A 0.500 g sample of impure CaCO3 is dissolved in 50.00 mL of 0.0985 M HCl. After the reaction is complete, the excess acid is titrated by 6.0 mL of 0.105 M NaOH. Find the percent purity of the CaCO3 sample. (CaCO3 = 100 g/mole)

section 2 , question 2 For humans to maintain the pH of their blood close to 7.4, the body uses three buffer systems: carbonate, phosphate and proteins. If the pH of the blood goes below 6.8 or above 7.8, death may result. Imagine you want to make a phosphate buffer in the lab that resembles the bloods phosphate buffer. You start by mixing 1.0 M NaH2PO4 and 0.70 M Na2HPO4. H2PO4-  + H2O ⇌  HPO42-  + H+          pKa= 7.21                From the reaction above indicate the weak acid and the conjugate base involved in making this buffer. Calculate the pH of the phosphate buffer solution. Provide your answer with 2 decimal places If the pH calculated in b. was the blood pH of an actual patient, would you say this patient’s blood pH is in the normal range or are they suffering from alkalosis or acidosis?

For humans to maintain the pH of their blood close to 7.4, the body uses three buffer systems: carbonate, phosphate and proteins. If the pH of the blood goes below 6.8 or above 7.8, death may result. Imagine you want to want to make a phosphate buffer in the lab that resembles the bloods phosphate buffer. You start by mixing 0.40 M NaH2PO4 and 0.60 M Na2HPO4. H2PO4-  + H2O ⇌  HPO42-  + H+          pKa= 7.21    From the reaction above indicate the weak acid and the conjugate base involved in making this buffer.  Calculate the pH of the phosphate buffer solution? Provide your answer to 2 decimal places  If the pH calculated in b was the blood pH of an actual patient, would you say this patient’s blood pH is within normal range or are they suffering from alkalosis or acidosis?

When 30.00 mL of 0.1011 M HCl in 50 mL of deionized water is titrated against 0.09889 M NaOH, the pH increases. - Find the pH when the volume of NaOH added is 0.02 mL less than the volume required to reach the equivalence point.   - Find the pH when the volume of NaOH added is 0.01 mL less than the volume required to reach the equivalence point.

The concentration of ammonia (17.031 g/mol) can be determined through back titration. After addition of 10.00 mmol of HCl into a sample of NH₃ solution, back-titration with Na₂CO₃ revealed that 5.00 mmol of HCl remained. What would be the concentration in %w/v of ammonia if the 10.00 mL of the sample was analyzed?

2. In the second titration, the molar mass of an unknown monoprotic acid is calculated.  If 0.100 grams of unknown solid are dissolved in water and titrated with 4.68 mL of the NaOH, what is the molar mass of the solid? HINTS: Molar mass is grams / moles. We have grams from the weighed mass, we can calculate moles of acid from the titration. Monoprotic means that it produces 1 H+ per molecule. This means that our moles of NaOH added and moles of acid are the same. You will need to use your concentration of NaOH from the previous question Which is 0.1186

120. A diprotic acid with a molar mass of 118.04 g/mol was titrated with 1.023 M NaOH. The equivalence point of the titration occurred at 22.16 mL. Calculate the mass of the acid.

For humans to maintain the pH of their blood close to 7.4, the body uses three buffer systems: carbonate, phosphate and proteins. If the pH of the blood goes below 6.8 or above 7.8, death may result. Imagine you want to make a phosphate buffer in the lab that resembles the bloods phosphate buffer. You start by mixing 1.0 M NaH2PO4 and 0.50 M Na2HPO4. H2PO4-  + H2O ⇌  HPO42-  + H+         pKa= 7.21 From the reaction above indicate the weak acid and the conjugate base involved in making this buffer. Calculate the pH of the phosphate buffer solution. Provide your answer with 2 decimal places. If the pH calculated in b. was the blood pH of an actual patient, would you say this patient’s blood pH is in the normal range or are they suffering from alkalosis or acidosis?

Please only answer the question below. The questions shown are already solved, you don’t need to add anything, they are there for context. 3) Calculate the [Ba2+] left in solution when Ca2+ starts to precipitate. Then calculate the % of the original Ba2+ left in solution when Ca2+ starts to precipitate. A) 1.1 • 10^-2 % B) 2.8 • 10^-7 % C) 9.3 • 10^-3 % D) 6.7 • 10^-3 % E) 5.6 • 10^-3 %

question 7!