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13 Dec 2019

Calculate the equilibrium concentration of FeSCN2+ that is expected in each of the solutions listed in Part A of the experimental procedure (blank nd solutions 1-5). Show all work and units.

Standard Solution

Blank A 0.2 M Fe(NO3)3 (in 0.1 M HNO3) 5x10-4 M KSCN (in 0.1 M HNO3) 0.1 M HNO3

1 30 drops 0 drops 45 drops

2 30 drops 5 drops 40 drops

3 30 drops 10 drops 35 drops

4 30 drops 30 drops 15 drops

5 30 drops 45 drops 0 drops

I know you have to use an ICE table, I'm just not exactly sure how to set it up. The equation I was given was Fe3+ + SCN- = FeSCN2+ and I'm not sure if this is used for the ice table. HELP please :)

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experment1 Electron Transfer (Redox) Reactions

1. Comparison of Br2/Brâ and Fe3+/Fe2+ Couples

Oxidant-reductant pairs from these couples will he tested, i.e. Br2 with the reductant Fe2+ and

Brâ with the oxidant Fe3+. A reaction will occur in only one of these cases because of the

difference in electrode potential between the two couples.

To test whether reaction has occurred, SCNâ will be used as a probe for Fe3+, with which it

gives a red complex. Note that the air rapidly oxidises Fe2+, so that a faint pink colour is often

obtained when SCNâ is added to Fe2+.

Solutions containing Fe2+ or Fe3+ tend to decompose (although in different ways); both are

stabilised by H+. The iron(II) solution provided is freshly prepared 0.1 M (NH4)2Fe(SO4)2 / 2

M H2SO4.

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of 0.1 M KBr and

stir. Place 1 drop of the mixture on a white tile, and test for Fe3+ by adding 1 drop 0.1 M

KSCN. A strong red colour indicates Fe3+.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M KBr.

Repeat using 5 drops of the 0.1 M iron(II) solution and adding 10 drops of 0.1 M Br2.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M Br2.

2.Use of the Br2/Brâ and Fe3+/Fe2+ Couples in an Electrochemical Cell

You are supplied with a pair of carbon rods joined by copper wire. Place two 100 mL beakers

side by side.

In one put 50 mL of the iron(II) solution and 1 mL of 0.1 M KSCN solution. In the other

beaker put 25 mL of 0.2 M Br2 solution and 25 mL of 4 M HCl.

Correct any difference in level between the solutions in the two beakers by adding water to

one. Fold an 11 cm filter paper into a strip about 1 cm wide. Bend the strip in the middle,

moisten with a little 4 M HCl, then insert the strip so that it links the two solutions. Insert one

carbon rod into each solution, then allow to stand undisturbed for 30 minutes while

proceeding with Experiment 3. Carefully observe the apparatus at intervals and record any

change that occurs.

3. Comparison of the Fe3+/Fe2+ and Sn4+/Sn2+ Couples

In this experiment the additional sources of ions are:

tin(IV) ion 0.1 M SnCl4/1 M HCl

tin(II) ion 0.1 M SnCl2/1 M HCl

(Both Sn2+ and Sn4+ solutions are stabilised by H+.)

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of the appropriate

0.1 M tin solution and stir. Place 1 drop of the mixture on a white tile, and test for Fe3+.

Record the results.

Similarly, test the 0.1 M iron(III) solution with an excess of the appropriate 0.1 M tin

solution.

You have now enough information to state whether the standard electrode potential of the

Fe3+/Fe2+ couple is more positive or less positive than that of the Sn4+/Sn2+ couple.

please answer following questions

1.Comparison of Br₂/Br^â and Fe³⁺/Fe²⁺ couples

Complete the following table.

Solutions	Observation	Fe³⁺ present (Yes/No)	Did a reaction occur?
Fe(II) and KBr
Fe(III) and KBr
Fe(II) and Br₂
Fe(III) and Br₂

2.Br₂/Br^â and Fe³⁺/Fe²⁺ couples in an Electrochemical Cell

a)What observations did you make over 3 minutes? Did Br₂ or Fe²⁺ react? Write the ion/electron half- equations for this reaction and the complementary reduction. Write the overall redox reaction.

b)Draw/scan/insert a circuit diagram showing the half-reactions that occur in each beaker, the direction of electron flow and the direction of flow of ions.

