Exp. 16: Volumetric Analysis: Redox Titration

Normality = eq wt of solute

L solution

Acid/bases: #eq = # H+ or OH- ionized

Redox reactions – transfer of e-

reduction – oxidation reactions

Exp. 16 – video (time: 1 hr and 23:08 minutes)

Redox reaction

Equivalent wt - one equivalent of any oxidizing agent

reacts with one equivalent of any reducing agent.

This means #eq/mol is equal to the number of e-

transferred.

MnO4-(aq) + 8H+

(aq) + 5e- Mn2+(aq) + 4H2O(l) (net)

MnO4- : 5eq same for KMnO4

mol MnO4-

Fe2+(aq) Fe3+

(aq) + 1e- 1eq

mol Fe2+

N M or M N

N (eq) = M (mol) x #eq

L L mol

Note: N equal to or greater than M

0.1 M KMnO4 N? Goal: eq KMnO4

L soln

MnO4-(aq) + 8H+

(aq) + 5e- Mn2+(aq) + 4H2O(l)

Calc:

4

Solubility Rules for Ionic Compounds (Dissociates 100%)

1.) All compounds containing alkali metal cations and the ammonium ion are soluble.

2.) All compounds containing NO3-, ClO4

-, ClO3

-, and C2H3O2

- anions are

soluble.

3.) All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg2

2+.

4.) All sulfates are soluble except those containing Hg22+, Pb2+, Ba2+, Sr2+,

or Ca2+. Ag2SO4 is slightly soluble.

5.) All hydroxides are insoluble except compounds of the alkali metals and Ca2+, Sr2+, and Ba2+ are slightly soluble.

6.) All other compounds containing PO43-, S2-, CO3

2-, CrO42-, SO3

2- and most other anions are insoluble except those that also contain alkali metals or NH4

+.

Generally, compound dissolves > 0.10 M - soluble (aq)

< 0.01 M - insoluble (s)

in between - slightly soluble

(this class we will assume slightly soluble as soluble)

Hg2Cl2 (s) insoluble

KI (aq) soluble

Pb(NO3)2 (aq) soluble

5

Strong Acids (Ionizes 100%)

HCl, HBr, HI, HClO4, HNO3, H2SO4

Strong Bases (Dissociates 100%)

NaOH, KOH, LiOH, Ba(OH)2, Ca(OH)2,

Sr(OH)2

6

• A molecular/formula unit equation is one in which the reactants and products are written as if they were molecules/formula units, even though they may actually exist in solution as ions.

Calcium hydroxide + sodium carbonate

M.E.

Ca(OH)2

Ions in Aqueous Solution Molecular and Ionic Equations

+ Na2CO3 CaCO3 + NaOH 2 (aq)

strong base strong base soluble salt insoluble salt

(aq) (s) (aq)

s solid

l liquid

aq aqueous (acid/bases and soluble salts dissolve in water)

g gases

7

• An total ionic equation, however, represents strong electrolytes as

separate independent ions. This is a more accurate representation of the

way electrolytes behave in solution.

– A complete ionic equation is a chemical equation in which strong

electrolytes (such as soluble ionic compounds, strong acids/bases) are

written as separate ions in solution. (note: g, l, insoluble salts (s), weak

acid/bases do not break up into ions)

M.E.

Ca(OH)2 (aq) + Na2CO3 (aq) CaCO3 (s) + 2 NaOH (aq)

Total ionic

Ions in Aqueous Solution

Molecular and Ionic Equations

Ca2+ (aq) + 2OH- (aq)

strong base soluble salt insoluble salt strong base

+ 2Na+ (aq) + CO32-

(aq) CaCO3 (s) + 2Na+ (aq) + 2OH- (aq)

8

Net ionic equations.

– A net ionic equation is a chemical equation from

which the spectator ions have been removed.

– A spectator ion is an ion in an ionic equation that

does not take part in the reaction. M.E.

Ca(OH)2 (aq) + Na2CO3 (aq) CaCO3 (s) + 2 NaOH (aq)

Total Ionic Ca2+ (aq) + 2OH- (aq) + 2Na+ (aq) + CO3

2- (aq) CaCO3 (s) + 2Na+ (aq) + 2OH- (aq)

Net

Ca2+ (aq) + CO32-

(aq) CaCO3 (s)

9

Types of Chemical Reactions

• Oxidation-Reduction Reactions (Redox rxn)

– Oxidation-reduction reactions involve the

transfer of electrons from one species to another.

– Oxidation is defined as the loss of electrons.

– Reduction is defined as the gain of electrons.

– Oxidation and reduction always occur

simultaneously.

10

27.1 Reduction and Oxidation

Redox reactions – transfer of e-

reduction – oxidation reactions

Reduction – gain of e- / gain of H / lost of O

Fe3+ + 1e- Fe2+ (lower ox state)

note: must balance atoms and charges

11

Oxidation - loss of e- / loss of H / gain of O

Fe2+ Fe3+ + 1e- (higher ox state)

H2O + BrO3- BrO4

- + 2H+ + 2e-

(Br oxidized: charge 5+ 7+)

2H+ + 2e- H2 (H reduced: charge 1+ 0)

Oxidizing agent is species that undergoes reduction.

Reducing agent is species that undergoes oxidation.

Note: need both for reaction to happen; can’t have

something being reduced unless something else is being

oxidized.

Br + 3(-2) = -1

Br = -1 +6 = +5

Br + 4(-2) = -1

Br = -1 +8 = +7

12

27.3 Balancing Redox Reactions

- Must know charges (oxidation numbers) of species

including polyatomic ions

- Must know strong/weak acids and bases

- Must know the solubility rules

Oxidation Numbers – hypothetical charge assigned to the

atom in order to track electrons; determined by rules.

