Electrochemistrys

• all chemical species can attract and hold electrons in their outer shell

• the study of relationships between chemical reactions and the flow of electrons

Electronegativity

• an elements ability to attract electrons to itself (what’s their pull game like)

• Electronegativity increases from left to right and from the bottom to the top

  • basically increases the closer you get to chlorine

• Most electronegative atoms are

  • fluorine

  • oxygen

  • chlorine

  • nitrogen

  • bromine

• Least electronegative atoms are

  • radium

  • francium

  • cesium

  • rubidium

Two ways for electrons to move between species

1) Covalently bonded molecules

• Covalently bonded elements

  • equal sharing between species

  • ex. N2 H2

• Covalently bonded compounds

  • unequal sharing between species

  • both species involved have high electronegativity with a small difference

  • ex. H20 CO2

2) Ionic Bonds

• large difference in electronegativity

• electrons are transferred (no sharing involved)

• oxidation reduction reactions (redox reaction) are reactions where one or more electrons are transferred

  Oxidation and Reduction

• Oxidation Is Loss and Reduction is Gain

  • mnemonic device → OIL RIG

• redox reactions involve the transfer of electrons from one reaction to another

• if one substance is oxidized another substance in the same reaction must be reduced

• oxidation

  • where electrons are removed from and atom or ion

  • the species that loses electrons is oxidized

    • X → (X^+) + (e-)

    • Here X is oxidized because it loses e-

• reduction

  • where electrons are gained from an atom or ion

  • the species t hat gains electrons is reduced

• oxidizing agent

  • causes the oxidation of another species

  • the reduced species is the oxidizing agent since it’s what made the other species lose electrons

• reducing agent

  • causes the reduction of another species

  • the oxidized species is the reducing agent since it’s what made the other species gain electrons

    • Y + (e-) → (Y^-)

    • Y is reduced because it loses e-

redox reactions

• examples of redox reactions

  • reactions with batteries

  • burning of wood

  • corrosion of metals

  • ripening of fruit

  • combustion of gasoline

• 4 categories of reactions for redox

  • single replacement

    • all single replacement reactions are redox reactions

  • hydrocarbon combustion

  • synthesis

  • decomposition

Oxidation Numbers

• rules used to assign charges (+/-) to see of an electron transfer has happened in covalently bonded compounds. These charges are used follow the changes that occur in redox reactions

• an increase in the oxidation number means the substance is oxidized

• a decrease in the oxidation number means the substance is reduced

• usually metals are only positive and non-metals are negative and positive

• the highest oxidation number an element can have is their group number on the periodic table

• if there’s no change in oxidation numbers then it isn’t a redox reaction

  Oxidation Rules

  1. free elements have an oxidation number of 0

  2. monatomic ions’ oxidation number is equal to the charge on the ion

  3. all alkali metals are +1

  4. all alkaline earth metals are +2

  5. aluminum is +3

  6. in about 90% of compounds, oxygen is -2

  7. hydrogen is 1+ with non-metals, and 1- with metals

  8. fluorine is -1

  9. halogens are negative unless paired with oxygen, then they’re positive

  10. in neutral molecules, the oxidation number adds up to 0

  11. oxidation numbers don’t have to be integers, they can be fractions too

Activity Series and Standard Reduction Table

• purpose is to see possible redox reaction with different metals and metallic ions

• Spontaneous reaction happens without any added energy

• spontaneous reaction ex.

  • Zn + 2HCl→ ZnCl + H2

  • hydrogen must be a strong enough oxidizing agent to remove the electrons from zinc

  • zinc must have an electron affinity low enough for hydrogen to remove its electrons

  • in order for a reaction to be spontaneous, hydrogen must be a strong enough oxidizing agent to remove the electrons from zinc

• the standard reduction potential table lists the equilibrium reacts between species and the voltage (E^0) for each reaction

  • the forward reaction

    • moves left to right

    • reduction half reaction

    • reaction that is gaining electrons

    • ex. Al^3+ + 3e- → Al

  • the reverse reaction

    • moves right to left

    • oxidation half reaction

    • reaction that is losing electrons

    • ex. Al → Al^3+ + 3e-

• a substance’s tendency to gain electrons is it’s reduction potential

• in every redox reaction

  • the half reaction that’s more positive will continue as reduction reactions

  • the half reaction that’s more negative will continue as the oxidation reaction

• oxidizing agents are in the left-hand column

  • strongest oxidizing agent is F2

  • higher up on the left side column, the stronger the oxidizing agent will be

• reducing agents are in the right-hand column

  • strongest reducing agent is Li

  • farther down the right side column, the stronger the reducing agent will be

• some species are in both columns and can be both an oxidizing agent of a reducing agemt

• How to determine if a reaction is spontaneous or not?

  • if the oxidizing agent is higher in the Reduction Table than the reducing agent, the reaction will be spontaneous

  • OA \ RA = spontaneous

  • OA / RA = not spontaneous

Balancing redox rxns

method 1) using standard reduction table

method 2) writing balanced 1/2 rxn

method 3) oxidation numbers and 1/2 rxn

method 4) balancing whole equation