Download CHEMICAL REACTIONS

CHEMICAL REACTIONS
You will need to know/ memorize the following:
Solubility rules
The strong acids and bases
Oxidation states
The common oxidizing and reducing agents (and what the products are).
You should also review: naming and formulas, polyatomic ions, the diatomic elements, organic formulas
and naming, and balancing equations.
Net Ionic Equations:
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Ionic Equations: Instead of writing formulas for the substances as molecular formulas, they are
written as the ions that would be present in an aqueous solution.
Substances written as ions are: Soluble salts, strong acids, strong bases.
Net Ionic Equations: Once the ionic equation is written, any spectator ions (ions that appear in the
same form on both sides of the equation) are deleted, leaving only the ions that are actually involved
with the reaction.
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Ex: Pb(NO3)2 (aq) + 2KI(aq) → PbI2 (s) + 2KNO3 (aq) (molecular equation)
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) → PbI2(s) + 2K+(aq) + 2NO3-(aq)
(ionic equation)
Pb2+(aq) + 2I-(aq) → PbI2(s) (net ionic equation)
Major Reaction Types:
A.
B.
C.
D.
E.
Combination (synthesis) reaction: A + B →AB
Decomposition (analysis) reaction: AB → A + B
Combustion reaction: CxHy + O2 → CO2 + H2O
Single replacement (displacement) reactions: A + BC → AC + B
Double replacement (metathesis) reactions: AB + CD → AD + CB
A. Combination/Synthesis Reactions
1. metal and nonmetal form a salt (binary ionic compound)
• e.g. 2 Na + F2 Æ 2 NaF
2. two nonmetals form a binary covalent compound
• e.g. S + F2 Æ SF2
3. metal or nonmetal combines with oxygen to form binary ionic or covalent compound
• e.g. 2 Mg + O2 Æ 2 MgO
4. nonmetallic element combines with a binary covalent compound
• e.g. oxygen and nonmetallic oxide (sulfur dioxide + oxygen Æ sulfur trioxide)
• e.g. nonmetallic halide and additional halogen ( chlorine trifluoride and fluorine Æ chlorine
tetrafluoride)
5. Combination of two compounds
a. metallic oxide and nonmetallic oxide form an ionic compound (salt) with a polyatomic ion
• e.g. Na2O + CO2 Æ Na2CO3
b. metal oxides react with water to form bases (hydroxides)
• e.g. CaO + H2O Æ Ca(OH)2
c. nonmetal oxides react with water to form acids
• e.g. SO2 + H2O Æ H2SO3
d. hydrates result when anhydrous compounds react with water to form hydrates
• e.g. CuSO4 + 5 H2O Æ CuSO4·5H2O
e. Lewis acid (e- pr acceptor) and Lewis base (e- pr. donor)
• e.g. boron trifluoride (e- deficient) and ammonia (lone pair on N) : BF3 + NH3 Æ BF3NH3
B. Decomposition/ Analysis Reactions
1. decomposition (not necessarily heated):
a. hydrogen peroxide decomposes into water and oxygen.
b. ammonium hydroxide decomposes into ammonia and water.
2. thermal decomposition (are heated)
a. hydrogencarbonates (bicarbonates) yield carbonates, water and carbon dioxide (relatively low
temps)
• e.g. NaHCO3 Æ Na2CO3 + H2O + CO2
b. metallic carbonates decompose into metallic oxides and carbon dioxide (high temp.)
• e.g. CaCO3 Æ CaO + CO2
c. ammonium carbonate decomposes into ammonia, water and carbon dioxide
d. sulfites yield oxides and sulfur dioxide
• e.g. FeSO3 Æ FeO + SO2
e. oxides, chlorates and perchlorates yield oxygen
• e.g. 2HgO Æ Hg + O2
• e.g. 2KClO3 Æ 2KCl + 3O2
• e.g. NaClO4 Æ NaCl + 2O2
f. hydroxides, hydrates and some oxyacids release water
• e.g. Ca(OH)2 Æ CaO + H2O
• e.g. CuSO4-5H2O Æ CuSO4 + 5 H2O
• e.g. H2SO3 Æ H2O + SO2
• e.g. H2CO3 Æ H2O + CO2
3. Electrolysis - use of electricity to decompose compounds
Binary compounds decompose into two elements
• e.g. 2 NaCl Æ 2 Na + Cl2
C. Combustion Reactions
Combustion of hydrocarbons produces water and carbon dioxide
• e.g. 2 C6H14 + 19 O2 Æ 14 H2O + 12 CO2
D. Single Replacement/ Displacement Reactions
A more active element replaces a less active element in a compound.
