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
- %%free elements have an oxidation number of 0%%
- %%monatomic ions’ oxidation number is equal to the charge on the ion%%
- %%all alkali metals are +1%%
- %%all alkaline earth metals are +2%%
- %%aluminum is +3%%
- %%in about 90% of compounds, oxygen is -2%%
- %%hydrogen is 1+ with non-metals, and 1- with metals%%
- %%fluorine is -1%%
- %%halogens are negative unless paired with oxygen, then they’re positive%%
- %%in neutral molecules, the oxidation number adds up to 0%%
- %%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 = spontaneousOA / RA = not spontaneous