Determine the pH of a solution of HCIO₂ by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-3 before submitting your answer. 1 2 3 NEXT > A 0.020 M aqueous solution of HCIO2 is prepared. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products.. Initial (M) Change (M) Equilibrium (M) -2x 1.2 x 10-²-2x 0 0.020 + x HCIO₂(aq) 0.020 0.020 - x + H₂O(l) 1.2 x 10-² 0.020 + 2x 0.010 0.020 - 2x H3O+ (aq) +X 1.2 x 10-2 + x + CIO₂ (aq) 1.2 x 10-2-x RESET +2x 1.2 x 10-² + 2x

Determine the pH of a solution of HCIO₂ by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-3 before submitting your answer. 1 2 3 NEXT > A 0.020 M aqueous solution of HCIO2 is prepared. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products.. Initial (M) Change (M) Equilibrium (M) -2x 1.2 x 10-²-2x 0 0.020 + x HCIO₂(aq) 0.020 0.020 - x + H₂O(l) 1.2 x 10-² 0.020 + 2x 0.010 0.020 - 2x H3O+ (aq) +X 1.2 x 10-2 + x + CIO₂ (aq) 1.2 x 10-2-x RESET +2x 1.2 x 10-² + 2x

Chemistry for Engineering Students

4th Edition

ISBN:9781337398909

Author:Lawrence S. Brown, Tom Holme

Publisher:Lawrence S. Brown, Tom Holme

Chapter12: Chemical Equilibrium

Section: Chapter Questions

Problem 12.11PAE

See similar textbooks

Similar questions

Use Le Chátelier's principle to explain what happens to the equilibrium H2O(l)H+(aq)+OH(aq) when a few drops of HCl are added to pure water.
7. Describe a buffered solution. Give three examples of buffered solutions. For each of your examples, write equations and explain how the components of the buffered solution consume added strong acids or bases. Why is buffering of solutions in biological systems so important?
Hydrogen iodide gas decomposes to hydrogen gas and iodine gas: 2HI(g)H2(g)+I2(g)To determine the equilibrium constant of the system, identical one-liter glass bulbs are filled with 3.20 g of HI and maintained at a certain temperature. Each bulb is periodically opened and analyzed for iodine formation by titration with sodium thiosulfate, Na2S2O3. I2(aq)+2S2O32(aq)S4O62(aq)+2 I(aq)It is determined that when equilibrium is reached, 37.0 mL of 0.200 M Na2S2O3 is required to titrate the iodine. What is K at the temperature of the experiment?
A small quantity of a soluble salt is placed in water. Equilibrium between dissolved and undissolved salt may or may not be attained. Explain.
Consider citric acid, H3C6H5O7, added to many soft drinks. The equilibrium constants for its step-wise ionization areKa1=7.5104 ,Ka2=1.7105, andKa3=4.0107. Write the overall net ionic equation and calculate K for the complete ionization of citric acid.
Determine the pOH of a weak acid (HA) solution by constructing an ICE table and using this information to determine the pOH of the solution. Complete Parts 1-2 before submitting your answer. 1 A 0.200 M aqueous solution of HA is 3.0% dissociated at equilibrium. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products. Initial (M) Change (M) Equilibrium (M) HA(aq) + H₂O(l) 2 H3O+ (aq) 2 NEXT > + A-(aq)
ermine the pH of a solution by constructing a BCA table, constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-4 before submitting your answer. 2 NEXT > Two solutions are mixed: 40.0 mL of 0.500 M NH, and 25.0 mL of 0.300 M HCI. Fill in the table with the appropriate value for each involved species to determine the moles of reactant and product after the reaction of the acid and base.. NH3(aq) NH, (aq) (be),H Before (mol) Change (mol) After (mol) 5 RESET 40.0 0.500 25.0 0.300 0.300 -x 0.0075 -0.0075 0.0100 -0.0100 0.0125 -0.0125 0.0200 -0.0200 Based on the result of the acid-base reaction (Part 1), set up the ICE table in order to determine the unknown concentrations of all reactants and products. NH3(aq) H,O(1) OH (aq) NH, (aq) Initial (M) Change (M) Equilibrium (M) 5 RESET 0.500 0.300 0.192 0.313 0.115 0.500 -x 0.300 -x 0.192 +* 0.192 -x 0.500 + x 0.300 +x 0.313 + x 0.313 -x 0.115 +x 0.115-x