Activity series for metals:
Most active Li Ca Na Mg Al Zn Fe Pb [H2] Cu Ag Pt
Least active
Activity series for non-metals
Most active F2 Cl2 Br2 I2
Least active
Note that the activity series can also be determined by looking at the standard electrode potentials
reference table supplied on the free response section of the AP exam.
1. Active metal replaces hydrogen in water or an acid
a. metal + acid Æ salt and hydrogen gas
• e.g. Mg(s) + 2 HCl(aq) Æ MgCl2(aq) + H2(g)
(Net : Mg + 2H+ Æ Mg2+ + H2)
b. Group 1A or 2A metal and water Æ base and hydrogen gas
• e.g. 2Li(s) + 2H2O(l) Æ 2LiOH(aq) + H2(g) (Net : Li + 2H2O Æ Li+ + OH- + H2)
2. a more active metal replaces a less reactive metal ion from solution(see activity series)
• e.g. ZnCl2(aq) + Mg(s) Æ MgCl2(aq) + Zn(s) ( Net : Zn2+ + Mg Æ Mg2+ + Zn )
3. a halogen replaces another halogen
• e.g. MgBr2(aq) + F2(g) Æ MgF2(aq) + Br2(l) ( Net : 2 Br - + F2 Æ 2 F - + Br2 )
4. Hydrides of alkali metals react with water to form hydroxides :
• e.g. LiH(s) + H2O(l) Æ LiOH + H2(g) (Net : LiH + H2O Æ Li+ + OH- + H2)
E. Double Replacement/ Metathesis Reactions
For a double replacement reaction to occur, ions must be removed from the solution by formation of a
precipitate, a gas, or a molecular compound.
Memorize solubility rules, strong acids and bases, and the following gas formers:
Common Gas Forming Reactions:
H2S
CO2
SO2
NH3
Common Gases
Any sulfide (S ) + any acid form H2S(g) and a salt
Any carbonate (CO32-) + any acid form CO2(g), H2O, and a salt
Any sulfite (SO32-) + any acid form SO2(g), H2O, and a salt
Any ammonium salt (NH4+) + a soluble strong hydroxide react
upon heating to form NH3(g), H2O, and a salt
2-
e.g: HC2H3O2(aq) + NaHCO3(aq) → NaC2H3O2(aq) + H2CO3(aq)
H2CO3(aq) → H2O(l) + CO2(g)
1. formation of an insoluble ionic compound (precipitate).
• e.g. BaCl2(aq) + MgSO4(aq) Æ BaSO4(s) + MgCl2(aq) ( Net :Ba2+ + SO42- Æ BaSO4(s))
2. neutralization reaction between an acid and a base (forms water)
a. strong acid/strong base :
• e.g. HCl + NaOH Æ H2O + NaCl (net : H+ + OH- Æ H2O)
b. strong acid/weak base
• e.g. HCl + NH3 Æ NH4Cl (net : H+ + NH3 Æ NH4+)
• e.g. HCl + CH3NH2 Æ CH3NH3Cl (net : H+ + CH3NH2 Æ CH3NH3+)
c. weak acid/strong base
• e.g. HC2H3O2 + NaOH Æ H2O + C2H3O2- (net : HC2H3O2 + OH- Æ H2O +
C2H3O2-)
d. weak acids and weak bases
• e.g. NH3 + HF Æ NH4+ + F- (net)
3. reactions with acids :
a. carbonates or bicarbonates and acids form a salt, water and CO2
• e.g. 2HCl + Na2CO3 Æ 2 NaCl + H2O + CO2 (net : H+ + CO32- Æ H2O + CO2)
b. sulfites and acids form a salt, water and SO2
• e.g. 2 HCl + Na2SO3 Æ 2 NaCl + H2O + SO2 (net : H+ + SO32- Æ H2O + SO2)
c. metallic sulfides and acids form H2S and a salt
• e.g. 2 HCl(aq) + Na2S(aq) Æ 2 NaCl(aq) + H2S(g) (net : H+ + S2- Æ H2S )
d. metallic hydrides and acids form H2 and a salt
• e.g. HCl(aq) + LiH(aq) Æ LiCl(aq) + H2(g) (net : H+ + H- Æ H2)
4. strong acids with salts of weak acids (salts containing anions such as : OCl-, ClO2-, F-, NO2-, CN-,
C2H3O2-)
a. metallic acetates and acids form acetic acid and a salt
• e.g. HCl(aq) + NaC2H3O2(aq) Æ NaCl(aq) + HC2H3O2(aq) (net : H+ + C2H3O2- Æ
HC2H3O2)
b. metallic nitrites and acids form nitrous acid and a salt
• e.g. HCl(aq) + NaNO2(aq) Æ HNO2(aq) + NaCl(aq) (net : H+ + NO2- Æ HNO2)
5. reactions with bases :
a. ammonium salts and soluble bases yield ammonia, water and a salt
• e.g. NH4Cl + LiOH Æ NH3 + H2O + LiCl (net : NH4+ + OH- Æ NH3 + H2O)
6. Amides with water
• e.g. sodium amide and water produce sodium hydroxide and ammonia
• Net : NaNH2 + H2O Æ Na+ + OH- + NH3
F. Oxidation Reduction Reactions (redox):
Memorize oxidation states and common oxidizing and reducing agents.
