Download Chapter 4: Solution Chemistry and the Hydrosphere

Oxidation-Reduction (Redox)
Reactions
In oxidation-reduction (abbreviated as “redox”)
reactions, electrons are transferred from one reactant
to another.
Oxidation
I
Lose electrons
Reduction
I
Gain electrons
Redox Reactions
In the reaction between Na and Cl2:
Na
Na+
Na lost an electron, it
has been oxidized
Cl
Cl-
Cl gained an electron, it
has been reduced
electron
(e-)
2 Na (s) + Cl2 (g)  2 NaCl (s)
Redox Reactions
What about the reaction between Al and O2?
O
O2-
O gained two electrons,
it has been reduced
Al
Al3+
Al lost 3 electrons, it
has been oxidized
electrons
(e-)
Al (s)
(s) +
+ 3O2O(g)
 Al2O
4 Al
Al3 2(s)
O3 (s)
2 (g)  2
Oxidation Numbers
Oxidation Number (or Oxidation State): actual or
hypothetical charge of an atom in a compound if it
existed as a monatomic ion
Common Oxidation Numbers:
H+ = +1
Cl- = -1
O2- = -2
Al = 0
Na = 0
Na+ = +1
Oxidation numbers can also be assigned to atoms with
in a more complex molecule.
Assigning Oxidation Numbers
1. The oxidation number of an element in its natural form is 0. Examples:
the oxidation number is zero for each element in H2, O2, Cl2, P4, Na, etc.
2. The oxidation number of a monatomic ion is the charge on the ion.
Examples: Na3N, the ions are Na+ and N3–, so oxidation #’s: Na = +1 and
N = -3. In Al2O3, the ions are Al+3 and O2–, so oxidation #’s: Al = +3 and
O = -2
3. In a compound or polyatomic ion,
– Group I elements are always +1.
– Group II elements are always +2.
– Fluorine is always -1.
– Oxygen is usually -2 (except in the peroxide ion, O22–, when O is -1)
– Hydrogen is usually +1 (except when it is with a metal, like NaH or
CaH2, then it is -1)
4. In a neutral compound, the sum of all oxidation numbers must equal 0.
In a polyatomic ion, the sum of all oxidation numbers must equal the
charge.
Assigning Oxidation Numbers
Examples: Determine the oxidation number for each
element in the following:
a.
b.
c.
d.
e.
f.
CrO42–: Cr: ____, O: ____
H2SO4: H: ____, S: ____, O: ____
NO3-: N: ____, O: ____
CaCr2O7: Ca: ____, Cr: ____, O: ____
C2O42–: C: ____, O: ____
C3H8: C: ______________, H: ____
Redox Reactions
In a redox reaction:
– One reactant Loses Electrons/is Oxidized (LEO)
– Another reactant Gains Electrons/is Reduced (GER)
An easy way to remember is “LEO the lion goes GER!”
(Though I prefer OIL RIG, it’s your choice).
The element or reactant that is oxidized is the reducing
agent.
The element or reactant that is reduced is the oxidizing
agent.
Examples
a. Zn(s) +
AgNO3(aq)

b. Al(s) +
HCl(aq)

AlCl3(aq) +
O2(g)

CO2(g) +
c. C2H2(g) +
Zn(NO3)2(aq) +
Ag(s)
H2(g)
H2O(g)
Examples
d. Ca(s) +
e. H2O2(aq) +
H2O(l)

Mn(OH)2(aq)
Ca(OH)2(aq) + H2(g)

Mn(OH)3(aq)
Solution Concentration
solution: homogeneous mixture of substances present
as atoms, ions, and/or molecules
solute: component present in smaller amount
solvent: component present in greater amount
Note: Unless otherwise stated, the solvent for most
solutions considered in this class will almost always be
water!
Aqueous solutions are solutions in which water is the
solvent.
How do we measure
concentration?
• A concentrated solution has a large quantity of
solute present for a given amount of solution.
• A dilute solution has a small quantity of solute
present for a given amount of solution.
amount of solute
amount of solvent
The more solute in a given amount of solution  the
more concentrated the solution
Example: Explain the difference between the
density of pure ethanol and the concentration of an
ethanol solution.
SOLUTION CONCENTRATION =
How do we measure
concentration?
Concentration can be measured a number of ways:
• ppm (parts per million) – one part in a million parts
• ppb (parts per billion) – one part in a billion parts
• g/kg (grams per kilogram) – one gram solute per one
kilogram of solvent
The chemical standard most used is Molarity
Molarity =
moles of solute
liters of solution
units: M (molar = mol/L)
Solution Concentration
1. Find the molarity of a solution prepared by
dissolving 1.25 g of KOH in 150.0 mL of solution.
2. Find the molarity of a solution prepared by
dissolving 5.00 g of copper(II) sulfate in 250.0 mL of
solution
Ion Concentrations
• When an ionic compound is dissolved in water, the
concentration on the individual ions is based on their
molecular formula…
• For example:
– 1 M NaCl solution contains 1 M Na+ and 1 M Cl– 2 M NaCl solution contains 2 M Na+ and 2 M Cl– 1 M CaCl2 solutions contains 1 M Ca2+ and 2 M Cl– 2 M CaCl2 solutions contains 2 M Ca2+ and 4 M Cl-
Solution Concentration
3. Indicate the concentration of barium and chloride ions in a
1.00M barium chloride solution.
4. Indicate the molarity of each ion in the solutions indicated
below:
a. In a 0.125M Na2SO4(aq) solution
[Na+]=____________ and [SO42-]=____________.
b. In a 0.500M Fe(NO3)3(aq) solution
[Fe3+]=____________ and [NO3–]=___________.
c. In a 1.250M Al2(SO4)3(aq) solution
[Al+3]=____________ and [SO42-]=___________.
Solving Concentration
Problems
Keep in mind that if molarity and volume are both
given, you can calculate # of moles since:
volume  molarity = volume (in L)  moles of solute
liters of solution
so volume units will cancel  # of moles!
If you are given volume and molarity for a solution,
multiply them together to get # of moles!
Solving Concentration
Problems
Calculate the mass of NaCl needed to make 1.00 L of a
1.00 M solution.
Preparing Solutions
Examples
Calculate the mass of barium hydroxide required to make 250.0
mL of a 0.500M barium hydroxide solution.
What volume (in mL) of a 0.125M silver nitrate solution
contains 5.00 g of silver nitrate?
Examples
Calculate the molarity of hydroxide ion in a solution
prepared by diluting 50.0 mL of 1.50M potassium
hydroxide with 100.0 mL of 0.500M calcium hydroxide.