As you know, water undergoes the equilibrium reaction:
#H_2O rightleftharpoons H^+ + OH^-#
(These days, it's a little bit more common to speak of the hydronium ion, #H_3O^+#: #2H_2O rightleftharpoons H_3O^+ + OH^-#) Both equations represent the autoprotolysis of water. Now, this equilibrium has been extensively studied, and under standard conditions; the equilibrium constant for the reaction, #K_w = 10^(-14)# #=# #[H^+][OH^-]#.
If the solution is neutral, then #[H^+]=[OH^-]# (and #pH = pOH = 7#, where #pH = -log_10[H^+]# and #pOH = -log_10[OH^-]#).
#K_w# is quoted under standard conditions of #1# #atm# and #298K#. What do you think would happen under non-standard conditions, say at a temperature of #350# #K#; would #pK_w# go up or go down? (Be careful with the negative sign!)