1. Chemical Reaction Types
PS21 September 23, 2016
ACOS #5, p. 40…”observe and analyze properties of substances before and after chemical reaction – Chemical Reaction types.”

2. Aqueous Reaction Types
Will focus on reactions in aqueous solution and their driving forces
Precipitation Reactions
Gas-forming Reactions
Acid-base Reactions
Redox Reactions

3. Solutions
Solutions are defined as homogeneous mixtures of two or more pure substances.
The solvent is the substance present in greatest abundance.
All other substances are solutes.

4. Dissociation
When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them.
This process is called dissociation.

5. Electrolyte
An electrolyte is a substances that dissociates into ions when dissolved in water.
A nonelectrolyte does not dissociate when dissolved in water

6. Electrolytes
A strong electrolyte dissociates completely when dissolved in water.
A weak electrolyte only dissociates partially when dissolved in water.

7. Electrolytes Strong Electrolyte Weak Electrolyte Nonelectrolyte
Ionic All None None
Molecular Strong acids Weak acids & All other
bases compounds .

8. Electrolytes Nonelectrolytes
Soluble ionic compounds are normally electrolytes
Nonelectrolytes
Molecular compounds tend to be nonelectrolytes, except for acids and bases.

9. Solution Chemistry
It is helpful to pay attention to exactly what species are present in a reaction mixture (i.e., solid, liquid, gas, aqueous solution).
If we are to understand reactivity, we must be aware of just what is changing during the course of a reaction.

10. AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
Molecular Equation
The molecular equation lists the reactants and products in their molecular form:
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

11. Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)
Ionic Equation
In the ionic equation, all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions.
This more accurately reflects the species that are found in the reaction mixture:
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)

12. Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)
Net Ionic Equation
To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right:
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq)  AgCl(s) + K+(aq) + NO3−(aq)

13. Net Ionic Equation
The only things left in the equation are those things that change (i.e., react) during the course of the reaction:
Those things that didn’t change (and were deleted from the net ionic equation) are called spectator ions:
K+ and NO3- are spectator ions
Ag+(aq) + Cl−(aq)  AgCl(s)
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq) 
AgCl(s) + K+(aq) + NO3−(aq)

14. Writing Net Ionic Equations
Write a balanced molecular equation.
Dissociate all strong electrolytes.
Cross out anything that remains unchanged from the left side to the right side of the equation.
Write the net ionic equation with the species that remain.

15. Reaction Types in Aqueous Solution
Precipitation Reactions
Gas-Forming Reactions
Acid-base Reactions
Redox Reactions
Can be identified from the net ionic equation

16. Metathesis (Exchange) Reactions
Metathesis comes from a Greek word that means “to transpose.”
It appears as though the ions in the reactant compounds exchange, or transpose, ions:
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
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Being Careful
Molecular equation: 2 KCl(aq) + Cu(NO3)2 (aq)  2 KNO3 (aq) + CuCl2(aq) Ionic equation: 2 K+(aq) + 2 Cl-(aq) + Cu2+(aq) + 2 NO3-(aq)  2 K+(aq) + 2 NO3-(aq) + Cu2+(aq) + 2 Cl-(aq) Net ionic equation: none – no reaction
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18. Precipitation Reactions
When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), a precipitate is formed.
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19. Example Precipitation Reaction
Molecular equation: Pb(NO3)2(aq) + K2CrO4(aq)  2 KNO3(aq) + PbCrO4(s) yellow Ionic equation: Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + CrO42-(aq)  2 K+(aq) + 2 NO3-(aq) + PbCrO4(s) Net ionic equation: Pb2+(aq) + CrO42-(aq)  PbCrO4(s) Precipitation is the driving force

20. Solubility Guidelines
All Compounds of Na+, K+, NH4+ (most Li+)
All Compounds of NO3-, NO2-, ClO3-, ClO4-, {CH3COO- Slightly Soluble}
All Compounds of Cl-, Br -, I-, SCN- (except Ag+, Hg22+ ,Pb2+ {PbCl2 soluble in hot water})   
All compounds of SO42- (except Sr2+, Ba2+, Pb2+, {Ag+, Ca2+, Hg2+ Slightly})
Hydroxides: LiOH, NaOH, KOH, NH4OH, Ca(OH)2, Sr(OH)2, Ba(OH)2

21. Acids
The Swedish physicist and chemist S. A. Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water.

22. Strong Acids There are only six strong acids (strong electrolytes):
Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Perchloric (HClO4)
(although some books include chloric acid, HClO3)

