Download some reaction rules File - District 196 e

Reaction Types (Hints)
I) Precipitation Reactions (double displacement reactions)
Simple Rules for the Solubility of Salts in Water
(Minor revisions to Zumdahl Table 4.1 are highlighted in bold print.)
1. Nitrate, acetate, chlorate salts are soluble.
2. Salts of alkali metals (Li, Na, K, Rb, Cs) and ammonium are soluble.
3. Chloride, bromide, iodide salts are soluble except those containing Ag+, Pb+2, Hg2+2.
4. Sulfate salts are soluble except BaSO4, PbSO4, Hg2SO4, CaSO4, SrSO4, Ag2SO4.
5. Hydroxides (and oxides) are insoluble except for alkali metals and ammonium. Hydroxides of
barium, strontium, calcium are marginally soluble.
6. Sulfides are insoluble except for those with ammonium, alkali metals, alkali earth metals.
7. Carbonates, chromates, phosphates, silicates are insoluble except for those with ammonium and
alkali metals.
II) Acid-Base Reactions (double displacement reactions)
Strong Acids –
Some Weak Acids –
Strong Bases -
HCl, HBr, HI, HClO4, HClO3, HNO3, H2SO4
HC2H3O2, H3PO4, H2CO3, H2S, HF, HNO2
LiOH, NaOH, KOH, RbOH, CsOH
Ba(OH)2, Sr(OH)2, Ca(OH)2
(soluble)
(marginally soluble)
The hydroxide ion (OH-) is such a strong base that it can be assumed to react completely with any
weak (and strong) acid. Therefore, the weak acid is not dissociated on the reactant side of the equation,
but is dissociated on the product side of the equation.
Example:
OH- (aq) + HF (aq)  H2O (liq) + F- (aq)
A similar argument can be made for weak bases in strong acids. The weak base is not dissociated on the
reactant side of the equation, but is dissociated on the product side of the equation.
Example:
H+ (aq) + NH4OH (aq)  H2O (liq) + NH4+ (aq)
III) Redox Reactions (includes synthesis, decomposition, single displacement and combustion)
Rules for Assigning Oxidation Numbers (from Zumdahl table 4.2)
1. Oxidation number of the atom of a free element is zero.
2. Oxidation number of a monoatomic ion equals its charge.
3. In compounds – oxygen has an oxidation number of -2 (except peroxides where it is -1).
4. In compounds – hydrogen has an oxidation number of +1.
5. In compounds – fluorine has an oxidation number of -1.
6. The sum of the oxidation states of an electrically neutral compound must be zero.
Activity Series (reactivity decreases from beginning to end)
Metals –
Li
K
Ba
Sr
Na
Ca
Ni
Sn
Pb
H2
Sb
Cu
Nonmetals - F
Cl
Br
I
Mg
Hg
Al
Ag
Zn
Pd
Cr
Pt
Fe
Au
Redox Reaction Prediction
Important Oxidizers
MnO4- (acid solution)
MnO4- (base solution)
MnO2 (acid solution)
Cr2O7-2 (acid)
CrO4-2
HNO3 (conc)
HNO3 (dilute)
H2SO4 (hot, conc)
Metallic ions
Free halogens
HClO4
Na2O2
H2O2
Formed in Reaction
Mn (II)
MnO2
Mn (II)
Cr (III)
Cr (III)
NO2
NO
SO2
Metallous ions
Halide ions
ClOHO2
Important Reducers
Halide ions
Free metals
Metallous ions
Nitrite ions
Sulfite ions
Free halogens (dilute, basic soln)
Free halogens (conc, basic soln)
C2O4-2
Formed in Reaction
Halogens
Metal ions
Metallic ions
Nitrate ions
Sulfate ions
Hypohalite ions
Halate ions
CO2
Redox reactions involve the transfer of electrons. The oxidation numbers of at least 2 elements
must change. Single displacement, synthesis (combination), decomposition and combustion
reactions are often redox reactions.
To predict the products of a redox reaction, look at the reagents given to see if there is both an
oxidizing agent and reducing agent. When a problem mentions an acid or base solution – it is
probably a redox reaction.
Balancing Redox Reactions
LEO the lions says GER
Lose Electrons – Oxidation – an increase in oxidation number
Gain Electrons – Reduction – a decrease in oxidation number
Oxidizing Agent – substance that is reduced (electron acceptor)
Reducing Agent – substance that is oxidized (electron donor)
Example:
Cu (s) + 2 Ag+ (aq)  2 Ag (s) + Cu+2 (aq)
Cu (s) is oxidized
Ag+ (aq) is reduced
Ag+ (aq) is the oxidizing agent
Cu (s) is the reducing agent
3 Methods for Balancing Redox Reactions
#1 “Condensed Oxidation State” Method (Zumdahl section 4.10)
Relies on inspection method. Limited use - does not account for reactions that need to be done
under acidic or basic conditions in order to balance them.
#2 -
“Half Reaction” Method (Zumdahl section 18.1 and Science Geek step by step section)
Works for balancing any redox reaction. Time consuming.
#3 -
“Hybrid” Method (shown below)
“Hybrid” Method
1. Write down charge numbers for each element in the equation – identify what is oxidized and what
is reduced.
2. Write a separate half reaction for the oxidation reaction and a separate half reaction for the
reduction reaction.
3. In each half reaction:
 Balance the element that is oxidized (or reduced)
 Add electrons lost (oxidation reaction) or electrons gained (reduction reaction) to the
appropriate side of each half reaction
 Rebalance each half reaction equalizing electrons
4. Return to original equation using coefficients obtained from your half reaction.
5. Balance charge by adding the H+ (acidic conditions) or OH- (basic conditions) to the appropriate
side of the equation.
6. Balance O and/or H by adding H2O to the appropriate side of the equation.
7. Recheck entire equation for balance (elements, charge, electron transfer).
Class Examples
MnO4- + Fe+2  Fe+3 + Mn+2
Cr2O7-2 + C2H5OH  Cr+3 + CO2
Decomposition Reactions
Following are some common decomposition reactions. Many of these reactions are initiated by
heat. The reverse of many of these reactions are synthesis reactions.
Ion in compound
Products
Example
Carbonate
CO3-2
oxide, CO2
CaCO3  CaO + CO2
Hydroxide
OH-
oxide, H2O
Ca(OH)2  CaO + H2O
Sulfate
SO4-2
oxide, SO3
CaSO4  CaO + SO3
Sulfite
SO3-2
oxide, SO2
CaSO3  CaO + SO2
Clorate
ClO3-
chloride, O2
2 KClO3  2 KCl + 3 O2
Ammonium
NH4+
acid, NH3
NH4Cl  HCl + NH3
Bicarbonate
HCO3-
H2O, CO2
H2CO3  H2O + CO2
How about the decomposition of ammonium carbonate?
 Expect ammonium to produce NH3
 Expect carbonate to produce CO2
 Figure out the final product based on what’s left over