This is a multiple-correct-answer question where we need to evaluate each statement about Group 15 elements.

Statement A: $B{i_2}{O_5}$ is more basic than ${N_2}{O_5}$

1. Analyze the elements: Nitrogen (N) and Bismuth (Bi) are both in Group 15. Nitrogen is at the top of the group, and Bismuth is at the bottom.
2. Trend in properties: As we move down a group in the p-block, the metallic character of the elements increases. Nitrogen is a non-metal, while Bismuth is a metal.
3. Nature of oxides: Oxides of non-metals are typically acidic, while oxides of metals are typically basic. For oxides of elements in the same group and the same oxidation state (+5 in this case), the acidic character decreases and the basic character increases as we move down the group.
4. Comparison: $N_2O_5$ is the oxide of a non-metal and is strongly acidic. $Bi_2O_5$ is the oxide of a metal. Following the trend, $Bi_2O_5$ is significantly less acidic than $N_2O_5$. Being less acidic implies it is relatively more basic. The order of acidity is $N_2O_5 > P_2O_5 > As_2O_5 > Sb_2O_5 > Bi_2O_5$. Therefore, the statement that $Bi_2O_5$ is more basic than $N_2O_5$ is correct.

Statement B: $N{F_3}$ is more covalent than $Bi{F_3}$

1. Analyze the compounds: We are comparing the covalent character of nitrogen trifluoride ($NF_3$) and bismuth trifluoride ($BiF_3$).
2. Apply Fajan's Rules: Covalent character is favored by a small cation and a large anion. Here, the anion is the same ($F^-$). The cations are $N^{3+}$ and $Bi^{3+}$.
3. Cation size: Atomic size increases down the group, so the ionic radius of $N^{3+}$ is much smaller than that of $Bi^{3+}$.
4. Polarizing power: A smaller cation has a higher charge density and thus greater polarizing power. $N^{3+}$ will polarize the electron cloud of the fluoride ion more effectively than $Bi^{3+}$. Greater polarization leads to a higher degree of covalent character.
5. Conclusion: Therefore, $NF_3$ is more covalent than $BiF_3$. This statement is correct.

Statement C: $P{H_3}$ boils at lower temperature than $N{H_3}$

1. Analyze the compounds: We are comparing the boiling points of ammonia ($NH_3$) and phosphine ($PH_3$).
2. Intermolecular forces: Boiling point is determined by the strength of intermolecular forces.
3. Hydrogen Bonding in $NH_3$: Due to the high electronegativity and small size of the nitrogen atom, $NH_3$ molecules can form strong intermolecular hydrogen bonds.
4. Forces in $PH_3$: Phosphorus is less electronegative and larger than nitrogen. The P-H bond is not polar enough for hydrogen bonding. The dominant intermolecular forces in $PH_3$ are weaker van der Waals forces (dipole-dipole and dispersion forces).
5. Comparison: The hydrogen bonds in $NH_3$ are much stronger than the van der Waals forces in $PH_3$. Consequently, more energy is required to overcome these forces in $NH_3$, resulting in a significantly higher boiling point. The boiling point of $NH_3$ is $-33.34 ^\circ C$, while that of $PH_3$ is $-87.7 ^\circ C$.
6. Conclusion: $PH_3$ boils at a lower temperature than $NH_3$. This statement is correct.

Statement D: The $N-N$ single bond is stronger than the $P-P$ single bond

1. Analyze the bonds: We are comparing the bond strength of a nitrogen-nitrogen single bond and a phosphorus-phosphorus single bond.
2. Lone Pair Repulsion: The nitrogen atom is very small. In an N-N single bond, the lone pair of electrons on each adjacent nitrogen atom are close to each other. This results in significant lone pair-lone pair repulsion, which weakens the bond.
3. Comparison with P-P bond: The phosphorus atom is larger than the nitrogen atom. The distance between the lone pairs on adjacent phosphorus atoms in a P-P single bond is greater, leading to much weaker lone pair-lone pair repulsion.
4. Bond Enthalpy Data: The experimental bond enthalpy of the N-N single bond is approximately 160 kJ/mol, whereas the bond enthalpy of the P-P single bond is approximately 200 kJ/mol.
5. Conclusion: The P-P single bond is stronger than the N-N single bond. This is a known anomaly in Group 15. Therefore, the statement is incorrect.

Final Result: The correct statements are A, B, and C.