Give an example of a solid solution in which the solute is a gas.

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Q. Give an example of a solid solution in which the solute is a gas.

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Concept understanding — Types Of Solutions

Types of Solutions: From Everyday Life to Chemistry

You already know what a solution is — sugar dissolved in water, salt in water, even the air you breathe. But not all solutions behave the same way. Some dissolve easily, some refuse to dissolve beyond a point, and some can hold more solute than they normally should. That difference is what we classify as types of solutions based on how much solute is dissolved.

The Intuition: A Cup of Tea

Imagine making a cup of tea. You add one spoon of sugar — it dissolves completely. You add a second spoon — still dissolves. A third spoon — maybe it dissolves, maybe it doesn't. At some point, no matter how much you stir, the sugar just sits at the bottom.

That moment — when no more sugar dissolves — is the saturation point. Before that, you have an unsaturated solution. At that exact point, you have a saturated solution. And if you carefully heat the tea, dissolve more sugar, then cool it down without disturbing it — you might get a supersaturated solution, where more sugar stays dissolved than should be possible at that temperature.

That's the entire idea. Three types, defined by how much solute is dissolved relative to the maximum possible.

The Precise Statement

A solution is a homogeneous mixture of a solute (the substance being dissolved) and a solvent (the substance doing the dissolving). Based on the amount of solute dissolved relative to its solubility at a given temperature, solutions are classified into three types:

Types of Solutions (by saturation) Unsaturated solution — contains less solute than the maximum that can be dissolved at that temperature. Saturated solution — contains exactly the maximum amount of solute that can be dissolved at that temperature. Supersaturated solution — contains more solute than the maximum normally possible at that temperature (a metastable state).

Breaking Down Each Type

Unsaturated solution — the most common type. You can still add more solute and it will dissolve. The concentration is below the solubility limit. If you have a glass of water at room temperature and add a pinch of salt, you get an unsaturated solution. Add more salt — still unsaturated, until you hit the limit.

Saturated solution — the solute and undissolved solute are in dynamic equilibrium. At the molecular level, the rate at which solute particles dissolve equals the rate at which they crystallize out. No net change. If you keep adding salt to water and it stops dissolving, the liquid above the undissolved salt is a saturated solution. The concentration is fixed at the solubility value for that temperature.

Watch out

A common mistake: thinking a saturated solution is always "thick" or "concentrated." Not true. Saturation depends on the solute's solubility. Lead(II) chloride saturates at about 0.45 g per 100 mL water — that's a very dilute saturated solution. Saturation ≠ high concentration.

Supersaturated solution — this is a tricky one. You create it by heating the solvent, dissolving more solute than normally possible, then carefully cooling it. The excess solute stays dissolved because there's no nucleation site (no scratch, no dust particle) to trigger crystallization. It's unstable — the slightest disturbance (a dust speck, a scratch on the glass, even a sudden jolt) causes the excess solute to crystallize out instantly.

Tip

Supersaturated solutions are the reason "hot ice" (sodium acetate) hand warmers work. You click a metal disc inside, which creates a nucleation site, and the entire solution crystallizes in seconds, releasing heat.

A Quick Comparison

Type	Solute amount vs. solubility	Can more solute dissolve?	Stability
Unsaturated	Less than maximum	Yes	Stable
Saturated	Equal to maximum	No (at equilibrium)	Stable
Supersaturated	More than maximum	No (excess will crystallize)	Metastable

Why This Matters

In exams, you'll often be asked to identify the type of solution from a given scenario — like "50 g of salt dissolved in 100 g water at 30°C, given solubility is 36 g per 100 g water." That's a supersaturated solution (50 > 36). Or you might be asked what happens when you add a seed crystal to a supersaturated solution — it triggers crystallization.

The key is always: compare the actual amount dissolved to the solubility at that temperature. That single comparison gives you the type.

Important

Solubility is temperature-dependent. A solution that is saturated at 20°C becomes unsaturated if heated to 50°C (because solubility usually increases with temperature). Always check the temperature condition given in the problem.

Searches such as "types of solutions saturated unsaturated supersaturated" and "solutions class 12 chemistry notes" align directly with the Solutions chapter of the NCERT/CBSE Class 12 Chemistry curriculum. Identifying which type a given scenario describes is a common short-answer question in board exams.

