Ionic, Covalent and Metallic Hydrides — Explained

On this page

1. Conceptual Foundation: Hydrogen's Bonding Versatility
2. I. Ionic Hydrides (Saline or Salt-like Hydrides)
3. Formation and Characteristics:
4. II. Covalent Hydrides (Molecular Hydrides)
5. Formation and Characteristics:
6. Classification of Covalent Hydrides (based on electron count around the central atom):
7. III. Metallic Hydrides (Interstitial Hydrides)
8. Formation and Characteristics:
9. Common Misconceptions:
10. NEET-Specific Angle:

Detailed Explanation

Hydrogen, with its unique electronic configuration ($1s^1$), occupies a special position in the periodic table, exhibiting properties akin to both alkali metals (by losing an electron to form $H^+$) and halogens (by gaining an electron to form $H^-$).

This dual nature allows hydrogen to form a diverse array of binary compounds called hydrides, where it combines with nearly all other elements. The classification of hydrides into ionic, covalent, and metallic types is fundamentally based on the electronegativity difference between hydrogen and the bonding element, which dictates the nature of the chemical bond formed.

Conceptual Foundation: Hydrogen's Bonding Versatility

Hydrogen's electronegativity value is approximately 2.20 on the Pauling scale. This value is intermediate, meaning it can form different types of bonds:

1
Ionic bond (as $H^-$ anion) — When hydrogen combines with highly electropositive elements (electronegativity much lower than 2.20), it gains an electron to form a hydride ion ($H^-$). This typically occurs with Group 1 and Group 2 metals.
2
Covalent bond — When hydrogen combines with elements having similar or higher electronegativity (electronegativity close to or higher than 2.20), it shares electrons to form a covalent bond. This is characteristic of p-block elements.
3
Metallic/Interstitial bond — With transition metals and inner transition metals, hydrogen atoms can occupy interstitial sites within the metal lattice, forming non-stoichiometric compounds that retain metallic properties.

I. Ionic Hydrides (Saline or Salt-like Hydrides)

Formation and Characteristics:

Ionic hydrides are formed by the most electropositive elements, specifically Group 1 (alkali metals) and Group 2 (alkaline earth metals, except Be and Mg which form polymeric covalent hydrides). In these compounds, hydrogen acts as a hydride ion ($H^-$), which is a strong reducing agent and a strong base. The general formula is $MH$ for Group 1 metals and $MH_2$ for Group 2 metals.

Bonding — Predominantly ionic, involving the complete transfer of an electron from the metal to hydrogen. For example, in NaH, it exists as $Na^+H^-$.
Physical Properties — They are typically white, crystalline solids at room temperature, resembling ionic salts. They have high melting points and boiling points due to strong electrostatic forces between ions. They are non-volatile and non-conductive in the solid state. However, in the molten state or when dissolved in suitable non-aqueous solvents (like molten salts), they conduct electricity, and hydrogen gas is liberated at the anode, confirming the presence of $H^-$ ions.
Chemical Properties

* Reactivity with Water: Ionic hydrides react vigorously and exothermically with water to produce hydrogen gas and a strong base. This reaction is often violent due to the strong basicity of $H^-$ and the reducing nature of the hydride ion.

$$NaH(s) + H_2O(l) \rightarrow NaOH(aq) + H_2(g)$$

$$CaH_2(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + 2H_2(g)$$

* Reducing Nature: They are powerful reducing agents. For instance, lithium aluminium hydride ($LiAlH_4$) and sodium borohydride ($NaBH_4$), though complex hydrides, derive their reducing power from the hydride ion.

Simple ionic hydrides can reduce metal oxides to metals.

$$2NaH(s) + TiO_2(s) \rightarrow Ti(s) + 2NaOH(s)$$

* Reaction with Acids: They react with acids to liberate hydrogen gas.

$$NaH(s) + HCl(aq) \rightarrow NaCl(aq) + H_2(g)$$

* Thermal Stability: Generally stable at high temperatures, but some decompose upon strong heating (e.

g., $LiH$ is very stable, $CaH_2$ also stable, but $BaH_2$ less so).

II. Covalent Hydrides (Molecular Hydrides)

Formation and Characteristics:

Covalent hydrides are formed by hydrogen combining with p-block elements (e.g., C, N, O, F, Si, P, S, Cl) and some s-block elements like Be and B. These compounds consist of discrete molecules where hydrogen shares electrons with the other element via covalent bonds. Their physical state (gas, liquid, or solid) depends on the intermolecular forces (van der Waals forces, dipole-dipole interactions, hydrogen bonding).

