Chapter 4 Solutions and Chemical Reactions
I.
Water
A.
Importance
1. Life (as we know it) depends on water
2. Human civilization requires water for many purposes
3. Many important chemical reactions occur in Aqueous Solutions, where
other compounds are dissolved in water
B. The nature of water
1. Bent shape and unequal sharing of electrons makes water polar
2. This aids water in dissolving ionic compounds (cations and anions)
3. Water hydrates the ions by interacting with its oppositely charged ends
4.
5.
The ionic substance breaks up into independent cations and anions
Nonionic compounds can also dissolve in water if they are polar
Ethanol
6.
II.
Nonpolar substances generally don’t dissolve in water: grease, oils, skin
Electrolytes
A.
Solutions
1. A solution is a homogeneous mixture the same throughout
2. We can vary the composition by adding more or less of the components
3. Solvent = usually a liquid; the most abundant component of a solution
4. Solute = the lesser abundant component(s) of a solution
NaCl(s) -----> Na+(aq) + Cl-(aq)
B. Solutions and Electrical Conductance
1. A substance allowing current to flow through it is electrically conductive
2. Pure water does not conduct electricity
3. Different solutes dissolved in water help it to be conductive
a. Strong electrolyte = completely ionized; strongly conductive solution
b. Weak electrolyte = partially ionized; somewhat conductive solution
c. Nonelectrolyte = not ionized; nonconductive solution
Arrhenius (1859-1927) found
that the more ions present, the
better the conductivity
C. Strong Electrolytes
1. Completely ionized when dissolved in water
2. Many salts (ionic compounds) are strong electrolytes
3. Strong Acids are strong electrolytes
a. Acid = substance that produces H+ when dissolved in water
b. Strong Acids completely ionize in solution
i. Hydrochloric Acid HCl(g) -------> H+(aq) + Cl-(aq)
ii. Nitric Acid HNO3(g) -------> H+(aq) + NO3-(aq)
iii. Sulfuric Acid H2SO4(l) -------> H+(aq) + HSO4-(aq)
4. Strong Bases are strong electrolytes
a. Base = substance that produces OH- when dissolved in water
b. Strong bases completely ionize in solution
NaOH(s) -------> Na+(aq) + OH-(aq)
KOH(s) -------> K+(aq) + OH-(aq)
D. Weak Electrolytes
1. Only partially ionized when dissolved in water
2. Weak Acids are weak electrolytes
a. Weak acid only produces a few H+ ions
b. Acetic acid is a weak acid
c. HC2H3O2(aq) -------> H+(aq) + -C2H3O2(aq)
d. Only 1 molecule in a 100 dissociates
3. Weak Bases are weak electrolytes
a. Weak base produces only a few OH- ions
b. Ammonia is a weak base
c. NH3(aq) + H2O(l) -----> NH4+(aq) + OH-(aq)
d. Only 1 molecule in 100 reacts
E. Nonelectrolytes
1. Does not ionize when dissolve in water
2. Sugar is a nonelectrolyte
3. C12H22O11(s) -------> C12H22O11(aq)
III. Solution Concentration
A.
The Stoichiometry of Chemical Reactions
1. We must know what the reactants and products are
2. We must know the amounts of the reactants and products
3. How do we describe the amounts in a solution?
B.
Molarity
1. Unit for the concentration of a solute in a solution
2. M = moles solute/liters of solution
3. 1.0 M NaCl = 1 mole of NaCl dissolved in 1 L of solution
a. Any volume having the same concentration is also 1.0 M NaCl
b. 500 ml (0.500 L) of 1.0 M NaCl would contain 0.5 mol NaCl
0.5mol 1mol

 1M
0.5L
L
4. Example: Calculate M of 11.5 g NaOH in 1.5 L of total solution.
 1mol  1  0.288mol  0.192mol 

 
  0.192M
11.5g 
 


1.5L
L


 40.00 g  1.5L 
5. Example: M = ? for 1.56 g HCl in a total of 26.8 ml of solution?
 1mol 
1  1000mL   1.60mol 




