reaction)

• write equation for molar concentration

Redox reaction formula

(L_(given)/1) x (Mol_(given)/1) x (Mol ratio_(ionic)/Mol ratio_(ionic)) x (Mol ratio_(wanted)/Mol ratio_(given)) x (1/L_(wanted))

REDOX reaction

electron transfer from one to another

Predict Spontaneity step one - materials

list materials present

• weak acids stay as molecules

• ionic compounds disociate

• aqueous solutions have H₂O(l)

• strong acids have H⁺

Predict Spontaneity step two - OA and RA

identify oxidation agents and reduction agents

Predict Spontaneity step three - SOA and SRA

• use redox table to find the SOA and SRA

• SOA: top left, reduction half reaction

• SRA: bottom right, oxidation half reaction

Predict Spontaneity step four - balance half reaction

• balance half reaction based on electron transfer

• multiply one or both reaction with whole number coefficient

Predict Spontaneity step five - net ionic equation

• write net ionic equation

• use spontanrity rule to see if it reacts or not

oxidation number method step one - oxidation numbers

write oxidation numbers

oxidation number method step two - change in oxidation number

connect the atoms changing in oxidation number

oxidation number method step three - subscripts

identify oxidized and reduced and balance based on subscripts

oxidation number method step four - oxidation and reduction

use coefficients to make the increase in oxidation = reduction

oxidation number method step five - oxygen/hydrogen

balance oxygen and hydrogen with H₂O and H⁺

oxidation number method step six - check

check balancing for mass and charge

Disproportionation Reactions

reaction where same species is oxidized and reduced

OA and RA substances

• a material that is both an OA and an RA

• may react with itself to produce two different substances

Balancing Disproportionation Reactions

• write oxidation numbers

• increase or decrease in oxidation numbers

• identify oxidized and reduced and write half reaction (add electrons)

• balance hydrogen and oxygen

• check atom and charge balancing

Electronegativity

tendency of an atom or molecule to attract electrons

table of selected electrode potential

• left side: oxidation agent (gain electron)

• right side: reducing agent (lose electron)

top of table of selected electrode potential

• strongest OA

• weakest RA

bottom of table of selected electrode potential

• weakest OA

• strongest RA

positive Eº

• easily reduced

• better OA

negative Eº

• easily oxidized

• better RA

Spontaneity Rule

• oxidizing agent must be above the reducing agents

• left hand above right hand

spontaneous REDOX reaction

• standard electrode potential for the REDOX reaction is positive

• E°_{(REDOX Reaction)} = E°_{(Reduction ½ Reaction)} + E°_{(Oxidation ½ Reaction)}

• E°_{(REDOX Reaction)} = E°_{(Cathode/SOA)} – E°_(Anode/SRA)

Building a REDOX table

• look for gain/loss of electrons

• spontaneous: OA above RA

• non-spontaneous: RA above OA

photosynthesis

• 6 CO₂(g) + 6 H₂O(l) + energy → C₆H₁₂O₆(g) + 6 O₂(l)

• oxidized: H₂O

• reduced: CO₂

• oxidizing agent: CO₂

• reducing agent: H₂O

• 24 electrons lost/gained

• endothermic

cellular respiration

• C₆H₁₂O₆(g) + 6 O₂(l) → 6 CO₂(g) + 6 H₂O(l) + energy

• oxidized: C₆H₁₂O₆

• reduced: O₂

• oxidizing agent :O₂

• reducing agent: C₆H₁₂O₆

• 24 electrons lost/gained

• exothermic - releasing small amounts (liquid)

rust

• 4 Fe(s) + 3 O₂(g) + 6 H₂O(l) → 4 Fe(OH)₃(s) + energy

• oxidized: Fe

• reduced: O₂

• oxidizing agent :O₂

• reducing agent: Fe

• 12 electrons lost/gained

Corrosion

reaction of a metal with moisture/oxygen in the environment to produce a metal ore

corrosion process

• outside of water has more dissolved oxygen, attracting electrons

• iron metal loses electrons, transferring to the outside of the water

• iron ions enter the water and the middle starts to pit

corrosion prevention

• painting/coating the metal

• impressed current/cathodic prevention

• galvanizing the metal

• sacrificial anode

coating the metal

• paint or wax prevents water from reacting with iron

• small openings can cause rusting

sacrificial anodes

• highly reactive metals

• strong reducing agents

• consumed instead of the protected metal

• more negative electrochemical potential than protected metal

galvanizing the metal

• zinc is used as a protective oxide coating over the iron

• if it is scratched, zinc acts as a sacrificial anode

corrosion prevention/cathodic protection

• forces iron to act as a cathode

• saturating irons surface with electrons

• water takes electrons from coated layer rather than the metal

rusting nails

• water: rusts

• water + NaCl: rusts more and flakes

• water + oil: no rust

• CaCl₂: no rust

rusting nails explaination

• water: exposed to oxygen/water

• water + NaCl: exposed to oxygen/water/salt

• water + oil: exposed to water

• CaCl₂: exposed to oxygen

REDOX chemistry

• reduction-oxidation chemistry

• electron transfer

• oxidation and reduction occur together

oxidation

• half reaction - electron is a product

• anodes

• increase oxidation number

• reducing agents

• loss of electrons

reduction

• decrease oxidation number

• oxidizing agents

• gain electron

• cathodes

• half reaction - electron is a reactant

oxidation numbers

• single/molecular ion - amount of electrons have been lost/gained

• positive - loss of electrons

• negative - gain of electrons

• number given to each atom/ion in a compound to keep track of amount of electrons

rule one - free atoms

oxidation number is zero

rule two - fixed ox^#

atoms/ions have fixed oxidation number in compounds

• group one = +1

• group two = +2

• fluorine = -1

• hydrogen = +1

• oxygen = -2

excpetion to hydrogen

oxidation number is -1 in a metal hydride (H is anion)

exception to oxygen

• -1 in peroxide (O₂)

• +2 with fluorine

rule three - ionic compounds

ion’s oxidation number is its charge in ionic form

rule four - compounds/polyatomic ions

sum of all the oxidation numbers must equal the total charge

rule five - complex ions

oxidation number can be determined using charge on complex ion

rules six - electronegativity

more electronegative element is given priority for the Ox^# rule

reducing agent

• oxidized (loses electrons)

• reduce another atom by causing it to gain electrons

• oxidation number incerases

• facilitates oxidization

• metal atom (monoatomic)

• nonmetal/polatomic ion (anions)

oxidizing agent

• reduced (gain electrons)

• oxidizes another atom by causing it to lose electrons

• oxidation number of decreases

• facilitates reduction

• metal ions (cation)

• non-metal atoms (polyatomic)

half reactions

• show gain and loss of electrons

• cannot occur on its own

• electrons must equal the same

• balance mass

original definition

• oxidation is gain of oxygen

• reduction is loss of oxygen

second definition

• reduction is loss of oxygen and gain hydrogen

• oxidation is gain of oxygen and loss of hydrogen

current definition

• oxidation is the loss of electrons and increase in oxidation number

• reduction is gain of electrons and decrease in oxidation number