Why this formula?

Types of Solutions: Why the Key Formulae Hold

Understanding why the formulae work is essential for Indian exams (JEE, NEET, CBSE). Let's break down the reasoning behind the most important relationships.

1. The Basic Classification: What Makes a Solution?

A solution is a homogeneous mixture of two or more substances. The key idea is intermolecular forces between solute and solvent particles.

Ideal Solution: Solute-solvent interactions are identical to solute-solute and solvent-solvent interactions. Why? No net energy change on mixing — the molecules "fit" perfectly.
Non-Ideal Solution: Interactions differ, leading to deviation from Raoult's law.

2. Raoult's Law: The Foundation

Formula:

P_{\text{solution}} = x_{\text{solvent}} \cdot P_{\text{solvent}}^0

Why does this hold?

Imagine a pure solvent surface. The vapour pressure $P^0$ comes from molecules escaping the liquid. When you add a non-volatile solute, solute molecules occupy some surface area, blocking solvent molecules from escaping.

The fraction of surface available to solvent = mole fraction of solvent ( $x_{\text{solvent}}$ ).
Therefore, the rate of escape (vapour pressure) is proportional to that fraction:

P_{\text{solution}} \propto x_{\text{solvent}}

At the limit $x_{\text{solvent}} = 1$ , $P_{\text{solution}} = P^0$ , so the constant is $P^0$ .

Key insight: Raoult's law is a surface-area argument, not a volume argument.

3. Relative Lowering of Vapour Pressure

Formula:

\frac{P^0 - P}{P^0} = x_{\text{solute}}

Derivation in one line:

From Raoult's law:

P = x_{\text{solvent}} \cdot P^0

Since $x_{\text{solvent}} + x_{\text{solute}} = 1$ ,

P = (1 - x_{\text{solute}}) P^0

\Rightarrow P^0 - P = x_{\text{solute}} \cdot P^0

\Rightarrow \frac{P^0 - P}{P^0} = x_{\text{solute}}

Why is this useful?

It depends only on the mole fraction of solute, not on its identity — making it a colligative property.

4. Elevation of Boiling Point

Formula:

\Delta T_b = K_b \cdot m

Why does boiling point rise?

Boiling occurs when vapour pressure = atmospheric pressure.
Adding a non-volatile solute lowers vapour pressure (Raoult's law).
To reach atmospheric pressure again, you must raise the temperature.
The shift $\Delta T_b$ $Δ T_{b}$ is proportional to the molality $m$ $m$ (moles of solute per kg of solvent), because:
- More solute → greater vapour pressure lowering → more temperature needed.
- $K_b$ (ebullioscopic constant) is a property of the solvent only.

5. Depression of Freezing Point

Formula:

\Delta T_f = K_f \cdot m

Why does freezing point drop?

At the freezing point, solid and liquid solvent are in equilibrium.
Adding solute disrupts this equilibrium — solute molecules interfere with the orderly crystal formation of the solvent.
To re-establish equilibrium, you must lower the temperature.
Again, $\Delta T_f \propto m$ , and $K_f$ depends only on the solvent.

Common exam trap: Both $\Delta T_b$ and $\Delta T_f$ are colligative — they depend on number of solute particles, not their nature.

6. Osmotic Pressure

Formula:

\Pi = i \cdot C \cdot R \cdot T

Why does this hold?

Osmosis is the net movement of solvent from low solute concentration to high solute concentration across a semipermeable membrane.
The solvent moves to dilute the higher concentration — this is a entropy-driven process (mixing increases disorder).
Osmotic pressure $\Pi$ is the external pressure needed to stop this flow.
It behaves like an ideal gas law for solute particles:

\Pi V = nRT \quad \Rightarrow \quad \Pi = \frac{n}{V}RT = CRT

The van't Hoff factor $i$ accounts for dissociation/association of solute (e.g., NaCl gives $i \approx 2$ ).

7. The van't Hoff Factor $i$

Formula:

i = \frac{\text{observed colligative property}}{\text{expected colligative property}}

Why is $i$ needed?

