Order these substances from highest to lowest #"pH"#?

#"Sr"("OH")_2(aq) > "NaOH"(aq) > "NH"_3(aq) > "HBrO"(aq) > "HBr"(aq)#

Here's what we should know:

• The stronger the base, the higher the #"pH"#, since it dissociates more in solution than weaker bases, thus making it more basic and raising the #"pH"#.

• The stronger the acid, the lower the #"pH"#, since it dissociates more in solution than weaker acids, thus making it more acidic and lowering the #"pH"#.

Now let's figure out where the acids and bases fall on the pH scale. The strong bases within this list are:

• #"Sr"("OH")_2#
• #"NaOH"#

The weak base within this list is #"NH"_3#.
The strong acid within this list is #"HBr"#.
The weak acid within this list is #"HBrO"#.

(You have to memorize the common strong acids and bases.)

Furthermore, the strong bases with more hydroxides per formula unit dissociate more hydroxides into solution. That means #"Sr"("OH")_2# is twice as concentrated with #"OH"^(-)# as #"NaOH"# is.

So far, we thus have that pH varies as follows when these are placed into solution:

#"Sr"("OH")_2(aq) > "NaOH"(aq) > ? > ? > "HBr"(aq)#

We now know that #"HBrO"# is a weak acid, so it must dissociate less than #"HBr"#, meaning that it decreases the #"pH"# by less from about #7#. That means the #"pH"# of #"HBrO"(aq) > "pH"# of #"HBr"(aq)#.

Now we have:

#"Sr"("OH")_2(aq) > "NaOH"(aq) > ? > "HBrO"(aq) > "HBr"(aq)#

Since #"NH"_3# is a weak base, it increases the #"pH"# by less than #"NaOH"#, but since it increases the #"pH"# and #"HBrO"# decreases the #"pH"#, it means that in general the same concentration of both implies that the #"pH"# is higher for #"NH"_3(aq)# than #"HBrO"(aq)#.

Therefore, our result is:

#bb("Sr"("OH")_2(aq) > "NaOH"(aq) > "NH"_3(aq) > "HBrO"(aq) > "HBr"(aq))#