1. Chapter 4: Aqueous Reactions
AP Chemistry
Unit 2

2. Types of Aqueous Reactions
Oxidation–Reduction Reactions
Acid–Base Reactions
Precipitation Reactions

3. Acids
The Swedish physicist and chemist S. A. Arrhenius defined acids as substances that increase the concentration of H+ when dissolved in water.
Both the Danish chemist J. N. Brønsted and the British chemist T. M. Lowry defined them as proton donors.

4. Strong vs. Weak Acid = extent of ionization
Strong Acids – 100% ionize
ex:
Weak Acids - only a small fraction ionizes (establishes equilibrium)

5. Acids There are only seven strong acids: Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Chloric (HClO3)
Perchloric (HClO4)

6. Bases
Arrhenius defined bases as substances that increase the concentration of OH− when dissolved in water.
Brønsted and Lowry defined them as proton acceptors.

7. Bases The strong bases are the soluble metal salts of hydroxide ion:
Alkali metals
Calcium
Strontium
Barium

8. Acid-Base Reactions
In a Bronsted-Lowry acid–base reaction, the acid donates a proton (H+) to the base .

9. Arrhenius Reactions (Neutralization Reactions)
Generally, when solutions of an acid and a base are combined, the products are a salt and water:
CH3COOH(aq) + NaOH(aq) CH3COONa(aq) + H2O(l)

10. Neutralization Reactions
When a strong acid reacts with a strong base, the net ionic equation is
HCl(aq) + NaOH(aq) 

11. Neutralization Reactions
When a strong acid reacts with a strong base, the net ionic equation is
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq) 
Na+(aq) + Cl−(aq) + H2O(l)
H+(aq) + OH−(aq)  H2O(l)

12. Neutralization Reactions
Weak acid with Strong Base

13. Precipitation Reactions
When one mixes ions that form compounds that are insoluble (as could be predicted by the solubility guidelines), a precipitate is formed.

14. Precipitation Reaction
A double displacement reaction in which a solid forms and separates from the solution.
When ionic compounds dissolve in water, the resulting solution contains the separated ions.
Precipitate – the solid that forms.

15. The Reaction of K2CrO4(aq) and Ba(NO3)2(aq)
Ba2+(aq) + CrO42–(aq) → BaCrO4(s)

16. Precipitates
Soluble – solid dissolves in solution; (aq) is used in reaction equation.
Insoluble – solid does not dissolve in solution; (s) is used in reaction equation.
Insoluble and slightly soluble are often used interchangeably.

17. Simple Rules for Solubility
Most nitrate (NO3-) salts are soluble.
Most alkali metal (group 1A) salts and NH4+ are soluble.
Most Cl-, Br-, and I- salts are soluble (except Ag+, Pb2+, Hg22+).
Most sulfate salts are soluble (except BaSO4, PbSO4, Hg2SO4, CaSO4).
Most OH- are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble).
Most S2-, CO32-, CrO42-, PO43- salts are only slightly soluble, except for those containing the cations in Rule 2.

18. Metathesis (Exchange) Reactions aka Double Displacement Reactions
Metathesis comes from a Greek word that means “to transpose.”
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

19. 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)

20. 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)

21. 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)

22. Ag+(aq) + Cl−(aq)  AgCl(s)
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.
The only things left in the equation are those things that change (i.e., react) during the course of the reaction:
Remove spectator ions
Ag+(aq) + Cl−(aq)  AgCl(s)

23. Solution Stoichiometry
10.0 mL of a 0.30 M sodium phosphate solution reacts with 20.0 mL of a 0.20 M lead(II) nitrate solution (assume no volume change).
What precipitate will form?
What mass of precipitate will form?
What is the concentration of the nitrate ion left after the reaction?
The balanced molecular equation is: 2Na3PO4(aq) + 3Pb(NO3)2(aq) → 6NaNO3(aq) + Pb3(PO4)2(s).
mol Na3PO4 present to start and mol Pb(NO3)2 present to start. Pb(NO3)2 is the limiting reactant, therefore mol of Pb3(PO4)2 is produced. Since the molar mass of Pb3(PO4)2 is g/mol, 1.1 g of Pb3(PO4)2 will form.

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

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

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

27. Activity Series