dipole–dipole interactions: OCR Sample Paper wrong?
Permanent Dipoles occur when two atoms bonded to each other have significantly different electronegativities. When the atomic number of an element increases the number of electrons present in them also increases. These electrons are constantly moving around in the atom. At some point, there are more electrons on one side of the atom than the other. This creates an imbalance of charge, creating a small charge of $\delta ^+$ and $\delta ^-$. These small charges induce its neighbors to create additional charges of $\delta ^+$ and $\delta ^-$, thus creating a chain of induced dipoles. These induced dipoles are temporary as the electrons can move, and hence can be reverted. They are, therefore, known as induced dipole forces, or better, known as London Dispersion Forces. In reality, they are not really forces, but rather attraction between molecules.

Naturally, more the number of electrons in the atom, more the amount of London dispersion forces, and hence greater the attraction. The greater the attraction, the more energy you have to put to break the bonds apart. Between the considered hydrogen halides ($\ce{HBr}$ and $\ce{HI}$), iodine has more electrons than bromine, and hence $\ce{HI}$ has a greater attraction force between molecules than $\ce{HBr}$. As a consequence, $\ce{HI}$ has a greater boiling point between them. Appropriate values are listed in following table:

$$\begin{array}{cccccc}
\hline
\ce{H-X} &amp; \text{BP}^\text{Ref.1} &amp; \ce{X}\text{'s EN} &amp; \mu ^\text{Ref.2} &amp; \text{Polarizability }(\pu{10^{-24} cm^3})  &amp; (r_\ce{H-X})^\text{Ref.3} \\ \hline
\ce{HCl} &amp; \pu{-85.1 ^\circ C} &amp; 3.16 &amp; \pu{1.09 D} &amp; 2.63 &amp; \pu{1.27 Å} \\
\ce{HBr} &amp; \pu{-67.1 ^\circ C} &amp; 2.96 &amp; \pu{0.82 D} &amp; 3.61 &amp; \pu{1.42 Å} \\
\ce{HI} &amp; \pu{-35.1 ^\circ C} &amp; 2.66 &amp; \pu{0.45 D} &amp; 5.44 &amp; \pu{1.60 Å} \\ 
\hline
\end{array}$$

Moreover, a related question you can find here. Orthocresol has given very reliable answer, which included two equations to calculate the strength of permanent dipole interactions ($U_\mathrm{dipole}$) and the strength of dispersion interactions ($U_\mathrm{dispersion}$). The author has stated that:


  If you plug all the numbers in for the hydrogen halides $\ce{HX}$, you will find that the magnitude of the dispersion forces is much larger than the magnitude of the dipole-dipole attractions, i.e. $\left|U_{\mathrm{dispersion}}\right| \gg \left|U_{\mathrm{dipole}}\right|$.


For example, for $\ce{HCl}$, dispersion forces contribute $86\%$ to the intermolecular attractions, and for $\ce{HI}$, they contribute $99\%$.



References:


J. F. Ogilvie, W. R. Rodwell, R. H. Tipping, “Dipole moment functions of the hydrogen halides,” J. Chem. Phys. 1980, 73(10), 5221 – 5229 (https://doi.org/10.1063/1.439950).
N. N. Greenwood, A. Earnshaw, “Chapter 17 – The Halogens: Fluorines, Chlorinee, Bromine, Iodine, and Astatine,” In Chemistry of the Elements; 2nd Edn.; Elsevier: Boston, MA, 1997, pp. 789–887. 
K. Fajans, N. Bauer, “Electronic Structure and Stability of Hydrogen Halides and of Complex Ions $\ce{XO4}$,” J. Chem. Phys. 1942, 10(7), 410 – 415 (https://doi.org/10.1063/1.1723742).