Qualitative attributes like strong, weak, very weak, very very weak (acid/base) have no widely accepted agreement/definition what it exactly means quantitatively. As consequence, different sources may give different boundaries between them.
That is the reason why you have ended with different evaluation for acids and bases.
In context of water solutions, acids have the characteristic acidity parameter $\mathrm{p}K_\mathrm{a}$. It is equal to $\mathrm{pH}$ when $\pu{50 \%}$ of the acid is dissociated.
\begin{array} {|l|c|} \hline \mathrm{pH} & \text{% of acid dissociation} \\ \hline \mathrm{p}K_\mathrm{a} - 3 & \approx \pu{0.1 \%} (\frac{1}{1001}) \\ \hline \mathrm{p}K_\mathrm{a} - 2 & \approx \pu{1 \%} (\frac{1}{101}) \\ \hline \mathrm{p}K_\mathrm{a} - 1 & \approx \pu{9 \%} (\frac{1}{11}) \\ \hline \mathrm{p}K_\mathrm{a} & \pu{50 \%} (\frac{1}{1}) \\ \hline \mathrm{p}K_\mathrm{a} + 1 & \approx \pu{91 \%} (\frac{10}{11}) \\ \hline \mathrm{p}K_\mathrm{a} + 2 & \approx \pu{99 \%} (\frac{100}{101}) \\ \hline \mathrm{p}K_\mathrm{a} + 3 & \approx \pu{99.1 \%} (\frac{1000}{1001}) \\ \hline \end{array}\begin{array} {|l|c|} \hline \mathrm{pH} & \text{% of acid dissociation} \\ \hline \mathrm{p}K_\mathrm{a} - 3 & \approx \pu{0.1 \%} (\frac{1}{1001}) \\ \hline \mathrm{p}K_\mathrm{a} - 2 & \approx \pu{1 \%} (\frac{1}{101}) \\ \hline \mathrm{p}K_\mathrm{a} - 1 & \approx \pu{9 \%} (\frac{1}{11}) \\ \hline \mathrm{p}K_\mathrm{a} & \pu{50 \%} (\frac{1}{2}) \\ \hline \mathrm{p}K_\mathrm{a} + 1 & \approx \pu{91 \%} (\frac{10}{11}) \\ \hline \mathrm{p}K_\mathrm{a} + 2 & \approx \pu{99 \%} (\frac{100}{101}) \\ \hline \mathrm{p}K_\mathrm{a} + 3 & \approx \pu{99.1 \%} (\frac{1000}{1001}) \\ \hline \end{array}
Mutual reactions of acids with bases have to be assessed in sense what $\mathrm{pH}$ wrt the acid $\mathrm{p}K_\mathrm{a}$ is set by the base excess.
In context of Broensted-Lawry acid-base theory, basedbases accept H+ ion:
$$\ce{B(aq) + H+(aq) <=> BH+(aq)\text{(conjugate acid)}}$$
where $\ce{B}$ may be a molecule or a ion.
If the base has the basicity constant $\mathrm{p}K_\mathrm{b}$ then its conjugate acid has the acidity constant (at $\pu{25 \circ C}$$\pu{25 ^\circ C}$)
$\mathrm{p}K_\mathrm{a}$$\mathrm{p}K_\mathrm{a, conj}$ = 14 - $\mathrm{p}K_\mathrm{b}$
We can then compare the relative strength of the given acid and the conjugate acid of the given base. (That applies on a solvent too, if it can act as a base like water)
For the stoichiometric ratio of the given acid and base, we can estimate the degree of neutralization.
\begin{array} {|c|r|} \hline \mathrm{p}K_\mathrm{a,conj} - \mathrm{p}K_\mathrm{a} & \text{Neutralization degree} \\ \hline 8 & \pu{0.01 \%} \\ \hline 6 & \pu{0.1 \%} \\ \hline 4 & \pu{1 \%} \\ \hline 2 & \pu{9 \%} \\ \hline 0 & \pu{50 \%} \\ \hline -2 & \pu{91 \%} \\ \hline -4 & \pu{99 \%} \\ \hline -6 & \pu{99.9 \%} \\ \hline -8 & \pu{99.99 \%} \\ \hline \end{array}