Colligative properties depend on number of particles.
If a solute dissociates (e.g., $\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-$ ), the effective particle count doubles.
If it associates (e.g., benzoic acid in benzene forms dimers), the count halves.
$i$ corrects for this:

\Delta T_f = i \cdot K_f \cdot m

Quick Summary Table

Property	Formula	Why it works
Raoult's law	$P = x_{\text{solvent}} P^0$	Surface area blocking by solute
Relative lowering	$\frac{\Delta P}{P^0} = x_{\text{solute}}$	Direct algebraic consequence
Boiling point elevation	$\Delta T_b = K_b m$	Need higher temp to overcome vapour pressure drop
Freezing point depression	$\Delta T_f = K_f m$	Solute disrupts crystal formation
Osmotic pressure	$\Pi = iCRT$	Analogy to ideal gas law for solute particles

Final takeaway: Every formula in "Types of Solutions" flows from Raoult's law (for vapour pressure) and the particle-counting principle (for colligative properties). Understand these two roots, and you can reconstruct the rest.

When a gas dissolves into a solid matrix and forms a homogeneous mixture at the molecular level, we get a solid solution with a gaseous solute. Hydrogen in palladium is the classic example.

Understanding Solid Solutions with Gaseous Solutes

A solution doesn't have to be liquid. Any homogeneous mixture where one substance (the solute) is uniformly dispersed in another (the solvent) at the molecular or atomic scale qualifies as a solution, regardless of the final state of matter.

When we talk about a solid solution with a gas as the solute, we mean that gas molecules or atoms have been absorbed into the crystal lattice of a solid and distributed uniformly throughout. The gas doesn't just sit on the surface or in visible pockets—it occupies interstitial spaces (gaps between atoms) or substitutional sites in the solid's structure.

The Classic Example: Hydrogen in Palladium

The most important and widely cited example is palladium metal absorbing hydrogen gas.

The absorption process: When palladium (Pd) is exposed to hydrogen gas ( $\ce{H2}$ ), the metal has a remarkable ability to absorb large quantities of it. The hydrogen molecules dissociate into individual hydrogen atoms at the palladium surface.
Formation of the solid solution: These hydrogen atoms then diffuse into the palladium crystal lattice, occupying the octahedral interstitial sites—the small spaces between the palladium atoms. At room temperature, palladium can absorb up to 900 times its own volume of hydrogen gas.
Homogeneity: The hydrogen atoms distribute themselves uniformly throughout the palladium structure, creating a true solid solution. The resulting material is often written as $\ce{PdH_x}$ , where $x$ can vary depending on temperature and pressure.
Why this matters: This property makes palladium extremely useful for hydrogen storage, purification (palladium membranes selectively allow hydrogen to pass through), and catalysis. The hydrogen can be released by heating or reducing the pressure.

Note

Other metals also form solid solutions with hydrogen, including platinum, nickel, and certain alloys. However, palladium's capacity and reversibility make it the textbook example.

Other Examples

While hydrogen in palladium is the standard answer, you might also encounter:

Hydrogen in platinum ( $\ce{Pt}$ )
Hydrogen in titanium ( $\ce{Ti}$ ) or zirconium ( $\ce{Zr}$ )—these form metal hydrides that can be considered solid solutions at certain compositions

The key feature in all cases is that the gas atoms become part of the solid's structure, not merely trapped in voids or adsorbed on the surface.

✓Final answer

\ce{PdH_x}

Method: Classification by Physical State of Solute and Solvent

This method uses the physical states (solid, liquid, gas) of both components to identify the type of solution.

Steps

Identify the solvent — the component present in larger amount (the medium).
Identify the solute — the component present in smaller amount (dissolved substance).
Check the physical state of each at room temperature.
Classify the solution based on the combination of states.

Example: Solid solution with gaseous solute

Solute: Gas

Solvent: Solid

Real example: Hydrogen gas ( $H_2$ ) dissolved in palladium metal ( $Pd$ )

Solvent: Palladium (solid metal)
Solute: Hydrogen gas
Result: A solid solution where gas molecules occupy interstitial spaces in the metal lattice.

Key point: This is possible because palladium can absorb large volumes of hydrogen gas into its crystal structure, forming a homogeneous solid solution.

