Why are rates of reaction faster for ionic compounds than for covalent compounds?
Not all ionic reactions are not faster, i.e. have a larger rate constant than those between neutral species, however, at similar activation energies there are effects that only ionic reactions suffer. There are two considerations, the dielectric constant of the solvent used, water, hexane, ethanol, acetone etc., and the relative charges on the ions, similar or opposite as well as their value, $\pm1,\pm 2 $ etc.

For slow reactions, that is not diffusion controlled, (see below) the 'primary kinetic salt effect' is important. This shows that the rate constant can increase or decrease depending on the ionic strength of the solution and the charge on the ions. It is based on the Debye-Huckel theory of ions in solution which relates activity coefficients in dilute solution to the ionic strength.

The rate constant for a bimolecular reaction  has the form 
$$ \log_{10}\left(\frac{k_2}{k_2^0}\right) = 2\alpha z_az_b\sqrt I$$
where I is the ionic strength, the z the charges on the two ions, $(-1,+1),(-1,-1) $ etc. and $\alpha \approx 1.6 \cdot 10^{-2} \pu{ m^{3/2} mol^{-1/2}}$ is a constant originating in Debye-Huckel theory. At low ionic strength ($ \lt 0.5 \pu{mol dm^{-3}}$ ) the rate constants increase or decrease linearly with $\sqrt I$.

The effect of increasing the dielectric constant (relative permittivity) is to lower the rate constant, thus rate constants are lower in water ($\epsilon = 78$) than in cyclohexane ($ \epsilon = 2.2$) for the same reaction. This is due to the fact that, at a given distance, in a high dielectric solvent the electric field of the ion is greatly reduced compared to its value in a low dielectric solvent. Thus in a high dielectric the ions are far more isolated from one another and so rate constants are only affected when ions come in relatively close contact. Conversely in low dielectric solvents the ions 'feel' the presence of one another at a large separation, this can considerably increase or decrease the rate constant depending on relative charges.

The rate constant is given by 
$$ \ln(k_2) =\ln(k_2^{\infty}) - \frac{z_az_bq^2 }{4\pi \epsilon _0 \epsilon d_{ab}k_BT}$$
where $k_2^{\infty}$ is the limiting rate constant for a solvent of infinite dielectric constant. The permittivity of free space is $\epsilon _0$, the relative permittivity $\epsilon$, q is the charge on the electron, $d_{ab}$ the separation of the ions at reaction and $k_B$ the Boltzmann constant. A plot of the rate constant vs. $1/\epsilon$ yields a reasonably straight line  for all but the lowest $\epsilon$ values. 

When a reaction's activation energy  is low, less than $\approx 10 \pu{kJmol^{-1}}$ the reactants react so rapidly on contact that the rate constant in determined by how quickly they can approach one another in solution; the rate constant then becomes diffusion controlled and thus depends on the solution viscosity.

Using the Stokes-Einstein equation and assuming equal sized molecules the rate constant has the simple form 
$$ k_2= \frac{8RT}{3\eta}$$
where $\eta$ is the solvent viscosity. Water has a value $\approx 1\cdot 10^{-3} $ Pa s or $1$ centipoise (cP) at room temperature. In fluid solution such as hexane the rate diffusion controlled constant is typically $10^{10}$ to $10^{11} \pu{dm^3mol^{-1}s^{-1}}$.

The effects of ions changes the diffusion controlled rate constant for the reasons described above and is now 
$$ k_2= \frac{8RT}{3\eta}\left(\frac{x}{e^x -1}\right)$$
where 
$$ x=\frac{z_az_bq^2 }{4\pi \epsilon _0 \epsilon d_{ab}kT}$$
and is the term used before.

When the ionic charges are the same and the dielectric constant small the rate constant is greatly reduced, and conversely it is enhanced when the charges are opposite in sign to one another. In high dielectric solvents these changes are greatly reduced. These effects are similar to those described above.