According to Fajan's rule ionic character should increase down the group as the size of cation increase.
So it must be

$$\ce{LiH < NaH < KH < RbH < CsH}$$

However, the following two sources provide two different opinions. Which one is right?

Source 1

Explanation: As the size of the metal cation increases down the group, the lattice energy of the hydrides decreases down the group. Consequently, the reactivity of the alkali towards hydrogen decreases down the group.

iii. The ionic character of the alkali metal hydrides increases from Li to Cs.

Explanation: Since ionisation enthalpy of alkali metals decreases down the group, tendency to form cations as well as the ionic character of the hydrides increases.

Source 2

But it we are asked to compare the ionic character of hydrides (LiH, NaH, KH, RbH, CsH), we can not use the Fajan's rule. This is because, here the anion is very small ($\ce{H-}$) as compared to the cations (alkali metal cations). The ionic character is this case can thus be compared on the basis of Lattice energy.

Since lithium ion and the hydride ion are approximately of same size, the charge density is high between the two, thus the ionic bond so formed will be very strong and hence more will be the lattice energy.

On coming down the group, the difference in the size of cation and anion increases, thereby making the ionic bond weaker, thus reducing the lattice energy.

Also, lesser will be the lattice energy, lesser will be the ionic character.

Thus, down the group, the ionic character of the hydrides decreases.

  • 1
    $\begingroup$ Just a side note: it's preferable to post text as, well, text, and not as images; also it's a good habit to provide references to the external data sources. $\endgroup$ – andselisk Jan 15 '19 at 12:58

Well, the second source is certainly more wrong. There is a major difference between ionic character of a bond and the stability (lattice energy) of an ionic lattice. $\ce{LiH}$ is certainly more covalent than $\ce{CsH}$ (meaning $\ce{CsH}$ is more ionic), which can be justified simply by comparing charge densities on $\ce{Li}$, which is very high and draws electron density from the $\ce{H}$ atom, forming a more covalent bond.

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It is true that the ionic lattice energy goes down as we move from $\ce{LiH}$ to $\ce{CsH}$. But what happens to the competing mechanism of covalent bonding?

To make a covalent bond you overlaporbutals of the atoms you're bonding, but in $\ce{CsH}$ you have a huge, fluffy, low effective nuclear charge cesium $6s$ orbital meeting a compact hydrogen $1s$ orbital. This is not going to go over well; the shoe really does not fit here. Lithium offers much better covalent overlap woth its $2s$ orbital instead of $6s$, with the other (stable) alkali metals in between.

So yes, ionic bonding energy goes down as we proceed to heavier alkali metal hydrides, but covalent bonding fares even worse. Heavier alkali metal hydrides have less overall bonding energy and a greater part of what bonding energy they have cones from the ionic interaction.

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