How many grams of glycine amide (FM 74.08) should be used with 1.0 gram of glycine amide HCl (FM 110.54, pKa=8.20) to give 100 mL at pH 8.0?

The buffer equilibrium reaction is:
GlycineAmideHCl <--- ---> GlycineAmide + H+
which has the general form:
BH+ <--- Ka ---> B + H+

I'm assuming I need to use the henderson-hasselbalch equation which would give something like:

8 = 8.2 + log((B/BH+))

Though I'm not sure what the next step is or what to do with the B and BH+

Take you HH equation of

8.00 = 8.20 + log(B/A)
and solve for (B)/(A). I get about 0.63, then
base = 0.63A
You have 1 g of the acid and you know the molar mass, substitute for A which will allow you to calculate A, then solve for B, then determine how many grams that will be in 100 mL.

So 8.00 = 8.20 + log(B/A)

=> -0.2 = log(B/A)
=> 10^-0.2 = 10^log(B/A)
=> 0.63 = B/A
=> base = 0.63A as you said

Though I'm still having trouble understanding what you mean by "substitute for A which will allow you to calculate A, then solve for B"

Ohh, I got it now =)

good.

To solve this problem, you are correct that you need to use the Henderson-Hasselbalch equation, which relates the pH of a buffer solution to the ratio of the concentration of its conjugate base (B) and its weak acid (BH+).

In this case, the weak acid is glycine amide HCl (BH+) and the conjugate base is glycine amide (B).

The Henderson-Hasselbalch equation is:

pH = pKa + log([B] / [BH+])

You want to understand how much glycine amide (B) you need to add to glycine amide HCl (BH+) in order to obtain a final pH of 8.0.

Here's a step-by-step approach to solving the problem:

1. Start by rearranging the Henderson-Hasselbalch equation to isolate [B] / [BH+]:

[B] / [BH+] = 10^(pH - pKa)

2. Substitute the given pH (8.0) and pKa (8.20) values into the equation:

[B] / [BH+] = 10^(8.0 - 8.20)

3. Calculate the ratio [B] / [BH+]:

[B] / [BH+] = 10^(-0.20) [Note: 10^(-0.20) is approximately 0.63]

4. Now, let's use stoichiometry to determine the amount of glycine amide (B) required.

The molecular weight (FM) of glycine amide is 74.08 g/mol.

If we let "x" represent the mass of glycine amide (B) in grams, then the mass of glycine amide HCl (BH+) would be (1.0 - x) grams since you are adding 1.0 gram of glycine amide HCl initially.

The moles of glycine amide (B) can be calculated as:

moles of B = x / FM of glycine amide

Similarly, the moles of glycine amide HCl (BH+) can be calculated as:

moles of BH+ = (1.0 - x) / FM of glycine amide HCl

5. Use the stoichiometric ratio from the equilibrium reaction (1:1) to set up an equation relating the moles of B and BH+:

moles of B = moles of BH+

x / FM of glycine amide = (1.0 - x) / FM of glycine amide HCl

6. Solve this equation for x, which represents the mass of glycine amide (B) in grams:

x = [(1.0 - x) / FM of glycine amide HCl] * FM of glycine amide

Simplifying this equation will give you the value of x, which is the mass of glycine amide (B) that should be used.

By following these steps, you should be able to determine the amount of glycine amide (B) needed to achieve a pH of 8.0 in a 100 mL solution.