3. A typical vitamin C tablet contains pure ascorbic acid H2C6H606 (a weak, diprotic acid).If 2 tablets are dissolved in about a cup of water, the resulting concentration of theascorbic acid is 0.0115 M.a. Give the equation for the dissociation for the first ionization (loss of the first acidicarlt el todW baiH broosz arit otH). This ionization has a Ka of 8.00 x 10°.b. What is the equilibrium concentration for all species from the first ionization? c. Give the equation for the dissociation for the second ionization (loss of the secondacidic H from the CB of step a).(blsp sire Non)0 Mboon eniatrI4 5 497d. What is the equilibrium concentration for all species in the second ionization if the Kafor the second H is 1.6 x 10-122Drertamine eT srir men es llo nol nor onon murndilups orit ai toriW de. Combining your answers from b and d, what is the total H3O* concentration, pH, pOH,and % ionization?

a.) To write the dissociation equation , we just ionise the given compound into its ions . b.) To…

3. A typical vitamin C tablet contains pure ascorbic acid H2C6H6O6 (a weak, diprotic acid). If 2 tablets are dissolved in about a cup of water, the resulting concentration of the ascorbic acid is 0.0115 M. a. Give the equation for the dissociation for the first ionization (loss of the first acidic H). This ionization has a Ka of 8.00 x 10-5. b. What is the equilibrium concentration for all species from the first ionization?

3. A typical vitamin C tablet contains pure ascorbic acid H2C6H6O6 (a weak, diprotic acid). If 2 tablets are dissolved in about a cup of water, the resulting concentration of the ascorbic acid is 0.0115 M. a. Give the equation for the dissociation for the first ionization (loss of the first acidic H). This ionization has a Ka of 8.00 x 10-5. b. What is the equilibrium concentration for all species from the first ionization?