Classification of Covalent Hydrides (based on electron count around the central atom):

1
Electron-deficient Hydrides — These hydrides have too few electrons to form conventional covalent bonds (where each bond requires two electrons). A classic example is boranes, like diborane ($B_2H_6$). Boron has only three valence electrons, so $BH_3$ (which exists as a dimer $B_2H_6$) is electron-deficient. They often form 'bridge bonds' or 'three-center two-electron bonds' to compensate for the electron deficiency. They act as Lewis acids (electron acceptors).
2
Electron-precise Hydrides — These hydrides have the exact number of electrons required to form conventional covalent bonds, with no lone pairs on the central atom. Methane ($CH_4$) is the prime example. The central atom achieves a stable octet (or duet for H) through sharing. They are generally tetrahedral in geometry (if the central atom is sp3 hybridized) and non-polar.
3
Electron-rich Hydrides — These hydrides have excess electrons present as lone pairs on the central atom. Examples include ammonia ($NH_3$), water ($H_2O$), and hydrogen fluoride ($HF$). The presence of lone pairs influences their molecular geometry (e.g., pyramidal for $NH_3$, bent for $H_2O$) and often leads to strong intermolecular forces like hydrogen bonding, resulting in higher melting and boiling points than expected based on molecular weight alone.

Physical Properties — Can be gases ($CH_4$, $NH_3$, $HCl$), liquids ($H_2O$), or low-melting solids ($H_2S$). Their volatility is generally higher than ionic hydrides. Many are soluble in organic solvents.
Chemical Properties

* Acidity/Basicity: Varies widely. $HCl$, $HBr$, $HI$ are strong acids. $H_2O$ is amphoteric. $NH_3$ is a weak base. $CH_4$ is largely unreactive. * Reactivity with Water: Electron-rich hydrides like $NH_3$ and $H_2O$ are miscible with water.

Some covalent hydrides like $SiH_4$ can hydrolyze slowly. Halogen hydrides like $HCl$ dissolve in water to form acids. * Lewis Acid/Base Behavior: Electron-deficient hydrides (e.g., $BH_3$) act as Lewis acids.

Electron-rich hydrides (e.g., $NH_3$, $H_2O$) act as Lewis bases due to their lone pairs.

III. Metallic Hydrides (Interstitial Hydrides)

Formation and Characteristics:

Metallic hydrides are formed by many d-block and f-block elements (transition metals and inner transition metals). These are often referred to as interstitial hydrides because hydrogen atoms occupy the interstitial sites (voids) within the crystal lattice of the metal. The formation of these hydrides is complex and not fully understood, but it involves the absorption of hydrogen by the metal.

Bonding — The bonding is not purely ionic or covalent. It's more akin to metallic bonding, where the hydrogen atoms are thought to donate their electrons to the metal's conduction band, or exist as protons within the electron sea, or even as hydride ions in some cases. The exact nature is still debated.
Stoichiometry — A defining feature is their non-stoichiometric nature, meaning the ratio of hydrogen to metal is not a simple whole number (e.g., $TiH_{1.5-1.8}$, $PdH_{0.6-0.8}$). This is because the number of interstitial sites occupied by hydrogen can vary. However, some transition metals (e.g., Group 6-8 metals like Cr, Mn, Fe, Co, Ni) do not form hydrides under normal conditions, a region known as the 'hydride gap'.
Physical Properties — They generally retain the metallic luster, hardness, and electrical conductivity of the parent metals, though these properties might be slightly altered. They are typically hard, greyish-black solids.
Chemical Properties

* Hydrogen Storage: Many metallic hydrides can absorb large volumes of hydrogen and then release it upon heating, making them promising materials for hydrogen storage and transport. For example, $LaNi_5H_6$ can store a significant amount of hydrogen.

* Catalytic Activity: Some metallic hydrides, like palladium hydride, are used as catalysts in hydrogenation reactions. * Reducing Nature: They can act as reducing agents, though generally less reactive than ionic hydrides.

* Thermal Stability: Their stability varies, with hydrogen being released upon heating.

Common Misconceptions:

1
All hydrides are basic — Incorrect. While ionic hydrides are strongly basic, covalent hydrides can be acidic ($HCl$), neutral ($CH_4$), or weakly basic ($NH_3$).
2
Metallic hydrides are true compounds with fixed stoichiometry — Incorrect. Most metallic hydrides are non-stoichiometric, meaning their composition is variable within a range.
3
Hydrogen always acts as $H^-$ in hydrides — Incorrect. Only in ionic hydrides does hydrogen exist as $H^-$. In covalent hydrides, it shares electrons, and in metallic hydrides, its state is more complex.
4
All elements form hydrides — Incorrect. There is a 'hydride gap' in the d-block (Group 7, 8, 9) where elements do not form hydrides under normal conditions.

NEET-Specific Angle:

For NEET, the focus is often on comparative properties, reactivity patterns, and specific examples. Questions frequently test:

Classification — Identifying the type of hydride given its formula or properties.
Reactivity — Especially the reaction of ionic hydrides with water (producing $H_2$ and base) and the acidic/basic nature of covalent hydrides.
Physical properties — Melting points, conductivity, physical state, and how they relate to bonding.
Exceptions and Trends — For instance, BeH2 and MgH2 are considered polymeric covalent, not ionic, despite being Group 2 metals. The 'hydride gap' is also an important concept. The electron-deficient, precise, and rich classification of covalent hydrides is also a common area for questions.
Applications — Hydrogen storage using metallic hydrides.

Often confused with

Side-by-side differences the NEET paper likes to test.