Terminology:
Oxidation - the oxidation number of one or more elements increases (it loses electrons).
Reduction - the oxidation number of one or more of the elements decreases (it gains electrons).
Oxidizing agent - A substance that causes oxidation of another substance (it is, itself, reduced)
Reducing agent - A substance that causes reduction of another substance (it is, itself, oxidized).
(OIL RIG or LEO says GER)
The most active elements in the activity series are the most powerful reducing agents.
Redox Reactions - involve a transfer of electrons (something is oxidized and something else is reduced)
and a change of oxidation state. Single replacement and combustion reactions are always redox reactions.
Combination and decomposition reactions are sometimes redox reactions. If a reaction takes place in an
acidic or basic solution it is most likely a redox reaction. Reactions involving elements in their natural
states (e.g. Al, Fe, Cl2, O2, etc.) are redox reactions.
Balancing Oxidation-reduction reactions (Oxidation States Method):
1.
2.
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4.
5.
6.
Assign oxidation numbers to all atoms in the equation
Identify the substance being oxidized and determine the number of electrons lost.
Identify the substance being reduced and determine the number of electrons gained.
Use coefficients to balance the atoms in the substances oxidized and reduced.
Use coefficients to balance the electrons gained and lost.
Use coefficients to balance the non redox substances.
Example:
CdS + I2 + HCl → CdCl2 + HI + S
1.
2.
3.
4.
Oxid #'s
Cd2+ S2- + I20 + H+ Cl- -> Cd2+ Cl2- + H+ I- + S0
Oxidation
Sulfur changes from a 2- to a 0 oxidation state, 2 electrons are lost.
Reduction
Iodine changes from 0 to -1 oxidation state, 1 electron is gained.
Balance atoms Because iodine is diatomic on the reactant side and both atoms must be reduced,
there must be 2 HI molecules formed on the product side.
CdS + I2 + HCl -> CdCl2 + 2HI + S
5. Balance e's Since sulfur loses 2 electrons and each iodine gains 1 (for a total of two), the number
of electrons lost and gained is balanced.
6. Balance atoms Balance the hydrogen and chlorine atoms.
CdS + I2 + 2HCl -> CdCl2 + 2HI + S
Example: Balance the following redox reaction
Cl2 + Ca(OH)2 -> CaCl2 + Ca(ClO3)2 + H2O
Note: This is an example of a disproportionation reaction - the same reactant (Cl2) undergoes both
oxidation and reduction.
1. Simple redox reactions :
a. Hydrogen displacement
• e.g. Ca(s) + H2O Æ Ca(OH)2(aq) + H2(g)
b. Metal displacement
• e.g. Zn(s) + CuCl2(aq) Æ ZnCl2(aq) + Cu(s) (Net : Zn + Cu2+ Æ Zn2+ + Cu)
c. Halogen displacement
• e.g. Cl2(g) + MgBr2(aq) Æ MgCl2(aq) + Br2(l) (Net : Cl2 + Br- Æ Cl- + Br2)
d. Combustion reactions.
• e.g. CH4 + O2 Æ CO2 + H2O
e. Combination and decomposition reactions (see above)
2. Reactions involving transition metals with multiple oxidation states :
• e.g. tin(II) ion with Fe(III) ion (net : Sn2+ + Fe3+ Æ Sn4+ + Fe2+)
3. Free halogens in dilute basic solutions form hypohalite ions
• e.g. Cl2(g) + KOH(aq) Æ KClO(aq) + KCl(aq) + H2O(l) (Net : Cl2 + OH- Æ ClO- + Cl- + H2O)
4. Redox reactions involving oxyanions such as Cr2O72• e.g. 14H+(aq) + Cr2O72-(aq) + 6Fe2+(aq) Æ 2Cr3+(aq) + 7 H2O(l) + 6Fe3+
The only way to be able to predict these types of reactions is to memorize common oxidizing and
reducing agents.