23. Bases
Arrhenius defined bases as substances that increase the concentration of OH− when dissolved in water.

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Bases
The strong bases (strong electrolytes) are the soluble metal salts of hydroxide ion:
Alkali metals - LiOH, NaOH, KOH, CsOH, RbOH
Calcium – Ca(OH)2
Strontium – Sr(OH)2
Barium – Ba(OH)2
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25. Acid-Base/Neutralization Reactions
In an acid–base reaction, the acid donates a proton (H+) to the base.
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26. Neutralization Reactions
Generally, when solutions of an acid and a base are combined, the products are a salt and nonelectrolyte (often water):
CH3COOH(aq) + NaOH(aq)  CH3COONa(aq) + H2O(l) salt water
Net ionic equation:
CH3COOH(aq) + OH-(aq)  CH3COO-(aq) + H2O(l)
Driving force: formation of nonelectrolyte and salt

27. Neutralization Reactions
When a strong acid reacts with a strong base, the net ionic equation is H+(aq) + OH−(aq)  H2O(l)
E.g., Molecular equation: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) Ionic equation: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  Na+(aq) + Cl-(aq) + H2O(l)
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28. Gas-Forming Reactions
At least one of the products is a gas
Driving force is production of the gas

29. Gas-Forming Reactions
Some metathesis reactions do not give the product expected.
In this reaction, the expected product (H2CO3) decomposes to give a gaseous product (CO2):
CaCO3(s) + 2 HCl(aq) CaCl2(aq) + CO2(g) + H2O(l) Ionic equation: CaCO3(s) + 2 H+(aq) + 2 Cl-(aq)  Ca2+(aq) + 2 Cl-(aq) + CO2(g) + H2O(l) Net ionic equation: CaCO3(s) + 2 H+(aq)  Ca2+(aq) + CO2(g) + H2O(l)

30. Gas-Forming Reactions
When a carbonate or bicarbonate reacts with an acid, the products are a salt, carbon dioxide, and water:
CaCO3(s) + HCl(aq) CaCl2(aq) + CO2(g) + H2O(l) NaHCO3(aq) + HBr(aq) NaBr(aq) + CO2(g) + H2O(l)

31. Gas-Forming Reactions
Similarly, when a sulfite reacts with an acid, the products are a salt, sulfur dioxide, and water:
SrSO3(s) + 2 HI(aq) SrI2(aq) + SO2(g) + H2O(l)

32. Gas-Forming Reactions
This reaction gives the predicted product, but you had better carry it out in the hood, or you will be very unpopular!
But just as in the previous examples, a gas is formed as a product of this reaction:
Na2S(aq) + H2SO4(aq)  Na2SO4(aq) + H2S(g)

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Household Examples
Baking Soda
NaHCO3 + H+ → Na+ + CO2 + H2O
Baking Powder
14 NaHCO3 + 5 Ca(H2PO4)2 → 14 CO2 + Ca5(PO4)3OH + 7 Na2HPO4 + 13 H2O
Alka Seltzer
Citric acid + 3 HCO3-  citrate + 3 CO2 +
3 H2O
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34. Types Are Not Exclusive
By coincidence, all the gas forming reaction above are also???

35. Oxidation-Reduction (Redox) Reactions
An oxidation occurs when an atom or ion loses electrons.
A reduction occurs when an atom or ion gains electrons.
One cannot occur without the other.
Driving force is transfer of electrons

36. Oxidation Numbers
To determine if an oxidation–reduction reaction has occurred, we assign an oxidation number to each element in a neutral compound or charged entity.

37. Oxidation Numbers
Elements in their elemental form have an oxidation number of 0.
The oxidation number of a monatomic ion is the same as its charge.

38. Oxidation Numbers
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
Oxygen has an oxidation number of −2, except in the peroxide ion, in which it has an oxidation number of −1.
Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.

39. Oxidation Numbers
Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.
Fluorine always has an oxidation number of −1.
The other halogens have an oxidation number of −1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions.

40. Oxidation Numbers
The sum of the oxidation numbers in a neutral compound is 0.
The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

41. Identifying Redox Reactions
Look for element(s) with different oxidation number on each side of equation
Often easy to find by looking for elemental form of an element on only one side of equation
2 Ca(s) + O2(g)  2 CaO(s)
Oxidation numbers:
Ca O Ca +2 O -2

42. Displacement Reactions
In displacement reactions, ions oxidize an element.
The ions, then, are reduced.

43. Displacement Reactions
In this reaction,
silver ions oxidize
copper metal:
Cu(s) + 2Ag+(aq)  Cu2+(aq) + 2Ag(s)

44. Displacement Reactions
The reverse reaction,
however, does not
occur:
Cu2+(aq) + 2Ag(s)  Cu(s) + 2Ag+(aq) x
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45. Activity Series