1. Check option A: Group 17 are all gases?

Group 17 elements are halogens: $\mathrm{F,\ Cl,\ Br,\ I,\ At}$.
At room temperature:

• $\mathrm{F_2}$ and $\mathrm{Cl_2}$ are gases,
• $\mathrm{Br_2}$ is a liquid,
• $\mathrm{I_2}$ is a solid.

So, all Group 17 elements are not gases.
Hence, A is false.

2. Check option B: Group 13 elements are all metals?

Group 13 elements are $\mathrm{B,\ Al,\ Ga,\ In,\ Tl}$.
Here, boron ($\mathrm{B}$) is a metalloid, not a metal.

So, all Group 13 elements are not metals.
Hence, B is false.

3. Check option C: Group 16 elements have lower ionization enthalpy than Group 15 elements in the same period?

Compare electronic configurations:

• Group 15: $ns^2 np^3$ (half-filled $p$-subshell, extra stability)
• Group 16: $ns^2 np^4$

Because half-filled subshells are relatively more stable, removing an electron from Group 15 requires more energy than from Group 16 in the same period.

Therefore, $\text{Ionization enthalpy of Group 16} < \text{Ionization enthalpy of Group 15}$ within the same period.

Hence, C is true.

4. Check option D: For Group 15, stability of $+5$ oxidation state increases down the group?

In Group 15 ($\mathrm{N, P, As, Sb, Bi}$), due to the inert pair effect, the lower oxidation state $+3$ becomes more stable down the group, while the higher oxidation state $+5$ becomes less stable.

So, the stability of $+5$ oxidation state decreases down the group.

Hence, D is false.

Final conclusion

Only option C is correct.