7.3 – BALANCING REDOX REACTIONS USING OXIDATION NUMBERS UNIT 7 – REDOX REACTIONS & ELECTROCHEMISTRY.
Oxidation-Reduction (Redox) Reactions
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Transcript of Oxidation-Reduction (Redox) Reactions
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Oxidation-Reduction (Redox) Reactions
In oxidation-reduction (abbreviated as “redox”) reactions, electrons are transferred from one reactant to another.
OxidationILose electrons
ReductionIGain electrons
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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)
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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)
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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.
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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.
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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
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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
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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)
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Examplesd. Ca(s) + H2O(l) Ca(OH)2(aq) + H2(g)
e. H2O2(aq) + Mn(OH)2(aq) Mn(OH)3(aq)
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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.
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• 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?
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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?
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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
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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-
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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-]=___________.
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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!
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Solving Concentration Problems
Calculate the mass of NaCl needed to make 1.00 L of a 1.00 M solution.
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Preparing Solutions
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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?
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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.