[FutureDoc]

Can someone please explain what's going on in this problem, and how you'd set up the chemical equation and ICE table? Once I get there, I can solve it fairly easily, but I'm confused on this one. Thanks in advance to all help.

What is the final pH if 0.03 mol HCl is added to 0.500 L of a buffer solution that is 0.24 M NH₃ and 0.20 M NH₄Cl?

Again, I'm not asking anyone to solve the entire problem, but rather just help me set it up. Thanks!

[LQ43]

Which buffer component NH3 or NH4+ will the H+ be a common ion for?

Then use that equation to start your ice table

Or have you heard of the Henderson-Hasselbach equation?

[FutureDoc]

Which buffer component NH3 or NH4+ will the H+ be a common ion for?

How do I know which component the H+ is a common ion for?

Also, what did do you do with the HCl?

[LQ43]

HCl will provide H+ that will affect the buffer and so the pH

Can you write out the ionization equations for

NH3  +  H2O -->

NH4+  -->

then look at the products and see which one H+ is in - that is the equation you need to use for the ICE table. With HCl you are adding more H+ - that is the common ion when you add acid to a buffer.

[FutureDoc]

HCl will provide H+ that will affect the buffer and so the pH

First off, why do you not include the Cl^- at all in your equation?

Can you write out the ionization equations for

NH3  +  H2O -->

NH4+  -->

then look at the products and see which one H+ is in - that is the equation you need to use for the ICE table. With HCl you are adding more H+ - that is the common ion when you add acid to a buffer.

NH₃ + H₂O --> NH₄⁺ + OH^-

Not so sure on this next one:
N_H4⁺ --> NH₃^- + H⁺ ?

Sorry for all the questions, but I'm absolutely clueless on this one.

[LQ43]

First off, why do you not include the Cl^- at all in your equation?

Not so sure on this next one:
N_H4⁺ --> NH₃^- + H⁺ ?

Sorry for all the questions, but I'm absolutely clueless on this one.

Cl- is too weak of a base to affect the pH - its a spectator ion

there is no charge on NH3, other than that, (fix the subscript on H) the equation is correct

- your textbook must have an example of adding and acid or base to a buffer and subsequent calculations.

- look up the Henderson-Hasselbach equation, its very useful in buffer calculations

[FutureDoc]

Okay, here's what I get so far when I work it out:

NH₄Cl --> NH₄⁺ + Cl^-
(Cl- is conjugate of strong acid, so not included in reaction equation.)

     NH₃ + H₂O <--> NH₄⁺ + OH^-
I:  0.24                   0.20      0
C:  -x                       +x      +x
E: 0.24-x               0.20+x     x

K_b= 1.8x10^-5 = (0.20)x/0.24, so x=2.16x10^-5

After addition of 0.03 mol HCl:
             NH₃ + H⁺  <--->  NH₄⁺ +   OH^-
Initial: 0.12 mol    0         0.4 mol
Add:                +0.03
Change: -0.03   -0.03      +0.03
Ntrlzatn: 0.09      0          0.43
Molarity: 0.18 M   0 M      0.86 M

pH = pK_a + log ([base]/[acid])

Is all this correct so far? If so, where do I go from there?

[Borek]

E: 0.24-x               0.20+x     x

K_b= 1.8x10^-5 = (0.20)x/0.24, so x=2.16x10^-5

Note that x is so small - compared with total ammonia concentration - that it can be neglected. That means that you in fact already know concentrations and no ICE calculations are needed.

After addition of 0.03 mol HCl:
             NH₃ + H⁺  <--->  NH₄⁺ +   OH^-
Initial: 0.12 mol    0         0.4 mol
Add:                +0.03
Change: -0.03   -0.03      +0.03
Ntrlzatn: 0.09      0          0.43

These are your concentrations of ammonia and NH₄⁺ after neutralization. Which one is acid and which one is conjugated base?

pH = pK_a + log ([base]/[acid])

Just plug...