Electrochemistry and Nuclear Chemistry

OIL RIG

Oxidation is loss

Reduction is gain

Reduction/oxidation potential

Reduction potential: tendency to acquire electrons

Oxidation potential: tendency to lose electrons

Reduction in terms of O and H

Losing bonds to more EN atoms like O

Gaining bonds to less EN atoms like H

Oxidation in terms of O and H

Gaining bonds to more EN atoms like O

Losing bonds to less EN atoms like H

Oxidation number

More positive=more hungry for electrons

Rules:

-Oxidation state of any element in neutral state = 0

-Sum of oxidation states equals overall charge on molecule

-Group 1 metals have +1

-Group 2 metals +2

-Fluorine -1

-Oxygen -2

-H is +1 when bonded to more EN atom than H, -1 when less EN

-Halogens -1 and atoms of oxygen family -2

Inorganic compounds ranking

Think about oxidation states and least oxidized less oxygen

Least positive=most reduced

Most positive=least reduced

Transition metals like Mn variable charge so that is what you are checking!

K2MnO4 K2=2+ and so MnO4 has to be -2 and O4=-8 so Mn=+6

Organic compounds

CO2, CH2, HCOOH etc.

No oxidation states!

Think about how many bonds to oxygen

Greater amount of bonds to oxygen is more oxidized!

Oxidizing/reducing agent

Oxidizing agent: Cause others to become oxidized it itself is REDUCED

Neutral non metals. MOx (MnO4-) (CrO3)

Reducing agent: Cause others to become reduced it itself is OXIDIZED

H2 and neutral non metals like MeHx (LiAlH4 and NaBH4)

Reduction potentials

Table given

They are for the reactants

Strongest oxidizing agent=best at being reduced to most positive reduction potential

Strongest reducing agents=best at being oxidized, most negative reduction potential because will be REVERSED sign, looking at products!

Spontaneous

Ecell>0 positive

Ecell=

Ecell=Ered + E ox

AN OX RED CAT

Anode=oxidation

Cathode=reduction

deltaG=

-nFEcell

Ecell>0 than deltaG<0 so spontaneous

n=moles of electrons transferred

Reduction potential rule

If spontaneous subsequent molecules needs thigher reduction potential than what came before it

Example: Coating Zn2+ with Cu2+ know that it is spontaneous therefore know that Cu2+ has to have higher reduction potential! Elemental copper coating it

For a reaction to happen spontaneously, the "coating" or substance doing the oxidizing must have a higher reduction potential than the substance being oxidized.

Different problem example

Stainless steel is an alloy made by mixing iron with low percentages of carbon and chromium. The Cr forms a thin, protective layer of chromium(II) oxide on the surface of the metal. Which of the following regarding the relative reduction potentials (E°) of metallic ions must be true?

Cr goes to Cr3+ so oxidized so lower reduction potential

Time based calculations

2 A for 2 minutes, moles of Cu2+ can be plated out F=100,000 C/mol

1A=C/s!!!!

2 min (60 sec/1 min) (2C/S) (1mol/100,000 C) (1 mol Cu2+/2 mol e-)

All cells must consist of

Two or more electrodes (anode + cathode)

An electrolyte salt bridge or electrolyte solution (balance charge build up)

Circuit to connect

Galvanic means…

Spontaneous

NO external power source

Electrolytic means…

Non spontaneous

POWER source needed

Galvanic cell (new battery)

Anode (-) oxidation

Cathode (+) reduction

*Note + and - does not mean charge just an assignment

Electrons flowing from A to C

Current opposite of electron flow

Where do cations and anions flow?

Cations flow to CATHODE

Anions flow to ANODE

Electrolytic cell (DEAD battery)

Ecell=0

Equilibrium

Electrolytic cell (recharging battery)

Everything flipped from galvanic cell

+ and - assignments match sides of power source battery

Cathode and anode flipped so now cathode - and anode +

Cathode vs anode

Cathode

RED CAT

Cations flow

Plating

+ galvanic and -electrolytic

Anode

AN OX
Anions flow

Pitting

- galvanic + electrolytic

Think in galvanic cell think anode is oxidation which means they are losing and leaving behind electrons so - and cathode is reduction so gaining electrons and +

Example: Cu + 2Ag+ to Cu2+ + 2 Ag

Cu to Cu2+ is oxidation Cu anode (solids)

and Ag+ to Ag is reduction Ag cathode (solids)

Reverse changes so Cu cathode and Ag anode

Cell in equilibrium

E cell =0

E=Estandard state-RT/nF lnQ

Non zero standard cell potential

Estandard=RT/nF lnK

Redox titrations

Equivalence point (weighted average E sample + Etirant)is the exact moment when the amount of oxidizing or reducing titrant added is chemically equivalent (stoichiometrically equal) to the amount of analyte in the sample

Half equivalence point=E sample

Mid after equivalence point is E tirant

Example: 0.1 M tirant eq point is 50 L then 0.1 M (50 L) = moles than can know how many moles of each thing!

A and Z

A=Proton + neutron (mass number)

Z=Protons (atomic number)

Alpha decay

A=A-4

Z=Z-2

Very large nuclei

Least dangerous

Beta decay

A=same

Z=+1

Not enough protons

Nuclei with high neutron/proton ratio

More dangerous

Positron emission

A=same

Z=Z-1

Too many protons

Proton/neutron ratio high

Electron capture

Absorbed ny nucleus

Same as positron

A=same

Z=Z-1

P/N ratio high too many protons

Daughter vs parent

Daughter always more stable (less energy) than parent!!!

Gamma decay

-Nuclear rxns

-Most dangerous

Same A and Z just goes from Dy* to Dy

*=higher energy

Just about energy

Deutron

1 proton and 1 neutron

2H or D

Half lives

Longer half lives are safer

Past 10 half lives considered safe

t1/2=7 days 10% orginal amount \

Days % left

0 100%

7 50%

14 25%

21 12.5%

28 6.25%

10% between 21-28 days

Nuclear reactions exothermic not reversible, all about stability!

Energy released in terms of bonding energy

deltaE=BEparent-BEdaughter(s)

BE is given

Sample over time radioactive decay