Ionic, Covalent and Metallic Hydrides vs Ionic, Covalent, and Metallic Hydrides

Aspect	Ionic, Covalent and Metallic Hydrides	Ionic, Covalent, and Metallic Hydrides
Elements Involved	Ionic Hydrides (e.g., NaH, CaH2)	Covalent Hydrides (e.g., CH4, NH3, H2O)
Bonding Nature	Predominantly ionic (electron transfer, $M^+H^-$)	Covalent (electron sharing)
Physical State (at RT)	Crystalline solids, high melting points	Gases, liquids, or low-melting solids
Electrical Conductivity	Non-conductive in solid state; conductive in molten state	Non-conductive (insulators)
Stoichiometry	Stoichiometric (fixed, simple whole-number ratio)	Stoichiometric (fixed, simple whole-number ratio)
Reactivity with Water	Vigorous reaction, produces $H_2$ gas and strong base	Varies (e.g., $H_2O$ miscible, $HCl$ acidic, $CH_4$ inert)
Reducing/Oxidizing Nature	Strong reducing agents ($H^-$)	Varies (e.g., $NH_3$ weak reducing, $H_2O$ mild oxidizing/reducing)
Examples	$LiH, NaH, CaH_2, SrH_2$	$CH_4, NH_3, H_2O, HF, SiH_4, H_2S$

The three main types of hydrides—ionic, covalent, and metallic—are fundamentally distinguished by the nature of their chemical bonding, which in turn dictates their macroscopic properties and reactivity.

Ionic hydrides, formed with highly electropositive metals, involve electron transfer to hydrogen, resulting in salt-like solids with high melting points and strong reducing capabilities. Covalent hydrides, formed with p-block elements, involve electron sharing, leading to molecular compounds with diverse physical states and chemical behaviors (acidic, basic, or neutral).

Metallic hydrides, formed with d- and f-block metals, are characterized by hydrogen occupying interstitial sites, resulting in non-stoichiometric compounds that retain metallic properties and are crucial for hydrogen storage.

Why it is tested: For NEET, understanding these differences is critical for classification questions, predicting reactivity (especially with water), and identifying the physical properties associated with each hydride type. Questions often involve matching hydride types with their characteristic properties or explaining observed trends and exceptions (e.g., BeH2 being covalent).

Questions students ask

6 answered on this topic.

What is the primary factor determining whether a hydride is ionic, covalent, or metallic?

The primary factor is the electronegativity difference between hydrogen and the element it bonds with. If the element is highly electropositive (low electronegativity, like Group 1 and 2 metals), the bond will be predominantly ionic, forming $H^-$ ions.

If the electronegativity difference is small or the element is more electronegative than hydrogen (like p-block elements), the bond will be covalent. For transition and inner transition metals, hydrogen occupies interstitial sites, leading to metallic hydrides.

Why are ionic hydrides strong reducing agents?

Ionic hydrides contain the hydride ion, $H^-$, which has a strong tendency to donate its extra electron to achieve a stable $1s^0$ configuration (like Helium). This strong electron-donating ability makes $H^-$ a powerful reducing agent, capable of reducing other species by providing electrons. For example, it can reduce metal oxides or even water.

How do electron-deficient, electron-precise, and electron-rich covalent hydrides differ?

These classifications are based on the number of valence electrons available for bonding around the central atom. Electron-deficient hydrides (e.g., $B_2H_6$) have fewer electrons than needed for conventional bonds, often forming multi-center bonds.

Electron-precise hydrides (e.g., $CH_4$) have exactly enough electrons to form conventional bonds and complete the central atom's octet. Electron-rich hydrides (e.g., $NH_3$, $H_2O$) have surplus electrons as lone pairs on the central atom, which can participate in hydrogen bonding or act as Lewis bases.

What is the 'hydride gap' in the periodic table?

The 'hydride gap' refers to the observation that elements of Group 7, 8, and 9 (Mn, Fe, Co, Ni, etc.) in the d-block do not form hydrides under normal conditions. This is an exception to the general trend of transition metals forming metallic hydrides. The reasons are complex, often attributed to unfavorable thermodynamic or kinetic factors for hydrogen absorption in these specific metals.

Can metallic hydrides conduct electricity? If so, how do their conductivity compare to the parent metals?

Yes, metallic hydrides generally retain their metallic character and can conduct electricity. This is because the hydrogen atoms occupy interstitial sites within the metal lattice, and the valence electrons of the metal remain delocalized, forming an 'electron sea'. However, their electrical conductivity is often lower than that of the parent metals. The absorbed hydrogen atoms can introduce scattering centers or alter the electronic band structure, slightly impeding electron flow.

Why are BeH2 and MgH2 considered covalent, despite Be and Mg being Group 2 metals?

While most Group 2 metals form ionic hydrides, Beryllium and Magnesium are exceptions. Beryllium has a very small size and high polarizing power, leading to significant covalent character in its bonds. Magnesium also exhibits some covalent character in its hydride. Both $BeH_2$ and $MgH_2$ exist as polymeric structures with bridging hydrogen atoms, consistent with covalent bonding rather than discrete $M^{2+}H_2^-$ ionic units. This is a crucial point for NEET aspirants.