Question

An analytical chemist is titrating of a solution of ethylamine with a solution of . The of ethylamine is . Calculate the pH of the base solution after the chemist has added of the solution to it. Note for advanced students: you may assume the final volume equals the initial volume of the solution plus the volume of solution added. g

Answer 1

pH=11.

Step-by-step explanation:

Hello!

In this case, since the data is not given, it is possible to use a similar problem like:

"An analytical chemist is titrating 185.0 mL of a 0.7500 M solution of ethylamine(C2HNH2) with a 0.4800 M solution of HNO3.ThepK,of ethylamine is 3.19. Calculate the pH of the base solution after the chemist has added 114.4 mL of the HNO3 solution to it"

Thus, for the reaction:

`C_2H_5NH_2+H^+\rightleftharpoons C_2H_5NH_4^+`

Tt is possible to compute the remaining moles of ethylamine via the following subtraction:

`n_(ethylamine)=0.1850L*0.7500mol/L=0.1365mol\\\\n_(acid)=0.1144L*0.4800mol/L=0.0549mol\\\\n_(ethylamine)^(remaining)=0.1365mol-0.0549mol=0.0816mol`

Thus, the concentration of ethylamine in solution is:

`[ethylamine]=(0.0816mol)/(0.1850L+0.1144L)=0.2725M`

Now, we can also infer that some salt is formed, and has the following concentration:

`[salt]=(0.0549mol)/(0.1850L+0.1144L)=0.1834M`

Therefore, we can use the Henderson-Hasselbach equation to compute the resulting pOH first:

`pOH=pKb+log(([salt])/([base]) )\\\\pOH=3.19+log((0.1834M)/(0.2725M))\\\\pOH=3.0`

Finally, the pH turns out to be:

`pH=14-pOH=14-3\\\\pH=11`

NOTE: keep in mind that if you have different values, you can just change them and follow the very same process here.

Best regards!