please give answer and explain the process thank you c. The Pourbaix diagram for iron is shown in Figure 4.1.+1.0+0.52 0.03-0.5-1.0keyFe³+2soluble or not.Fe2+soluble4H₂O/H₂Fe6insolubleH₂O/O2pH008Figure 4.1 Simplified Pourbaix diagram (versus pH) for iron.Fe(OH)310Fe(OH)212VDescribe all of the changes which occur to the iron-containing species as the potentialis decreased gradually from +1.0V to -1.0 V at pH 2. Your answer should include theoxidation state of each of the Fe species and should indicate whether the substance isNERFAREN

The Pourbaix diagram shows the stability of different oxidation states of an element with pH. Based…

c. The Pourbaix diagram for iron is shown in Figure 4.1. +1.0- +0.5- 2 0.0 -0.5- -1.05 ⁹0 key Fe3+ 2 Fe2+ soluble H₂O/H2 Fe 6 insoluble H₂0/02 pH 8 Figure 4.1 Simplified Pourbaix diagram (versus pH) for iron. Fe(OH)3 10 Fe(OH)2 N 12 14 Describe all of the changes which occur to the iron-containing species as the potential is decreased gradually from +1.0 V to -1.0 V at pH 2. Your answer should include the oxidation state of each of the Fe species and should indicate whether the substance is soluble or not.

c. The Pourbaix diagram for iron is shown in Figure 4.1. +1.0- +0.5- 2 0.0 -0.5- -1.05 ⁹0 key Fe3+ 2 Fe2+ soluble H₂O/H2 Fe 6 insoluble H₂0/02 pH 8 Figure 4.1 Simplified Pourbaix diagram (versus pH) for iron. Fe(OH)3 10 Fe(OH)2 N 12 14 Describe all of the changes which occur to the iron-containing species as the potential is decreased gradually from +1.0 V to -1.0 V at pH 2. Your answer should include the oxidation state of each of the Fe species and should indicate whether the substance is soluble or not.

Introductory Chemistry: A Foundation

9th Edition

ISBN:9781337399425

Author:Steven S. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Donald J. DeCoste

Chapter16: Acids And Bases

Section: Chapter Questions

Problem 10ALQ

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A monoprotic organic acid that has a molar mass of 176.1 g/mol is synthesized. Unfortunately, the acid produced is not completely pure. In addition, it is not soluble in water. A chemist weighs a 1.8451-g sample of the impure acid and adds it to 100.0 mL of 0.1050 M NaOH. The acid is soluble in the NaOH solution and reacts to consume most of the NaOH. The amount of excess NaOH is determined by titration: It takes 3.28 mL of 0.0970 M HCl to neutralize the excess NaOH. What is the purity of the original acid, in percent?
Isocyanic acid (HNCO) can be prepared by heating sodium cyanate in the presence of solid oxalic acid according to the equation 2NaOCN(s)+H2C2O4(s)2HNCO(l)+Na2C2O4(s) Upon isolating pure HNCO(l), an aqueous solution of HNCO can be prepared by dissolving the liquid HNCO in water. What is the pH of a l00.-mL solution of HNCO prepared from the reaction of 10.0 g each of NaOCN and H2C2O4, assuming all of the HNCO produced is dissolved in solution? (Ka of HNCO = 1.2 l04.)
Consider 1000. mL of a 1.00 104-M solution of a certain acid HA that has a Ka value equal to 1.00 104. How much water was added or removed (by evaporation) so that a solution remains in which 25.0% of HA is dissociated at equilibrium? Assume that HA is nonvolatile.
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A friend asks the following: Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like NaOH is added, the HA reacts with the OH to form A. Thus the amount of acid (HA) is decreased, and the amount of base (A) is increased. Analogously, adding HCI to the buffered solution forms more of the acid (HA) by reacting with the base (A). Thus how can we claim that a buffered solution resists changes in the pH of the solution? How would you explain buffering to this friend?
A buffer is prepared in which the ratio [ H2PO4 ]/[ HPO42 ]is 3.0. (a) What is the pH of this buffer? (b) Enough strong acid is added to convert 15% of HPO42- to H2PO4-. What is the pH of the resulting solution? (c) Enough strong base is added to make the pH 7.00. What is the ratio of [H2PO4-] to [HPO42-] at this point?
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The titration of Na2CO3 with HCl bas the following qualitative profile: a. Identify the major species in solution at points AF. b. Calculate the pH at the halfway points to equivalence, B and D. (Hint: Refer to Exercise 113.)