3.Fe³⁺/Fe²⁺ and Sn⁴⁺/Sn²⁺ couples

Complete the following table.

Solutions	Observation	Fe³⁺ present (Yes/No)	Did a reaction occur?
Fe(II) and Sn(IV)
Fe(II) and Sn(II)

	Trial 1	Trial 2
Volume of Unknown Weak Acid	4 mL	4 mL
Total Number of Drops Required to Reach Equivalence Point	5	5
Average Number of Drops	5

Drops NaOH Added	pH Value Trial 1	pH Value Trial 2	pH Value (Average)
Half-Equivalent Point = (46 Drops)	2.5	2.5	2.5
Equivalent Point = (93 Drops)	5	5	5
0	3.5	3.5	3.5
10	4.5	4.5	4.5
20	5.5	5.5	5.5
30	7.0	7.0	7.0
40	8.0	8.0	8.0
50	10.0	10.0	10.0
60	10.0	10.0	10.0
70	10.0	10.0	10.0
80	10.0	10.0	10.0
90	10.0	10.0	10.0
100	10.0	10.0	10.0
110	10.0	10.0	10.0
120	10.0	10.0	10.0

pK_a of Unknown Weak Acid:	2.5
K_a of Unknown Weak Acid:	3.16 x 10 ^-3
Unknown Weak Acid Identity:	Citric Acid
% Error pK_a:	10.39%
% Error K_a:	4.9%

A. pH indicator strips work by changing color in the presence of solutions with varied pH values. Thinking about your procedure steps and results in Part 1, why do you think the pH of the unknown weak acid was not determined with pH indicator strips until Part 2.

B. Discuss possible causes of error in the experimental procedure. Why do you think there was a percent of error in the pK_a and K_a values of the unknown weak acid, in comparison to the values presented in Table 2?

C. Do you think using a pH meter instead of pH indicator strips would have created a larger or smaller percent error? Explain your answer.

D. Why was phenolphthalein a good indicator to use for determining the equivalence point between the unknown weak acid and strong base?

1) You prepare a mixture to use to create a colorimetry calibration curve for FeSCN²⁺by mixing these things in a 25.00 mL volumetric flask and then diluting to the volume of the flask.
10.00 mL 0.100 M KSCN
5.00 mL 2 M HNO3
4.00 mL 0.00100 M Fe(NO₃)₃

What is the concentration of FeSCN²⁺ in this mixture?

2) You prepare a mixture to use to measure the equilibrium constant of the reaction of Fe³⁺ and SCN^- ions. You mix these things in a 25.00 mL volumetric flask and then dilute to the volume of the flask.
3.00 mL 0.00400 M KSCN
5.00 mL 2 M HNO3
10.00 mL 0.00400 M Fe(NO₃)₃

What is the concentration of Fe³⁺ ions in this mixture initially

3) You prepare a mixture to use to measure the equilibrium constant of the reaction of Fe³⁺ and SCN^- ions. You mix these things in a 25.00 mL volumetric flask and then diluting to the volume of the flask.
3.00 mL 0.00400 M KSCN
5.00 mL 2 M HNO3
10.00 mL 0.00400 M Fe(NO₃)₃

What is the concentration of SCN^- ions in this mixture initially?

4)You prepare a mixture to use to measure the equilibrium constant of the reaction of Fe³⁺ + SCN^- --> FeSCN²⁺

You mix these things in a 25.00 mL volumetric flask and then diluting to the volume of the flask.
3.00 mL 0.00400 M KSCN
5.00 mL 2 M HNO3
10.00 mL 0.00400 M Fe(NO₃)₃

You use colorimetry to determine that, at equilibrium, the FeSCN2+ concentration is 2.4 x 10^-5 M.

Create an I-C-E table and find the K for this particular reaction under these particular conditions

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