5. Atypical redox reactions
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hydrogen reacts with a hot metallic oxide to produce the elemental metal and water
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a metal sulfide reacts with oxygen to produce the metallic oxide and sulfur dioxide
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chlorine gas reacts with dilute sodium hydroxide to produce sodium hypochlorite, sodium chloride,
and water
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Copper reacts with concentrated sulfuric acid to produce copper (II) sulfate, sulfur dioxide, and water
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Copper reacts with dilute nitric acid to produce copper (II) nitrate, nitrogen monoxide, and water
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Copper reacts with concentrated nitric acid to produce copper (II) nitrate, nitrogen dioxide, and water
G. Simple Organic Reactions
1. Addition reactions
Usually occur when halogens or hydrogens are added to alkenes or alkynes. The double or triple bond
is eliminated and the product contains single-bonds (an unsaturated compound becomes saturated)
• e.g. Br2 + CH2=CH2 Æ CH2Br-CH2Br
• e.g. H2 + CH2=CH2 Æ CH3-CH3
2. Substitution reactions
Occur when an atom attached to carbon is replaced by something else.
• e.g. Cl2 + CH4 Æ CH3Cl + HCl
3. Combustion Reactions
When an organic compound is oxidized. Complete combustion results in carbon dioxide and water.
• e.g. C2H5OH + 3O2 Æ 2CO2 + 3H2O
4. Esterification reactions
A carboxylic acid is combined with an alcohol to produce an ester and water. Reactions are usually
aided by the addition of a small amount of an inorganic acid such as sulfuric to act as a dehydrating
agent.
• e. g. CH3COOH + CH3OH Æ CH3COOCH3 + HOH
H. Complex ion reactions (coordination chemistry) :
Usually involve transition metals and common ligands (see below). Since many complex ion
reactions occur with the addition of an excess of a substance containing the ligand, also watch for the
terms excess and concentrated in the problem.
Common complex ion metals (Lewis acids) are: Fe, Co, Ni, Cr, Cu, Zn, Ag, Al
Common ligands (Lewis bases) are: NH3, CN-, OH-, SCNGeneral Rule: The number of ligands will often be twice the charge of the metal ion.
1. Formation of ammine complexes
a. Metal ion solutions react with an excess of concentrated ammonia to form ammine complex
ions
• e.g. Cu2+ + NH3 Æ [Cu(NH3)4]2+
b. solid metal hydroxides, when combined with concentrated ammonia solution, produce ammine
complex ions and hydroxide ions.
• e.g. NH3(aq) + Cu(OH)2(s) Æ [Cu(NH3)4]2+ + OH2. Formation of cyanide complexes
a. metal ion solutions react with an excess of cyanide solution to form cyano complex ions
• e.g. CN- + Ag+ Æ [Ag(CN)2]3. Formation of hydroxo complexes
a. metal ion solutions react with an excess of hydroxide solution to form hydroxo complex ions.
b. Solid metallic hydroxides when added to an excess of hydroxide solution, produce hydroxo
complex ions.
4. Formation of thiocyanato complexes.
a. Metal ion solutions react with thiocyanate solution to form thiocyanato complex ions.
5. Dissolution of a complex ion with an acid.
A complex ion solution is treated with a strong acid solution resulting in the free metal ion or a
metal precipitate and the ligand ion.
• e.g. A solution of diamminesilver(I) chloride is treated with dilute nitric acid
[Ag(NH3)2]+ + H+ + Cl- Æ AgCl(s) + NH4+
• e.g. Excess dilute nitric acid is added to a solution containing the tetraamminecadmium(II)
ion.
H+ + [Cd(NH3)4]2+ Æ Cd2+ + NH4+
6. Lewis acid-base reaction
An electron pair acceptor Lewis acid) reacts with an electron-pair donor (Lewis base) to form a
coordinate covalent compound.
• e.g. BF3(g)+ NH3 (g) Æ H3N:BF3
Predicting Reactions on the AP exam:
** Please note that the format of this question has been changed for the 2007 exam and may not be
reflected yet in AP review guides that you purchase.
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You will be given 3 reactions which describe only the reacting substances.
For each reaction you will need to write a balanced equation showing reactants and products.
All reactants are in aqueous solution unless otherwise stated (watch for words such as solid or
crystals, or a piece of..).
All equations must be written as net ionic equations. You must cancel out spectator ions.
Charges on ions must be correct.
You need not include states of matter (s, aq, l, g)
The majority of the points on this question will be awarded for the correct balanced equation.
The remaining credit for each question will be earned for correctly answering a question
about the reaction. Ex: What are the reactant spectator ions?
Which substance is being oxidized?
What is the change in the oxidation number of one of the elements?
What changes would you observe during the reaction?