Given:

5 moles of liquid X

10 moles of liquid Y

The total moles in the solution are 15 (5 moles X + 10 moles Y).

The mole fraction of X ($X_x$) is:

$ \frac{5}{15} = \frac{1}{3} $

The mole fraction of Y ($X_y$) is:

$ \frac{10}{15} = \frac{2}{3} $

The vapor pressures of pure liquids are:

$P_x^{\circ} = 63$ torr for liquid X

$P_y^{\circ} = 78$ torr for liquid Y

Using Raoult's Law for an ideal solution, the total vapor pressure ($P_{\text{total}}$) is calculated as:

$ P_{\text{total}} = X_x \cdot P_x^{\circ} + X_y \cdot P_y^{\circ} $

Substitute the values:

$ P_{\text{total}} = \left(\frac{1}{3} \times 63\right) + \left(\frac{2}{3} \times 78\right) $

Calculate each term:

$ = 21 + 52 $

$ = 73 \text{ torr} $

The observed vapor pressure of the solution is 70 torr, which is lower than the calculated ideal pressure of 73 torr. This indicates a negative deviation from Raoult's Law because the observed vapor pressure is less than expected for an ideal solution. Therefore, the solution exhibits stronger intermolecular attractions than those in an ideal solution, leading to a lower vapor pressure.