What mass of solid sodium formate (of MW

68.01) must be added to 115 mL of 0.57 mol/L
formic acid (HCOOH) to make a buffer solution
having a pH of 4.16? Ka = 0.00018 for
HCOOH.
Answer in units of g

millimols HCOOH = mL x M = approx 65 but you need to do all of these calculations more accurately.

mmols HCOONa = base

pH = pKa + log (base)/(acid)
Substitute and solve for (base) in millimols.
Change to mols, then g = mols x molar mass.
You know mols and molar mass

whats the answer?

Well, sodium formate is like the Batman to formic acid's Robin. Together, they make a great buffer solution team!

To find out how much sodium formate we need, we can use the Henderson-Hasselbalch equation:

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

We know the pH is 4.16 and the pKa is given as 0.00018. Let's plug in those values and solve for the ratio [A-]/[HA].

4.16 = -log(0.00018) + log([A-]/[HA])

Now, if there was a superhero duo called "Sodium Formate and Formic Acid", they would have the perfect partnership. You see, the ratio [A-]/[HA] in their team-up is 1:1, with equal amounts of their concentrations. So, we can rewrite the equation as:

4.16 = -log(0.00018) + log(1)

Now, let's solve for -log(0.00018):

4.16 + log(0.00018) = 0

Finally, we can plug this value into the equation below to find the mass of sodium formate (x):

x = (0.57 mol/L) * (115 mL) * (68.01 g/mol) / (1000 mL/L)

And the answer is... *drum roll*:

x = 0.447 grams

So, you'll need approximately 0.447 grams of solid sodium formate to make that pH 4.16 buffer solution. But remember, even though sodium formate and formic acid may be a great duo, they don't take requests for costume changes.

To calculate the mass of solid sodium formate needed to make a buffer solution, we need to consider the Henderson-Hasselbalch equation for buffer solutions:

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

Given:
pH = 4.16
Ka = 0.00018 (Note: pKa = -log(Ka))

From the equation, we can derive:
[A-]/[HA] = 10^(pH - pKa)

Given that the concentration of formic acid ([HA]) is 0.57 mol/L, we can calculate the concentration of sodium formate ([A-]) using the derived equation.

[A-]/[HA] = 10^(4.16 - (-log(0.00018)))
[A-]/[HA] = 10^4.16 / 0.00018
[A-]/[HA] ≈ 294926.71

Since sodium formate (HCOONa) dissociates into one sodium ion (Na+) and one formate ion (HCOO-), the concentration of sodium formate ([HCOONa]) will be the same as [A-], which is approximately 294926.71 mol/L.

Now, let's convert the volume of formic acid from milliliters (mL) to liters (L):
115 mL = 115/1000 L = 0.115 L

Next, let's calculate the number of moles of formic acid (HCOOH) present in the given volume:

moles of HCOOH = concentration × volume
moles of HCOOH = 0.57 mol/L × 0.115 L
moles of HCOOH ≈ 0.06555 mol

To create a buffer solution, we need equal moles of formic acid and sodium formate. Therefore, we'll need approximately 0.06555 mol of sodium formate.

Finally, we can calculate the mass of solid sodium formate needed using its molar mass:

mass of sodium formate = moles × molar mass
mass of sodium formate = 0.06555 mol × 68.01 g/mol
mass of sodium formate ≈ 4.46 g

Therefore, approximately 4.46 grams of solid sodium formate must be added to 115 mL of 0.57 mol/L formic acid to create a buffer solution with a pH of 4.16.

To find the mass of solid sodium formate that needs to be added to the formic acid solution, we need to use the Henderson-Hasselbalch equation. The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa and the ratio of the concentration of the conjugate base to the weak acid.

The Henderson-Hasselbalch equation is given by: pH = pKa + log([conjugate base]/[weak acid])

In this case, the weak acid is formic acid (HCOOH), and its conjugate base is the formate ion (HCOO-). The molecular weight (MW) of sodium formate is provided as 68.01 g/mol.

We are given that the pH of the buffer solution is 4.16, and the pKa of formic acid is given as 0.00018. Since the pKa is very small compared to the pH, we can assume that [conjugate base]/[weak acid] is approximately equal to 1.

Therefore, we can rearrange the Henderson-Hasselbalch equation to find the concentration of formate ion ([conjugate base]) in the solution:
pH = pKa + log([conjugate base]/[weak acid])
4.16 = 0.00018 + log(1)
4.16 - 0.00018 = log(1)
4.15982 = log(1)
Antilog(4.15982) = 1

Since the concentration of formate ion is equal to the concentration provided for the weak acid (0.57 mol/L), we can calculate the moles of formate ion in the solution:
moles of formate ion = concentration of formate ion x volume of formic acid solution
moles of formate ion = 0.57 mol/L x 115 mL = 0.57 mol/L x 0.115 L = 0.06555 mol

Finally, we can calculate the mass of sodium formate needed to supply this amount of moles:
mass of sodium formate = moles of formate ion x MW of sodium formate
mass of sodium formate = 0.06555 mol x 68.01 g/mol = 4.46 g

Therefore, approximately 4.46 g of solid sodium formate must be added to the formic acid solution to make a buffer solution with a pH of 4.16.