Other common examples:

Helium gas dissolved in platinum
Hydrogen dissolved in nickel

Quick Reference Table

Solute	Solvent	Solution Type	Example
Gas	Solid	Solid solution	$H_2$ in $Pd$
Gas	Liquid	Liquid solution	$CO_2$ in water (soda)
Gas	Gas	Gaseous solution	Air ( $O_2$ + $N_2$ )

Final answer: A solid solution with a gaseous solute is hydrogen gas dissolved in palladium metal.

Common Mistakes: Solid Solutions with a Gaseous Solute

Students often struggle with this concept because it challenges the everyday intuition that "solids can't dissolve gases." Here are the most frequent errors and how to avoid them.

✗ Mistake 1: Confusing "solute" with "solvent"

The error: Students list examples like "air" (gas in gas) or "salt water" (solid in liquid) — these are not solid solutions at all.

Why it happens: The phrase "solid solution" is misinterpreted as "a solution that is solid" rather than "a solution where the solvent is solid."

How to avoid: Always identify the solvent first. In a solid solution, the solvent (the majority component) must be a solid. The solute can be a gas, liquid, or another solid.

Key rule: Solid solution = solid solvent + any solute (gas, liquid, or solid).

✗ Mistake 2: Giving impossible examples

The error: Students write "ice cubes in air" or "sugar in a gas" — these are not solutions at all.

Why it happens: Lack of clarity about what constitutes a true solution (homogeneous mixture at molecular level).

How to avoid: Remember that a true solution is homogeneous. For a gas to dissolve in a solid, the gas molecules must fit into the interstitial spaces or lattice vacancies of the solid.

✗ Mistake 3: Forgetting the correct example

The error: Students cannot recall any valid example under exam pressure.

Why it happens: The only common example is specialized — hydrogen in palladium or hydrogen in platinum.

How to avoid: Memorise this one key example:

Correct answer: Hydrogen gas dissolved in palladium metal

(Palladium metal absorbs up to 900 times its own volume of hydrogen gas, forming a homogeneous solid solution.)

✓ Quick Reference Table

Component	Example	Valid?
Gas in solid	$\text{H}_2$ in Pd	✓ Yes
Gas in solid	$\text{O}_2$ in Pt	✓ Yes
Gas in solid	Air in ice	✗ No (not homogeneous)
Gas in solid	$\text{CO}_2$ in solid water	✗ No (forms clathrate, not true solution)

📝 Exam Tip

If asked to "give an example," write:

"Hydrogen gas ( $\text{H}_2$ ) dissolved in palladium metal (Pd) — a solid solution where the solute is a gas and the solvent is a solid."

This shows you understand both the concept and the terminology.

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Q.Give an example of a solid solution in which the solute is a gas.

Concept understanding — Types Of Solutions

Types of Solutions: From Everyday Life to Chemistry

The Intuition: A Cup of Tea

The Precise Statement

Breaking Down Each Type

A Quick Comparison

Why This Matters

Why this formula?

Types of Solutions: Why the Key Formulae Hold

1. The Basic Classification: What Makes a Solution?

2. Raoult's Law: The Foundation

Why does this hold?

3. Relative Lowering of Vapour Pressure

Derivation in one line:

4. Elevation of Boiling Point

Why does boiling point rise?

5. Depression of Freezing Point

Why does freezing point drop?

6. Osmotic Pressure

Why does this hold?

7. The van't Hoff Factor $i$

Why is $i$ needed?

Quick Summary Table

Solid Solution with Gaseous Solute

Understanding Solid Solutions with Gaseous Solutes

The Classic Example: Hydrogen in Palladium

Other Examples

Method: Classification by Physical State of Solute and Solvent

Steps

Example: Solid solution with gaseous solute

Quick Reference Table

Common Mistakes: Solid Solutions with a Gaseous Solute

✗ Mistake 1: Confusing "solute" with "solvent"

✗ Mistake 2: Giving impossible examples

✗ Mistake 3: Forgetting the correct example

✓ Quick Reference Table

📝 Exam Tip

Q.Give an example of a solid solution in which the solute is a gas.

Understanding Solid Solutions with Gaseous Solutes

The Classic Example: Hydrogen in Palladium

Other Examples

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