1. Aqueous Reactions and Solution Stoichiometry
Chapter 4
Aqueous Reactions and Solution Stoichiometry

2. Aqueous Solutions Solutions in which water is the solvent
Solution: homogeneous mixture of two or more substances in a single physical state
Solvent: greater quantity; acts as the dissolving medium
Solute: lesser quantity; is dissolved by solvent

3. Electrolytes
A substance that conducts an electrical current when it is dissolved; contributes ions to the solution
Soluble salts, acids and bases are electrolytes
Nonelectrolyte: will not form ions in solution; will not conduct an electrical current

4. When in water… Ionic: ions dissociate and disperse thru sol’n
Each ion is surrounded by water molecules
Forms a stable complex and prevents recombination
Polyatomic ions remain in groups

5. When in water…
Molecular: bonding is strong enough to prevent dissociation with water molecules
The entire molecular compound is surrounded by water molecules
Become uniformly distributed thru water
No ions formed = nonelectrolyte
Acids, though covalent in nature, will ionize in aqueous solution

6. Strong & Weak Electrolytes
Strong: produce many ions in sol’n; conducts current strongly
Weak: produce few ions in sol’n; will conduct, but weakly
Non: produces no ions in sol’n; will not conduct
Weak electrolytes will establish an equilibrium between ions & molecule

7. Precipitation Rxns One insoluble product formed from aqueous reactants
Precipitate: insoluble solid formed as the result of a rxn in sol’n
Solubility: amount of a substance that may be dissolved in a given amount of solvent
Solubility <0.01 mol/L is insoluble
Insoluble = attraction is too great to be broken by attraction for water molecules

8. Solubility Guidelines (Memorize These!!)
Soluble Ions:
Nitrate
Acetate
Halides
Sulfate
Hydrogen carbonate
Ammonium
Chlorate
Perchlorate
Group 1 ions
Exceptions:
None
Ag, Hg, Pb
Ag, Ca, Sr, Ba, Pb

9. Solubility Guidelines (Memorize These!!)
Insoluble Ions:
Carbonate
Chromate
Phosphate
Sulfide
Hydroxide
Exceptions:
Group 1 or ammonium
Group 1, NH4+, Ca, Mg
Group 1, ammonium, Ca, Ba, Sr

10. Writing Ionic Equations
Complete ionic equations show all electrolytes in their aqueous ionic form
Spectator ions are those ions which are unchanged over the course of the reaction
Net ionic equations remove spectator ions
The AP exam will require you (as will I) to write all reactants and products in their ionic form!

11. Acids Acids ionize (separate to form soluble ions) in aqueous solution
All acids contribute hydrogen as the only positive ion in solution
Monoprotic: donates 1 H+
Diprotic: donates 2 H+
Triprotic: donates 3 H+
Polyprotic acids donate in steps

12. Bases
Bases are substances that act to reduce the H+ ion concentration in a solution
Hydroxide ions are the prototypical base
Other compounds may also be classified as bases, including NH3

13. Strength of Acids & Bases
Determined by electrolyte strength
Strong acid = strong electrolyte
Memorize: HCl, HNO3, HBr, HI, HClO4, HClO3, H2SO4
Weak acid = weak electrolyte
Strong Base = strong electrolyte
Memorize: group 1 metal hydroxides and heavy group 2 metal hydroxides
Weak base = weak electrolyte

14. Neutralization Reactions
An acid and a base will react to produce water and a dissolved salt
Net ionic reaction:
H+ + OH-  H2O

15. Oxidation-Reduction Rxns
Transfer of electrons between reacting substances
Oxidation: loss of electrons, increase in ox#
Reduction: gain of electrons, decrease in ox#
Ox Agent: gets reduced
Red Agent: gets oxidized
LEO GER-OAR

16. Oxidation Numbers
Apparent charge assigned to all elements in a compound based on the assumption that all electrons are possessed by the most electronegative atoms
Givens:
All uncombined elements (incl. diatomic) = 0
Monatomic ions = charge
Compounds: sum of all ox #s = 0
Polyatomic ions: sum of all ox #s = charge

17. Assigning Oxidation Numbers
Group 1 metals = +1
Group 2 = +2
Hydrogen = +1 (usually)
Oxygen = -2 (usually)
Halides = -1 (last resort)
Constants:
Al = +3 F = -1

18. Activity Series
A list of half reactions ordered to show the ease of oxidation or reduction of a substance
All metals above hydrogen will be oxidized by acid
Group 1 and 2 metals will be oxidized by water

19. Concentration of Solutions
Molarity: moles of solute per liter of solution
Molality: moles of solute per kg of solvent
Mole Fraction: moles component per total moles of solution
Normality: equivalence per liter of solution
Percent composition (pph): mass of component per total solution mass x 100
Parts per million: mass of component per total solution mass x 1,000,000

20. Dilutions
Proportionally weakening the concentration of a solution through the addition of solvent
M1V1 = M2V2
Remember, electrolytes will increase the # moles in a solution
Aliquot: a small sample of solution, usually used to form a more dilute solution

21. Solution Stoich
We can interconvert between moles, mass, volume of solution, or concentration of solution depending upon our data
How many moles of water form when 25.0mL of M nitric acid is completely neutralized by sodium hydroxide?

22. Titrations
A carefully controlled neutralization reaction; may be used to determine unknown concentrations
Standard solution: a solution of known concentration
Equivalence point: [H+] = [OH-]
End point: indicator changes color
N1V1 = N2V2