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

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1. Calculate the pH of a solution prepared by mixing 27.0 mL of pure pyridine, C5H5N (d = 0.978 g/mL), with 65.0 mL of 2.75 M HCl. Kb for pyridine is 1.5 × 10-9 . 2. Calculate the pH of a solution prepared by mixing 2.15 g of butyric acid (HC4H7O2) with 0.450 g of NaOH in water. (Ka butyric acid = 1.5 × 10-5 ) 3. How many grams of potassium hydroxide KOH would have to be added to 130 mL of a 0.298 M nitrous acid HNO2, solution, in order to prepare a buffer with a pH of 3.10? (Ka = 6.0 × 10-4 )
Acetic acid (pKa = 4.75) is contained in vinegar. 1. Calculate the pH of an 0.1500 M aqueous solution of acetic acid. 2. What is the pKb of the conjugate base of acetic acid, the acetate anion? 3. Calculate the pH of the solution prepared by mixing 0.0500 mol of acetic acid with 0.0150 mol of sodium acetate in a 500.0 mL volumetric flask.
Tartaric acid is found in many fruits including grapes. Tartaric acid is partly responsible for the dry texture of certain wines. Tartaric acid, H₂C4H406, is a diprotic acid with the following Ka values: Ka1 = 1.00 x 10-3; K₁2 = 4.60 x 10-5 a. Write the equation for the 1st ionization of tartaric acid. b. Write the equation for the 2nd ionization of tartaric acid. c. Write the Ka expression for the 1st ionization. d. Write the Ka expression for the 2nd ionization. e. What is the equilibrium concentrations of the reactants and products from the 1st ionization of a 0.250M solution of tartaric acid? f. What is the equilibrium concentrations of the reactants and products from the 2nd ionization? g. What is the total [H3O+] and the overall pH of the solution of tartaric acid?
At 22 ˚C, an excess amount of a generic metal hydroxide, M(OH)₂, is mixed with pure water. The resulting equilibrium solution has a pH of 10.20. What is the Ksp of the salt at 22 °C? K sp= APR 9 $ 4 % mtv G Search or type URL 5 MacBook Pro < 6 ONT17 X & 7 = 11 * 00 66 8 + A ( 9 < *** m - ) O W P { + 17 [ = } 1 delete
1. 0.200 M acetic acid is added to water. What is the concentration of H3O* in solution if Kc=1.8x10-6? 2. After an all-night drinking binge, a man pukes 300.0mL of liquid into his trash can. Assuming the puke is 20% HCI and the rest does not affect the pH, what is the pH of the solution resulting from the puke being neutralized with 0.50M NH3? (Kb-1.8x105) 3. If the initial concentration of NH3 is 0.350 M and the concentration at equilibrium is 0.325 M, what is Kc for this reaction?
5. I added 0.015 M solutions of my reactants together. At equilibrium, I found a 0.008 M concentration of HPO42. (Hint: Does NOT require use of the quadratic equation) H2PO4° (ag)+ H20 m +→ H30* (ag) + HPO2°G (ag) A) What is the Ka of my solution? B) What is the pH of my solution at equilibrium?
b) What is th 1 с E 10. If a 50.0-mL sample of 0.250 M NaCN is titrated by 0.250 M HCI (strong acid), what is the pH of the solution at the equivalence point? Ka of Hydrocyanic Acid (HCN) at 25 °C is 6.2 x10-10. Hydrocyanic acid ce reacts with water to produce hydronium and cyanate.
10. What chemical species, besides water, is present in the highest concentration in an aqueous solution of the weak base (CH3)2NH? 11. In each of the following pairs, identify the stronger acid and state a reason for your answer: H2PO, HPO, HNO2 HNO3 PH3 NH3 H2SO4 H2SEO4 PH3 H2S 12. In each of the following pairs, circle the solution that would have the lower pH 0.100 М НСI 0.100 м НС2НЗО2 0.100 М НСI 0.200 М НСI 0.100 M HC2H;O2 (Ka = 1.8 x 10) 0.100 M HCIO3 (Ka = 5.0 x 10) 0.100 M HC2H3O2 (Ka = 1.8 x 10) 0.100 M HC2H3O2 mixed with 0.100 M NaCzH3O2 %3D 13. It took 32.15 mL of 0.200 M NaOH to titrate 1.50 g of an unknown acid, HA. Calculate the molar mass of HA.
10. One characteristic of blood that we rarely give a thought to is its pH. The pH of blood must be held remarkably constant, varying only by a few hundredths of a pH unit from 7.36 to 7.40. Blood pH is kept constant primarily through the buffering action of the bicarbonate ion, HCO3-, and dissolved carbon dioxide, commonly represented as H2CO3. The second ionization is negligible in blood so that H2CO3 can be treated as a monoprotic acid. A) If the Ka1 for H2CO3 is 4.2 * 10-7, what ratio of HCO3- to H2CO3 is necessary to produce a buffer of pH 7.40? B) What is the pH of a buffer that is 1.0 M NaHCO3 and 1.0 M H2CO3? C) If you have 1.00 L of the buffer described above (1.0 M NaHCO3 and 1.0 M H2CO3) and 10.0 mL of 0.100 M HCl is added to it, what is the new pH? D) Just to prove that the buffer helps minimize changes in pH, calculate the pH of a solution prepared by adding 10.0 mL of 0.1 M HCl to 1.00 L of pure water.
The pH of blood is 7.4. What is the ratio (to one decimal place) of bicarbonate (conjugate base) to carbonic acid (weak acid) in the blood's buffering equilibrium? Note, the K₂ of carbonic acid is 4.27 x 10-7. Write your answer this way: # to 1 Done Clear Incorrect. The ratio is 10.7 to 1, indicating that there is 10 times more bicarbonate than carbonic acid. Enter the correct answer.
A buffer solution is 0.418 M in HCN and 0.342 M in KCN. If Ka for HCN is 4.0 × 10^-10, what is the pH of this buffer solution? pH = A buffer solution is 0.387 M in H2S and 0.212 M in KHS. If Ka1 for H2S is 1.0 x 10^-7, what is the pH of this buffer solution? pH = A 1.00 liter solution contains 0.43 M hypochlorous acid and 0.56 M potassium hypochlorite. If 0.14 moles of barium hydroxide are added to this system, indicate whether the following statements are true or false. (Assume that the volume does not change upon the addition of barium hydroxide.): A. The number of moles of HCIO will increase. B. The number of moles of ClO-will decrease. C. The equilibrium concentration of H3O+ will remain the same. v D. The pH will remain the same. E. The ratio of [HClO]/[ClO-] will remain the same. A 1.00 liter solution contains 0.44 moles hydrofluoric acid and 0.34 moles sodium fluoride. If 0.17 moles of hydroiodic acid are added to this system, indicate whether the following statements are…
bha A buffer prepared by dissolving oxalic acid dihydrate (H2C2O4·2H2O) and disodium oxalate (Na2C2O4) in 1.00 L of water has a pH of 5.333. How many grams of oxalic acid dihydrate (MW = 126.07 g/mol) and disodium oxalate (MW = 133.99 g/mol) were required to prepare this buffer if the total oxalate concentration is 0.241 M? Oxalic acid has pKa values of 1.250 (pKa1) and 4.266 (pKa2).