1. Chemistry 141
Friday, September 29, 2017
Lecture 11
Solution Chemistry

2. Objectives for today
Begin to gain chemical intuition about reactions in solution
Learn the terminology we use to describe solutions
Perform stoichiometric calculations using solution concentrations and volumes
Understand the factors that determine solubility and solution conductivity (electrolytes)
Predict whether a precipitate will form in a reaction

3. The balanced reaction for the combustion of methanol is: 
2 CH3OH + 3 O2 → 2 CO2 +  4 H2O 
If 0.098 mol of O2 reacts with excess CH3OH, what is the theoretical yield of water vapor from this reaction (in g) if the formula weight of water is g/mol?

4. If the actual (experimental) yield of water was 1
If the actual (experimental) yield of water was 1.9 grams, what was this reaction's percent yield?
 (Theoretical yield = 2.4 g)

5. Chemistry in solution
Many reactions take place in solution (i.e. homogeneous mixtures usually in the liquid phase), particularly aqueous solution (dissolved in water)
Solution terminology:
solution: mixture of solute and solvent
solute: substance dissolved in liquid
solvent: liquid itself (most abundant)
solvent does not have to be water (and a solution does not have to be liquid!)
Examples:
coffee, salt water, vinegar, ethanol/n-propanol, fat-soluble vitamins (such as Vitamin A) in body fat

6. Units of concentration
Molarity [M], moles solute per liter of solution
Molality: moles solute per kg of solvent (rarely used!)
Parts per million: (usually by mass, but can be volume)

7. Using concentrations
How would you prepare 2.00 L of a 0.50 M solution of NiCl2 from the salt NiCl26H2O?

8. Dilutions
How much water must be added to 15.5 ml of a 7.16 M solution of H3PO4 to reduce its concentration to 1.00 M? 1.
15.5 ml of M solution
2. dilute with water
3. final 1.00 M
solution

9. Solubility
What makes one substance soluble in another?

10. Water, the universal solvent

11. Dissolving a covalent compound
Covalent (molecular) compounds:
Consist of neutral atoms (usually non-metals) bonded together in discrete molecules
Dissolve in ‘like’ solvents (polar ethanol molecule in polar water)
Dissociate very little

12. Dissolving an ionic compound (salt)
Consist of a lattice of + (metal) and – (non-metal) ions
Dissociate in water into ions
Conduct electricity (electrolyte)
Cl–
Na+
NaCl (s) Na+ (aq) + Cl– (aq)

13. Electrolytes
A strong electrolyte dissociates completely when dissolved in water; solution conducts electricity (e.g. soluble salts, strong acids and bases)
A weak electrolyte only dissociates partially when dissolved in water; conducts weakly (e.g., weak acids and bases)
A nonelectrolyte does NOT dissociate in water

14. Solubility rules Compounds containing the following ions are soluble:
- Li+, Na+, K+, Cs+, and NH4+
- nitrates (NO3–), chlorates (ClO3–), and acetates (C2H3O2–)
- chlorides (Cl–), bromides (Br–), and iodides (I–)
(except those of Ag+, Hg22+, Tl+, and Pb2+)
- sulfates (SO42–)
(except those of Ca2+, Ba2+, Pb2+, Ag+, and Hg22+)
Compounds containing the following ions are insoluble:
- sulfides (S2–)
(except those of alkali metals, NH4+, Ca2+, and Ba2+)
- carbonates (CO32–), phosphates (PO43–), and sulfites (SO32–)
(except those of alkali metals and NH4+)
- hydroxides (OH–)

15. Solubility trends Most compounds containing +1 and –1 ions are soluble
Most compounds containing +2 and –2 or more highly charged ions are insoluble
Why?
The coulombic (charge) forces holding these ions together is greater +1 –1 +2 –2
more force
less force

16. Compound types and solubility
Ionic (salt)
Covalent (molecule)
- Made up of + and – ions
- Made up of neutral atoms
- Consist of metal + nonmetal
- Consist of nonmetal + nonmetal
- Exist as a continuous lattice of alternating + and – ions
- Exist as discrete molecules
- Dissociate in water
- Dissociate very little in water
Solutions conduct electricity
(electrolytes)
Solutions do not conduct electricity
(non-electrolytes)
Ionic
Polar covalent
Non-polar covalent
(+2)(-2)
(+1)(-1)
ethanol
hexane
dissolves in water? NO
YES
increasing charge separation
(use solubility rules for each case)

17. Precipitation
What happens when aqueous solutions of KCl and AgNO3 are mixed?