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-
    • </li>

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%%

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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

  </p>

  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%%

• </p>

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

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