What is the pH of a 0.1 M acid solution?

Question

Natalie Anton · Accepted Answer

For a strong acid, #pH# #=# #-log_10[H_3O^+]# #=# #-log_10[0.1]=-log_10(10^-1)# #=# #-(-1)# #=# #1#

Assuming a strong acid, the equilibrium reaction is strongly to the right, as indicated below.

#HX(aq)+H_2O(l)rightleftharpoonsX^(-) + H_3O^(+)#

Why do we use #pH#? Well it is a holdover from the pre-electronic calculator days, which some of us can remember (not me of course, I am not a day over 29!).

When I write #log_ab=c#, I explicitly say that #a^c=b#. Common logarithmic bases are #10# and #e# (I think you cover this in A level mathematics these days). #pH# and #pK_a# scales use logarithms to the base #10#.

Anyway given what I have said #log_(10)100=2#, #log_(10)1000=3#, and #log_(10)1=0#, and #log_(10)0.1=log_(10)10^-1=-1#. If you can grasp this, you will find the #pH# concept (literally #"pouvoir hydrogene, power of hydrogen"#) fairly straightforward.

With weaker acids, say acetic acid, #H_3C-C(=O)OH#, the acid-base equilibrium lies to the left. The equilibrium constant, #K_a#, can be measured and often reported as #pK_a# #=# #-log_10K_a#.

Wyatt Atkins · Answer

undefined The pH of a 0.1 M acid solution depends on the specific acid. Without knowing the identity of the acid, it's not possible to determine the pH accurately. Different acids have different dissociation constants (Ka values), which influence the concentration of hydrogen ions (H⁺) in solution and, consequently, the pH.