1. Reactions in Aqueous Solutions

2. Types of Aqueous Solutions
Electrolytic—substances dissolve in water to conduct electricity.
Made by ionic compounds dissociating in water.
Nonelectrolytic—molecular compounds dissolve as intact molecules and do not conduct electricity.
Attraction of substance to water molecules overcomes attraction of substance to itself.

3. Strong vs. Weak Acids
An acid is a molecular compound that ionizes in water.
Strong acids—ionize completely
Examples: HCl, H2SO4, HClO4, HClO3, HI, HBr, HNO3
Produce strong electrolytic solutions
Weak acids—ionize partially
Examples—CH3COOH
Produce weak electrolytic solutions

4. Solubility Soluble—dissolves in water
Insoluble—does not dissolve in water.
Memorize solubility rules!

5. Precipitation Reactions
Forms a precipitate (solid) when two solutions are mixed.
Example—reaction between potassium iodide and lead (II) nitrate.
Draw reaction at particulate level.
Pay attention to ratio of molecules, size of ions, orientation of surrounding water molecules.

6. Equations for aqueous rxns
Molecular equation—shows complete, neutral formulas for every compound.
Complete ionic equation—shows all species as they are present in solution
Net Ionic Equation—shows only what reacts.

7. Give the net ionic equation for the reaction between ammonium carbonate and calcium chloride.

8. Acid-Base Reactions
Transfer of proton is made.

9. Acid-Base definitions
Arrhenius
Produces H+
Produces OH-
Bronsted-Lowery
Proton donor
Proton acceptor
Lewis
Electron acceptor
Electron donor
Amphoteric: can act as an acid or base (H2O)

10. Acids Ionize in water to produce H+ H+ = bare proton
Normally associates with water to form hydronium ion.
Polyprotic acids—contain more than one ionizable proton.
Diprotic—two ionizable protons
Triprotic—three ionizable protons

11. Bases Strong bases produce OH- ions in water.
Usually alkali metals and alkaline earth metals with hydroxide.
Weak bases—take protons from water molecules (only occurs to small extent).

12. Acid-Base reactions Produce water and salt.
Net Ionic equations show exactly what reacts.
Call them neutralization reactions because we neutralize the acidity and basicity to form water (neutral on pH scale).

13. Gas Evolution Reactions
Two aqueous solutions combine to form a gaseous product.
Often have intermediate products that are unstable and therefore, decompose.
Consider reaction between sodium bicarbonate and hydrochloric acid.

14. In general… Sulfides—hydrogen sulfide gas
Carbonates and bicarbonates—intermediate of hydrogen carbonate = CO2 evolved.
Sulfites and disulfites—intermediate of hydrogen sulfite = SO2 evolved
Ammonium—intermediate of NH4OH = NH3 evolved.

15. Oxidation-Reduction Reactions
AKA Redox reactions
Used daily for production of electricity.
Involves loss and gain of electrons

16. Redox reactions Oxidation = loss of electrons
Reduction = gain of electrons
Does not need to be a complete transfer of electrons like in an ionic compound.
Consider reaction between hydrogen gas and chlorine gas.

17. Oxidation State Our way of tracking electrons
Does not necessarily mean there is a charge on each atom.
Considered the charge of each atom if all electrons were assigned to the atoms with the greatest attraction for them.

18. Rules for assigning oxidation states
1. Oxidation state of a free element = 0
4. H = +1 in covalent compounds; -1 in metal hydrides.
2. Oxidation state of monatomic ion = charge
5. Fluorine = −1 in all compounds
3. Oxygen = −2 in covalent compounds (except in peroxides where it = −1)
6. Sum of oxidation states = 0 in compounds; Sum of oxidation states = charge of the ion for polyatomic ions

19. Assign oxidation states for each element in the following compounds:
1.) Cl2 2.) Na+ 3.) KF 4.) CO2 5.) SO42- 6.) K2O2

20. Identifying Redox reactions
Assign oxidation numbers to all atoms.
If there is a change in ox. numbers, a redox has occurred.
OIL RIG
LEO GER
Oxidation—increase in oxidation state
Reduction—decrease in oxidation state

21. Example
Consider reaction between Magnesium metal and water.

22. Oxidation and reduction MUST occur together.
Oxidizing agent—oxidizes the other substance.
Examples—oxygen, peroxides
Always reduced—why?
Reducing agent—reduces the other substance
Examples—Hydrogen, Group I and Group II
Always oxidized—why?

23. Practice
Identify what is reduced and what is oxidized in the following reaction:
H2O2 + C2H4  C2H4O + H2O
Identify the reducing agent and oxidizing agent.

24. Redox Reactions
A closer look…

25. Important Concepts Electricity = flow of electrons or ions.
Law of conservation of mass is always followed.

26. Balancing Redox reactions
Half-reaction method
1. Assign oxidation states and identify what is oxidized and reduced.
2. Separate overall rxn into two half-rxns.
3. Balance each half-rxn with respect to mass in this order:
Balance all elements other than H and O.
Balance O by adding H2O
Balance H by adding H+.
4.Balance each half-rxn with respect to charge.
5. Make electrons equal in both half-rxns.
6. Add both half-rxns together

27. Practice
The following reactions are under acidic conditions using the half-reaction method:
Al(s) + CuCl2(aq)  AlCl3(aq) + Cu(s)
Cu(s) + NO3-(aq)  Cu2+(aq) + NO2(g)

28. Practice Balance the following reaction under basic conditions:
I-(aq) + MnO4-(aq)  I2(s) + MnO2(s)