I don't really think the comment about context-dependency answers your question. They're right that energy is measured relatively, but that doesn't change the fact that there is an objective standard for measuring bond energies, and that definition is the one you're most likely to encounter as you're looking up this topic. They've also missed an important point: what we usually cite as the "bond strength" of an ionic interaction is actually the strength of multiple ionic interactions, while that for a covalent bond is for one discrete bond.

Context is still important though - namely, the context in which ionic "bonds" form compared to that in which covalent bonds do. Covalent bonding is a fundamentally quantum mechanical phenomenon that occurs between pairwise electron interactions. This means that if two electrons participate in one bond, they cannot form another without that first one being broken. Bonds do interact, and electrons do move around, but we can still model a given molecule as having a discrete number of covalent bonds. Consequently, each atom has a characteristic maximum amount of covalent bonds that it can (stably) form.

Ionic "bonds", however, are longer-range electrostatic interactions that do not involve significant overlap of electron energy orbitals. You don't need to invoke quantum mechanics to understand ionic bonding (at least at a basic level), since ionic interactions can be accurately modeled as classic charged-particle interactions of the sort you learn in an introductory physics course on electromagnetism. In contrast to the situation of covalent bonding, the number of ionic "bonds" that an ion can participate in is limited only by the available space around it. This is why we do not refer to ionic solids as "molecules" - in an ionic lattice, there is no clear way to define where one "molecule" ends and another begins. The formula "NaCl" is a molar ratio expressing the fact that a lattice of sodium and chloride will have a 1:1 ratio of those two ions.

How, then, should we describe the strength of a typical sodium-chloride interaction? The most common method is not to describe the potential energy between a single sodium ion and a single chloride ion in a vacuum, but instead to talk about the energy present per mole of lattice. Note that this is different from how we define covalent bond energies, which is per mole of bond. We define the molar lattice energy as the energy required to completely dissociate a mole of an ionic solid. For NaCl, this value is 786 kJ/mol (though this is usually expressed as a negative number, reflecting the energy liberated when the lattice forms from gaseous components), which is higher in magnitude than all but the strongest covalent bonds. It's certainly higher than any single covalent bonds, but it's not really reflecting the strength of one bond - it's more like six bonds added together. Each of the Na-Cl electrostatic interactions is weaker than the typical covalent bond, but the summed strength of all the electrostatic interactions between a sodium ion and the six chloride ions immediately surrounding it in a sodium chloride lattice is far higher.

Note that I'm simplifying a lot here. Ionic and covalent interactions are both emergent phenomena of the same underlying quantum-mechanical processes. These are human-made definitions, and there may be ambiguous edge-cases between them. But if we take concrete examples or ask specific questions like, "If ionic bonds are said to be stronger than covalent bonds, then why are ionic bonds (but not covalent ones) often easily broken in water?" we can find some relevant differences between these bonding behaviors that allow us to answer those questions.