Explain what the section will be like LATER and then do examples



Explain what the section will be like LATER and then do examples
Writing Reactions
Preparation for AP Exam Free
Response Question #4
In this question, you will be given just the reactants of a reaction, described in
words. You will have to write and balance the entire reaction (predicting
products) and also answer a question (????) about each reaction.
This section was different a couple years ago, but it is still the hardest section
for chemists and students. We’ll do our best to prepare!
You need to be familiar with the various reaction
types to help predict the products:
5 types of reactions
- Double Replacement
- Lewis Acid-Base
- Organic
- Redox
- Complex Ion Formation
Precipitation Reactions
(“double replacement”)
• two ionic compounds (or acids/bases) mix and
form two new compounds.
• one of the new combinations of ions prevents
the ions from reforming the original compounds
- because it forms a stable precipitate
- one product breaks apart forming gas, which
leaves the solution
Preparation & Hints
• Write water as HOH and acids with hydrogens out front
and recognize that it is H+ and OH• Memorize:
•special metal & polyatomic ions and charges
•solubility rules to predict precipitates
•strong acids and bases so that you recognize the
weak ones (write strong acids and bases as ions,
weak in molecular form)
•the double replacement products that break into
 you can think of or write chemicals in different ways for
different reactions. Ammonia may be NH3(aq) for
complex ions or NH4OH for double replacement or
acid-base reactions
A potassium chromate solution and barium nitrate solution are combined.
molecular equation:
K2CrO4(aq) + Ba(NO3)2 (aq) 
complete ionic equation: 2K+ + CrO42- + Ba2+ + 2NO3-  2K+ + CrO42- + Ba2+ + 2NO3Net ionic equation: CrO42-(aq) + Ba2+(aq)  BaCrO4(s)
Always write the net ionic equation! All strong electrolytes should be shown as free ions
(eliminate spectators) and weak electrolytes shown as precipitates (solids).
You need to balance equations but you don’t need to include the phases of the reactants
and products.
Solutions: split into ions (unless insoluble)
Solids, crystals, chunks: these are not ions, even if they are things that should be soluble
A carbonic acid solution is combined with a calcium chloride solution.
H2CO3(aq) + Ca2+ --->
2 H+(aq) + CaCO3(s)
Solutions of ammonia and sulfuric acid are mixed.
strA + wkB = salt + HOH + H+
H2SO4 + NH3 
H+ + SO42- + NH4OH  NH4+ +SO42- + H2O
H+ + NH4OH  NH4+ + H2O
wkA + strB = salt + HOH + OHstrA + strB  salt + HOH (Net ionic always: H+ + OH- =  H2O )
wkA + wkB = depends
Solid sodium oxide reacts with sulfuric acid
Na2O + H2SO4  Na2SO4 + H2O
Na2O + H+  Na+ + H2O
Sulfur trioxide gas is bubbled through a sodium hydroxide solution
2NaOH + SO3  Na2SO4 + H2O
OH- + SO3  SO42- + H2O
Know the strong acids &
Gas-Forming Reactions
Sulfuric acid erodes a limestone statue (solid calcium carbonate)
CaCO3(s) + H2SO4(aq) --->
CaSO4(s) + H2CO3(aq)
Carbonic acid is unstable: H2CO3(aq) ---> CO2 + water
CaCO3(s) + H+ (aq) ---> Ca2+ + CO2 + H2O
Lewis Acid-Base Reactions
• Two species come together and share electrons in a coordinate covalent
bond. The species who donates the electrons is the Lewis base and the
accepting species is the Lewis acid.
Preparation & Hints
• “Have pair will share” = Lewis base; e- acceptors are Lewis acids
• Aluminum ions and Fe3+ ions make a solution acidic because they are
Lewis acids, drawing e- away from O in H2O, making the O-H bond more
polar, causing H+ to pull off H2O molecules
• Watch for: NO3-, NO2-, CO32-, SO32-, SO42- ions in solids that are heated in a
vacuum (not air), this is not combustion. A gas (ex. CO2) and an anhydride
(ex. CaO) are formed.
• Two uncombined elements can form a salt.
• A Lewis acid and base can also form a salt: CO2 + CaO  CaCO3
• When metals and metal hydrides react with water, they usually form bases
and hydrogen gas.
