part 1 redox rxn

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 righ%%t 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%%
    1. %%in neutral molecules, the oxidation number adds up to 0%%
    2. %%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