Determine the pH of a buffer solution by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-3 before submitting your answer. 3 > A solution is prepared with 0.55 M HNO2 and 0.75 M KNO2. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products. Initial (M) Change (M) Equilibrium (M) -2x 0.75 -2x 0 1 HNO₂(aq) 0.55 + x 0.55 -X 0.55-x 0.55 0.55 - x 0.75 2 H₂O(l) III 0.55 + 2x 6.8 x 104 0.55 - 2x 1L +x H3O+(aq) 0.75 + x +x -X 0.75-x + NO₂ (aq) RESET +2x NEXT 0.75 + 2x +x
Determine the pH of a solution by constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-3 before submitting your answer. NEXT > A solution is prepared with 0.075 M (CH3)3N and 0.10 M (CH3)3 NHCI. Fill in the ICE table with the appropriate value for each involved species to determine concentrations of all reactants and products. Initial (M) Change (M) Equilibrium (M) -2x 0.10 - 2x 1 (CH3)3N(aq) 0 000 0.075 + x 0.075 0.075-x + 0.10 2 H₂O(l) 0.075 + 2x 5.2 x 104 0.075 - 2x 1L +x 3 OH (aq) 00 0.10 + x 0.10-x + (CH3)3NH+ (aq) RESET +2x 0.10 + 2x
Explain why we can simplify the calculations for K when the value of K is very small (less 4.1 than -10*). 4.2 HCN will react very slightly with water with a K = 6.2 x 10-10. HCN(aq) + H,O() H,O* (aq) and CN" (aq) A chemist starts with a 0.50 M HCN solution. What is the equilibrium concentration of H;O*?. Then use the concentration of H3O* to calculate the pH of the solution with the equation pH = -log[H3O*] (you can do this!!!). Put your final answers in the box. If you need more space to show your work, you can use the backside of this page. [H;O'] at equilibrium pH = -log(H,O] pH = % of HCN that reacted
Determine the Ka for an acid by constructing an ICE table and using this information to construct and solve the equilibrium constant expression. Complete Parts 1-2 before submitting your answer. Initial (M) Change (M) Equilibrium (M) 0.022 The pH for 0.0715 M solution of CCI3CO₂H is 1.40. Fill in the ICE table with the appropriate value for each involved species to determine the unknown concentrations of all reactants and products. 0 1 CCI CO₂H(aq) + 0.0715 1.40 2 -1.40 H₂O(1) 2 H3O+ (aq) + CCI3CO₂ (aq) NEXT > 0.040 -0.040 RESET 0.032 Determine the Ka for an acid by constructing an ICE table and using this information to construct and solve the equilibrium constant expression. Complete Parts 1-2 before submitting your answer. Initial (M) Change (M) Equilibrium (M) 0.000 The pH for 0.0715 M solution of CCI.COH is 1.40. Fill in the ICE table with the appropriate value for each involved species to determine the unknown concentrations of all reactants and products. 0 1 CCI.COH(aq) + H₂O(l)…
Determine the Kb for a base by constructing an ICE table and using this information to construct and solve the equilibrium constant expression. Complete Parts 1-2 before submitting your answer. Initial (M) Change (M) Equilibrium (M) NEXT The pH for a 0.0160 M solution of CH3C6H4NH₂ is 8.600. Fill in the ICE table with the appropriate value for each involved species to determine the unknown concentrations of all reactants and products. -2.50 × 10-9 1 CH3C6H4NH₂(aq) + 0 3.98 x 10-6 0.0160 -3.98 x 10-6 8.60 5.4 H₂O(1) -8.60 -5.4 2 = OH-(aq) 0.93 -0.93 + CH3C6H4NH3 + (aq) RESET 2.50 × 10-9