• If a gas compound is bubbled through a solution, water is acting like a
Lewis base or acid.
Sulfur dioxide gas is bubbled through water
SO2(g) + H2O(l) H2SO3(aq)
Nonmetal oxides (CO2 SO2 SO3 NO2 P4O10) form acids in
A calcium oxide solution is prepared.
CaO(s) + H2O(l) --> Ca(OH)2(aq)
Metal oxides (MgO, CaO, Na2O) form bases in water
Organic/Combustion Reactions
• Involves hydrocarbons, commonly combustion
but quite often other reactions
Preparation and Hints
• Combustion produces carbon dioxide and water
• If something reacts with an alkene or alkane, it
results in substitution for hydrogen
• Combustion or burned: O2 is a reactant
Methane gas is combusted in air
products depend on how much oxygen is available
O 2 plentiful :
CH 4  2O 2  CO2  2H 2 O
O 2 limited :
2CH 4  3O 2  2CO  4H 2 O
O 2 very limited : CH 4  O 2  C  2H 2 O
Methane reacts with chlorine gas
CH4 + Cl2  CH3Cl + HCl
cis-1,2 dichloroethene reacts with hydrogen gas
H Cl
+ H-H  Cl C C H
*Esterification with carb. acids
*Oxidation (they lose e- when lose a H) to form
aldehydes, ketones, or carb acids
Esters: Hydrolysis to yield carb. acids and alcohols
Ketones: Reduction to yield an alcohol
*Carbonyl groups (ketones AND aldehydes) react
with PCl5 to make acid chlorides
*Oxidation to form carb. acids
Amines: React with acids to form ammonium salts;
condensation to form amides
Complex Ion Formation
A transition metal (and some other metals) reacts and
gets surrounded by ligands, resulting in a charged
compound formed of multiple ions (ex Cu(NH3)42+)
Ligand: has lone pair to act as lewis base and bond with
metal ion( lewis acid)
Preparation and Hints
• A ligand can be a polar molecule (H2O, NH3) or a
negative anion: Cl-, OH-, SCN-, CN- , S2O32• ‘Excess’ and ‘concentrated’ may indicate complex ions,
not precipitates: AgNO3 + HCl forms a white precipitate
but concentrated HCl results in AgCl2- formation and
the solution clears.
• Watch for FeSCN2+
Central Metal Ion
Coordination number
Generally, you place twice
the number of ligands around
4 or 6
the central ion as the
on that central ion.
This is not always correct
but it will do.
Silver chloride dissolves in ammonia solution
AgCl(s) + 2NH3  [Ag(NH3)2]+ (aq) + ClAluminum hydroxide dissolves in a concentrated NaOH solution
Al(OH)3(s) + OH-  Al(OH)4- (aq)
Redox Reactions (“Oxidation-Reduction”)
•An oxidizing species accepts electrons from a reducing species. The
oxidation number of one element increasing as the other decreases.
Preparation & Hints
• Memorize:
-the strong oxidizers (ions with lots of oxygens) and what they turn
in to (ex HNO3 may be an acid or an oxidizer forming NO2,)
- the strong reducers and what they turn in to
• Look for:
- free neutral element in the reaction.
-battery reactions. Metal with the greatest E⁰ will reduce
- key words like ‘acidified solution’ or an acid included in the
reactants. The H+ will form H2O with oxygens from the oxidizer
• LeO says GeR
• When the hydrides of an alkali metal, Ca, Ba, or Sr dissolve in water,
hydroxides form and H2 gas is released (redox)
• When a hydroxide dissolves, hydroxides form but no gas is released
(not redox)
Redox Reactions
Single Replacement: Magnesium metal reacts with
hydrochloric acid Mg + 2H+ Mg2+ + H
Combination: sulfur reacts with oxygen
S + O2
+4 -2
Decomposition: potassium chlorate undergoes
decomposition 2KClO3
2KCl + 3O2
+1 +5 -2
+1 -1
Chlorine gas is bubbled through a strongly basic solution
Disproportionation Reaction: Element is simultaneously
oxidized and reduced
Cl2 + 2OHClO- + Cl- + H2O

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