oxidation-reduction (redox) reactions

Oxidation-Reduction (Redox) Reactions

In oxidation-reduction (abbreviated as “redox”) reactions, electrons are transferred from one reactant to another.

OxidationILose electrons

ReductionIGain electrons

Redox ReactionsIn the reaction between Na and Cl2:

Na

Cl Cl-

Na+

electron (e-)

Na lost an electron, it has been oxidized

Cl gained an electron, it has been reduced

2 Na (s) + Cl2 (g) 2 NaCl (s)

What about the reaction between Al and O2?

Redox Reactions

O

Al

electrons (e-)

Al lost 3 electrons, it has been oxidized

O gained two electrons, it has been reduced

O2-

Al3+

Al (s) + O2 (g) Al2O3 (s)4 Al (s) + 3 O2 (g) 2 Al2O3 (s)

Oxidation NumbersOxidation 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- = -1O2- = -2 Al = 0Na = 0 Na+ = +1

Oxidation numbers can also be assigned to atoms with in a more complex molecule.

Assigning Oxidation Numbers1. 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, O2

2–, 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. CrO4

2–: Cr: ____, O: ____b. H2SO4: H: ____, S: ____, O: ____c. NO3

-: N: ____, O: ____d. CaCr2O7: Ca: ____, Cr: ____, O: ____e. C2O4

2–: C: ____, O: ____f. C3H8: C: ______________, H: ____ C

CC

H

HH

H H

H

HH

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.

Redox Reactions

Examplesa. Zn(s) + AgNO3(aq) Zn(NO3)2(aq) +

Ag(s)

b. Al(s) + HCl(aq) AlCl3(aq) + H2(g)

c. C2H2(g) + O2(g) CO2(g) + H2O(g)

Examplesd. Ca(s) + H2O(l) Ca(OH)2(aq) + H2(g)

e. H2O2(aq) + Mn(OH)2(aq) Mn(OH)3(aq)

Solution Concentrationsolution: 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.

• 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.

SOLUTION CONCENTRATION = 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.

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 solventThe chemical standard most used is Molarity

Molarity =

units: M (molar = mol/L)

How do we measure concentration?

Solution Concentration1. 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 Concentration3. 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 [SO4

2-]=____________.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 [SO4

2-]=___________.

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

ExamplesCalculate 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?

ExamplesCalculate 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.