As [Paul has pointed out][1], there are many different ways of defining atomic radii.
[ron has also pointed out][2] that the trend goes in the *opposite* direction from what you said.
For instance, according to [Ptable][3],

    |Calculated radii |
    |  |r(pm)|  |r(pm)|
    |H | 53  |He| 31  |
    |F | 42  |Ne| 38  |
    |Cl| 79  |Ar| 71  |
    |Br| 94  |Kr| 88  |

I have added hydrogen here as the natural "halogen" for helium.
Empirical and covalent also follow the same trend as above.
But the van der Waals radii, as shown below, do not,

    |    vdW radii    |
    |  |r(pm)|  |r(pm)|
    |H | 120 |He| 140 |
    |F | 147 |Ne| 154 |
    |Cl| 175 |Ar| 188 |
    |Br| 185 |Kr| 202 |

I am assuming this is the data set you've mentioned.
The reason for the disagreement is that electron clouds may relax or contract depending on their chemical environments and so do atomic radii.
Thus, what this shows is that:

1. H, F, Cl and Br are **smaller** than He, Ne, Ar and Kr when those atoms **are not** bonded;
2. H, F, Cl and Br are **larger** than He, Ne, Ar and Kr when those atoms **are** bonded.

  [1]: http://chemistry.stackexchange.com/a/66588/40029
  [2]: https://chemistry.stackexchange.com/questions/66587/why-does-being-bonded-decrease-atomic-size#comment114980_66587
  [3]: http://www.ptable.com/
  [4]: http://dx.doi.org/10.1021/j100785a001