How does acid affect the pH and H concentration when added to water?

When an acid is added to water it is conceived to increase concentrations of the characteristic cation (of water!):

#HX(aq) + H_2O(l) rightleftharpoons H_3O^+ + X^-#

This affects the equilibrium that already operates in water under standard condtions:

#2H_2O rightleftharpoons H_3O^+ + HO^-#, where
#[H_3O^+][""^(-)OH]# #=# #10^(-14)#

Note that if we take negative logs to the base 10, we get an expression that should be familiar:

#-log_(10)[H_3O^+] -log_(10)[""^(-)OH]# #=# #-log_(10)10^-14# #=# #14#

OR, using the standard definition, #pH=-log_(10)[H_3O^+]#, and #pOH=-log_(10)[HO^-]#

#pH + pOH =14#

So, at neutrality, i.e. #[HO^-]=[H_3O^+]#, #pH=7#, and #pOH=7#

For strong acids, e.g. hydrogen halides, perchloric acid, sulfuric acid, the first equilibrium lies strongly to the right. Whatever the strength of the acid, #[H_3O^+]# is altered from neutrality.

And thus in an acidic solution, #pH<7#, whereas #pH>7# in an alkaline solution.

I acknowledge that you have been hit with a lot of facts and definitions. Review your notes, and ask for clarification if there are queries.