  
  1.60M
1.56g 

L
L
 

 36.46 g  26.8mL 
6. Molarity descriptions of a solution reflect composition before dissolution
a. 1.0 M NaCl actually contains no NaCl
b. 1.0 M NaCl is 1.0 M in Na+ and 1.0 M in Clc. 1.0 M CaCl2 is 1.0 M in Ca2+ and 2.0 M in Cld. CaCl2(s) -------> Ca2+(aq) + 2Cl-(aq)
7. Example: Give the concentration of each ion
a. 0.5 M Co(NO3)2 = 0.5M in Co2+ and 1.0 M in NO3b. Co(NO3)2 (s) -------> Co2+(aq) + 2NO3-(aq)
c. 1M Fe(ClO4)3 = 1M Fe3+ and 3M ClO48. Example: ??? moles of Cl- in 1.75L of 0.001 M ZnCl2
 0.002mol Cl 1.75L
L


  0.0035mol Cl 

9. Example: What volume of 0.14 M NaCl contains 1.0mg NaCl?

1g  1 mol NaCl 

  1.7 x 10 -5 mol NaCl


 1000mg  58.45g NaCl 
1.0mg NaCl




1L

  1.2 x 10 -4
1.7 x 10 mol NaCl 
 0.14mol NaCl 
-5
 1000mL 
  0.12ml
L
L


10. Standard Solution = concentration is accurately known
a. Accurate masses come from an analytical balance (0.4563g)
b. Accurate volumes are obtained using a Volumetric Flask
c. Example: How much K2Cr2O7 needed for 1.00 L of 0.200 M?
 0.200mol  294.20 g

1.00L
L

 mol

  58.8g K 2 Cr2 O 7

C. Dilution
1. Chemicals are often purchased or prepared as concentrated stock solutions
2. Dilution = adding water to stock solution to make a less concentrated one
3. M1V1 = M2V2 is a useful equation to calculate dilutions
 mol 
 mol 
 L 1  mol  
 L 2  M 2V2
M 1V1  
 L 1
 L 2
4. Example: ??? volume of 16 M H2SO4 is needed for 1.5 L 0.10M H2SO4
 16mol 
 0.10mol 
XL  
1.5L 
M 1V1  M 2V2  
L
 L 


XL   0.10mol 1.5L 1L   9.4 x 10-3 L 1000ml   9.4ml
L


 16mol 
 L 
IV. Precipitation Reactions
A. Definitions
1. When two solutions are mixed and a solid forms
2. Precipitate = solid that forms from a precipitation reaction
3. K2CrO4(aq) + Ba(NO3)2(aq) = 2K+(aq) + CrO42-(aq) + Ba2+(aq) +2NO3(aq)
a. K2CrO4 and Ba(NO3)2 are both soluble (all dissolve in water)
b. A yellow precipitate forms when these solutions are mixed
+
=
Precipitate
Spectator Ions
c. K2CrO4(aq) + Ba(NO3)2(aq) -----> BaCrO4(s) + 2KNO3(aq)
4. AgNO3(aq) + KCl(aq) ------> AgCl(s) + KNO3(aq)
B. Solubility Rules
Example: predict what will happen when you mix:
a. KNO3(aq) + BaCl2(aq) ------>
b. Na2SO4(aq) + Pb(NO3)2(aq) ------>
c. 3KOH(aq) + Fe(NO3)2(aq) ------>
C. Describing Reactions in Solution
1. Molecular Equation shows what compounds the ions came from
a. Does not give clear picture of what happens in solution
b. K2CrO4(aq) + Ba(NO3)2(aq) -----> BaCrO4(s) + 2KNO3(aq)
2. Complete Ionic Equation represents the form of the ions in solution
a. All strong electrolytes are represented as their ions
2K+(aq) + CrO42-(aq) + Ba2+(aq) +2NO3-(aq) ----> BaCrO4(s) + 2K+(aq) + 2NO3-(aq)
3. Net Ionic Equation shows only the ions participating in the reaction
a. The K+ and NO3- ions occur on both sides of the complete ionic eqn.
b. These spectator ions can be cancelled out of each side (algebra)
c. Ba2+(aq) + CrO42-(aq) -------> BaCrO4(s)
4. Example
a. 3KOH(aq) + Fe(NO3)3(aq) -----> Fe(OH)3(s) + 3KNO3(aq)
b. 3K+(aq) + 3OH-(aq) + Fe3+(aq) + 3NO3-(aq) ----> Fe(OH)3(s) + 3K+(aq) + 3NO3-(aq)
c. Fe3+(aq) + 3OH-(aq) -------> Fe(OH)3(s)