Which Bases Could Be Used In The Preperation Of A Buffer With A Ph Of 9.8

HELP! Acids, Bases, and Buffers Problem!?

Hey Diana,

Why don't you use your gen chem book if you still have it? You and I both go to GCU, so if you have your book, go to page 731 and it talks about preparing buffers with a specific solution. You would still use the Hasselbach equation, yes. But remember you have to write out the chemical equation its asking for. So you have Carbonic acid + either NaOH or HCL +H2O.

Not sure if this helps a little or not at all, but the book really helps!

Biochem Buffer question?

OK....in both cases, begin with the Henderson-Hasselbalch equation. For the first problem:

pH = pKa + log [NH3]/[NH4+]
9.4 = 9.26 + log [NH3]/[NH4+]
[NH3]/[NH4+] = 1.38

Now, you know from the working of the problem that [NH3] + [NH4+] = 0.05 M

You can rearrange this: [NH3]/[NH4+] = 1.38 to be

[NH3] = 1.38 [NH4+]. Now, substitute this into the equation above to get:

1.38 [NH4+] + [NH4+] = 0.05
2.38 [NH4+] = 0.05
[NH4+] = 0.021M

Using that and either of the equation, you can show [NH3] = 0.05 - 0.021 = 0.029M

Now, to make the buffer, you need to calculate the mass of NH4Cl that will give you 2 L of a 0.05 M solution:

0.05 mol/L X 2 L X 53.49 g/mol = 5.349 g NH4Cl

Now, when you dissolve that in water, the [NH4+] is essentially 0.05 M. But, you need to convert 0.029 M of that into NH3 by adding NaOH.

So, the amount of NaOH you have to add is:

0.029 mol/L X 2 L = 0.058 mol / 2.13 mol/L = 0.0272 L = 27.2 mL

So, weigh out 5.349 grams NH4Cl, dissolve in 1.8 L H2O or so, add 27.2 mL of the NaOH solution, and dilute to 2 L with water. That SHOULD give you the correct buffer.

Now, the second one is really a little easier. First, you can completely ignore the pK2, since you are making a buffer of pH = 2.4. Now, the second thing to recognize is that for this buffer, pH = pKa.

Looking at the H-H equation again,

pH = pKa + log [COO-]/[COOH]
Since pH = pKa, the ratio of conjugate acid to conjugate base must equal 1.

Since the buffer will contain 0.15 M glycine the absolute concentrations of the acid and base forms of glycine will both be 0.075 M. Now, in its solid form, glycine exists as the zwitterion, so all of the carboxyl groups are in the COO- form. You must add enough HCl to convert half of the glycine from COO- to COOH.

So, since you want 0.200 L of the buffer, you will need:

0.200 L X 0.15 mol/L = 0.030 mol glycine X 75 g/mol = 2.25 g glycine

Volume of HCl needed:

0.030 mol glycine / 2 = 0.015 mol HCl / 1 mol/L = 0.015 L = 15 mL HCl.

Combine these two things in a total final volume of 200 mL and you've got it...

Hope that makes sense...

What masses of HCN and NaCN should be used to prepare a buffer?

You can use the information about osmotic pressure to first calculate the total concentration of HCN and NaCN, and then using the Henderson Hasselbalch equation, you can calculate the ratio of CN- to HCN, and then the masses of each.

Osmotic pressure = π = iMRT where i is van't Hoff factor, in this case i = 3 because NaCN gives Na+ and CN-, and HCN give H+ and CN-. So you have H+, Na+ and CN- and i = 3. Solve for M.
M = π/(i)(R)T) = 1.72/(3)(0.082)(298) = 0.023 M
0.023 mol/L x 1.9 L = 0.0446 moles

pH = pKa + log [NaCN]/[HCN]
look up pKa for HCN, substitute in the equation and solve for the ratio [NaCN]/[HCN].
Once you have that, and you know you need 0.0446 moles, you can solve for the moles NaCN and moles of HCN, and from that you can calculate the mass of each.

A FEW BUFFER QUESTii0NS...?

N0T 2 C0MPLiiCATED...JUST N0T SURE WHERE 2 START FR0M...ANY HELP W0ULD BE GREAT...SH0W STEPS PLEASE && THANKS iiN ADVANCE
1. What is the pH of a solution prepared by dissolving 31 mmol of a weak base in 250 mL of a 0.139-M solution of its conjugate acid if pKa = 5.06 for the acid.

2. How many mL (to the nearest mL) of 0.182-M KF solution should be added to 360. mL of 0.189-M HF to prepare a pH = 2.60 solution?

3. How many grams (to the nearest 0.01 g) of NH4Cl (Mm = 53.49 g/mol) must be added to 700. mL of 1.294-M solution of NH3 in order to prepare a pH = 9.70 buffer?

4. What volume (to the nearest 0.1 mL) of 5.40-M NaOH must be added to 0.350 L of 0.350-M HNO2 to prepare a pH = 4.20 buffer?

5. What volume (to the nearest 0.1 mL) of 4.90-M HCl must be added to 0.750 L of 0.300-M K2HPO4 to prepare a pH = 7.00 buffer?

What would be acceptable conjugate acid-base pairs to use in the preparation of a pH=9.8 buffer?

NH4Cl / NH3 OR NH4OH IS THE PAIR THAT MAINTAINS THE pH FROM 9 -11

) What is the hydronium ion concentration in a solution prepared by mixing 50.00 mL of 0.10 M HCN with 50.0?

What you've made is a buffer solution. You can calculate the pH (and by extension, the hydronium concentration) of a buffer solution using the Henderson-Hasselbalch equation:

pH = pKa + log([A-] / [HA])

Now we just have to calculate the concentrations of HCN (the acid component of the buffer) and CN- (the base component). For that we can use the equation M1xV1 = M2xV2. If volumes are additive, the final volume of the solution will be 100.00 mL

For the acid, HCN:

M1xV1 = M2xV2

(0.10 M)(50.00 mL) = M2(100.00 mL)

M2 = 0.050 M

For the base, CN-

M1xV1 = M2xV2

(0.050 M)(50.00 mL) = M2(100.00 mL)

M2 = 0.025 M

Now use the H-H equation:

pKa = -log(4.9x10^-10) = 9.31

pH = 9.31 + log(0.025 M / 0.050 M)

pH = 9.31 + log(0.50)

pH = 9.31 + -0.30

pH = 9.00

Given the pH, we can find the hydronium concentration by taking the anti-log of 9.00

[H3O+] = 10^-9.00 = 1.0x10^-9 M

I hope that helps. Good luck!