13

Rules to balance redox

1.) Convert to net ionic form if equation is originally in molecular form

(eliminate spectator ions).

2.) Write half reactions.

3.) Balance atoms using H+ / OH- / H2O as needed:

– acidic: H+ / H2O put water on side that needs O or H (comes from

solvent)

– basic: OH- / H2O put water on side that needs H but if there is no H

involved then put OH- on the side that needs the O in a 2:1 ratio

2OH- / H2O balance O with OH, double OH, add 1/2 water to

other side.

4.) Balance charges for half rxn using e-.

5.) Balance transfer/accept number of electron in whole reaction.

6.) Convert equation back to molecular form if necessary (re-apply

spectator ions).

Zn(s) + AgNO3(aq) Zn(NO3)2(aq) + Ag(s)

Total ionic:

Net ionic:

Zn(s) + Ag+(aq) + NO3-(aq) Zn2+(aq) + 2NO3

-(aq) + Ag(s)

Zn(s) + Ag+(aq) Zn2+(aq) + Ag(s)

14

Net: Zn(s) + Ag+(aq) Zn2+

(aq) + Ag(s)

Oxidation:

Reduction:

Balanced net:

Balanced eq:

Zn(s) Zn2+(aq) + 2e-

Ag+(aq) Ag(s) 1e- +

Zn(s) + 2 Ag+(aq) Zn2+(aq) + 2 Ag(s)

[ ] 2

Zn(s)

15

+ 2 AgNO3(aq) Zn(NO3)2(aq) + 2 Ag(s)

H+

Net: MnO4-(aq) + Fe2+

(aq) Mn2+(aq) + Fe3+

(aq)

Ox:

Red:

Balanced net:

Fe2+(aq) Fe3+(aq) + 1e- [ ] 5

MnO4-(aq) Mn2+(aq) + H2O(l) 4 8 H+(aq) + 5e- +

8 H+(aq) + MnO4-(aq) + 5 Fe2+(aq) Mn2+(aq) + 5 Fe3+(aq) + 4 H2O(l)

16

KMnO4(aq) + NaNO2(aq) + HCl(aq) NaNO3(aq) + MnCl2(aq) + KCl(aq) + H2O(l)

Net:

Ox:

Red:

Balanced net:

Balanced eq:

MnO4-(aq) Mn2+(aq) + NO2

-(aq) NO3

-(aq) + + H+(aq) + H2O(l)

NO2-(aq) NO3

-(aq)

MnO4-(aq) Mn2+(aq) + 4 H2O(l) 8 H+(aq) +

H2O(l) + + 2 H+(aq)

5 e- +

+ 2 e- [ ] 5

[ ] 2

2 MnO4-(aq) + 5 NO2

-(aq) + 16 H+(aq) + 5 H2O(l) 2Mn2+(aq) + 8 H2O(l) + 5 NO3-(aq) +10 H+(aq)

2 MnO4-(aq) + 5 NO2

-(aq) + 6 H+(aq) 2Mn2+(aq) + 3 H2O(l) + 5 NO3-(aq)

2 KMnO4(aq) + KCl 2 17

+ 5 NaNO2(aq) + 6 HCl(aq) 2MnCl2(aq) + 3 H2O(l) + 5 NaNO3(aq)

Net: OH-

CrI3 (s) + Cl2 (g) CrO42-

(aq) + IO4-(aq) + Cl-(aq)

Ox:

Red:

Balanced net:

CrI3(s) CrO42-(aq) + IO4

-(aq)

Cl2(g) Cl-(aq) 2

3 32 OH-(aq) + + 16 H2O(l)

2 e- +

+ 27 e- [ ] 2

[ ] 27

64 OH-(aq) + 2 CrI3(s) + 27 Cl2(g) 2 CrO42-(aq) + 6 IO4

-(aq) + 54 Cl-(aq) + 32 H2O(l)

18

Exp 16:

S2O32-

(aq) + I2 S4O62-

(aq) + I-(aq)

thiosulfate ion iodine

Ox:

Red:

Balanced net:

Outside exercise II page 199 – posted on my website

S2O32-(aq) S4O6

2-(aq)

I2(aq) I-(aq) 2

2 + 2 e-

2 e- +

2 S2O32-(aq) + I2(aq) S4O6

2-(aq) + 2 I-(aq)

S2O32-

2eq = 1 eq

2mol S2O32- mol S2O3

2-

I2 2eq

mol I2

Exp today

First: Standardize thiosulfate against 0.100 N I2 standard

solution.

Changes in sample preparation:

10 mL I2, 30 mL deionized H2O, 1 mL starch (20 drops)

Starch – indicator (add from beginning)

Starch + I2 gives blue color

At end pt (all I2 consumed), solution will be colorless

Since using normality can use

NiodineViodine = NthiosulfateV thiosulfate

minimum 3 runs ± 0.005 N (around ± 0.5 mL)

report

Avg N ± s N thiosulfate ion (S2O32-)

Convert average N to M

Second: Same exact procedure as standardization

except using unknown conc. of I2.

minimum 3 runs ± 0.005 N (around ± 0.5 mL)

report

Avg N ± s N iodine (I2) unknown

Convert average N to M

Amount of chemicals to obtain in small beaker per

group:

Na2S2O3.5H2O – 150 mL (source of thiosulfate ions)

0.100 N I2 standard solution – 50 mL

Unknown I2 solution – 45 mL