Download - Net Ionic and Redox reactions
NET IONIC REACTIONS
All ions break up or “dissociate” in aqueous solutions
Net ionic reactions only show the elements that are actually reacting.
• Reaction = change states• ie. aqueous solid, precipitate, liquid, or gas
Spectator ions- are not changed (stay in an aqueous solution on both sides)
AgNO3(aq) + NaCl(aq) → AgCl(ppt) + NaNO3(aq)• The compounds in aqueous solution can really be written as
the separate ions, because they will break up in the solution.• ONLY the aqueous solutions break up
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(ppt) + Na+
(aq) + NO3-(aq)
• This is a complete ionic equation – shows all ions in the reacting solutions.
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(ppt) + Na+(aq) + NO3
-(aq)
Some of these ions do not participate in the reaction (they stay the same state on each side).
• These are called spectator ions. • You can remove them from the equation, giving the net ionic equation
Ag+(aq) + Cl-(aq) → AgCl(ppt)
In net ionic equations you must balance the atoms AND THE CHARGE.
HCl(aq) + ZnS(aq) → H2S(g) + ZnCl2(aq)
Pb(ClO4)2(aq) + NaI(aq) → PbI2(ppt) + NaClO4(aq)
Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g)
Ca(OH)2(aq) + H3PO4(aq) → Ca3(PO4)2(aq) + H2O(l)
Pb(C2H3O2)2(aq) + HCl(aq) → PbCl2(ppt) +
HC2H3O2(aq)
(NH4)2S(aq) + Co(NO3)2(aq) → CoS(s) + NH4NO3(aq)
REDOX REACTIONSRedox (reduction-oxidation) reactions occur when there is a transfer of electrons
• Oxidation reaction: Loss of electrons • Often see a gain of oxygen/ loss of
hydrogen• Reduction reaction: Gain of electrons
• Often see a loss of oxygen /gain of hydrogen
LEO the lion says GER
Reduction and Oxidation always happen TOGETHER
OXIDIZING AND REDUCING AGENTS
Reducing Agent - donates electrons (is oxidized)
Oxidizing Agent - accepts electrons (is reduced)
OXIDATION NUMBERSThe oxidation number of a monatomic (ie. not polyatomic) ion is equal in # and sign to its ionic charge.
The oxidation number of hydrogen in a compound is always +1 (except in metal hydrides where it is -1. ex: NaH)
The oxidation number of oxygen in a compound is always -2. (Except in peroxides, where it is -1).
The oxidation number of an uncombined element is zero.
For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal zero.
For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.
BALANCING REDOXStep 1: Write the equation in ionic form.
Step 2: Write separate half-reactions for oxidation and reduction.
Step 3: Balance the atoms in the half-reactions. Use H2O and H+ to balance oxygen and
hydrogen in an acid solution. Use H2O and OH- for reactions in a basic solution.
Step 5: Multiply each half-reaction by an appropriate number to make the electron charges equal.
Step 6: Add the half-reactions and subtract terms that appear on both sides of the equation.
Step 7: Check the final equation to be sure that atoms are conserved, charge is conserved, and all electrons have canceled. If spectator ions are known, add them to balance the equation.