Amongst the following compounds, identify which are insoluble, partially soluble and highly soluble in water?

Exercises · 1.25

Q. Amongst the following compounds, identify which are insoluble, partially soluble and highly soluble in water?

(i) phenol

(ii) toluene

(iii) formic acid

(iv) ethylene glycol

(v) chloroform

(vi) pentanol.

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Concept understanding — Solubility Rules

Solubility Rules – From Intuition to Precision

Imagine you drop a spoonful of sugar into a glass of water. Stir once, and it disappears. Now try the same with a spoonful of sand. It just sits at the bottom. Why? The sugar molecules are able to break apart and mingle with water molecules — we say sugar dissolves in water. Sand does not.

That "disappearing" act is solubility. But in chemistry, we don't just ask if something dissolves — we ask how much and under what conditions. For ionic compounds (salts), the answer is surprisingly predictable. That predictability is what we call the Solubility Rules.

The Core Idea: "Like Dissolves Like" — But Ionic Compounds Are Different

For molecular substances like sugar, the rule of thumb is "like dissolves like" — polar dissolves in polar, non-polar in non-polar. But ionic compounds are made of charged particles (cations and anions). When you drop an ionic solid into water, the water molecules (which are polar) try to pull the ions apart. Whether they succeed depends on a tug-of-war:

The water molecules want to surround and separate the ions (hydration energy).
The ions themselves are held together by electrostatic forces (lattice energy).

If the hydration energy wins, the salt dissolves. If the lattice energy wins, it stays solid.

Note

You don't need to calculate these energies for exams. The Solubility Rules are a shortcut — a set of patterns discovered by observing thousands of salts.

The Precise Statement: The Solubility Rules

Here are the rules as you'll use them in exams. They are hierarchical — the first applicable rule overrides later ones.

Solubility Rules for Ionic Compounds in Water

Rule	Statement	Examples
1	All nitrates ( $\text{NO}_3^-$ ) are soluble.	$\text{NaNO}_3$ , $\text{AgNO}_3$ , $\text{Pb(NO}_3)_2$
2	All acetates ( $\text{CH}_3\text{COO}^-$ ) are soluble.	$\text{NaCH}_3\text{COO}$ , $\text{AgCH}_3\text{COO}$
3	All chlorides ( $\text{Cl}^-$ ), bromides ( $\text{Br}^-$ ), and iodides ( $\text{I}^-$ ) are soluble, except with $\text{Ag}^+$ , $\text{Pb}^{2+}$ , and $\text{Hg}_2^{2+}$ .	Soluble: $\text{NaCl}$ , $\text{KBr}$ , $\text{CaI}_2$ Insoluble: $\text{AgCl}$ , $\text{PbI}_2$ , $\text{Hg}_2\text{Cl}_2$
4	All sulfates ( $\text{SO}_4^{2-}$ ) are soluble, except with $\text{Ba}^{2+}$ , $\text{Pb}^{2+}$ , $\text{Sr}^{2+}$ , and $\text{Ca}^{2+}$ (slightly soluble).	Soluble: $\text{Na}_2\text{SO}_4$ , $\text{CuSO}_4$ Insoluble: $\text{BaSO}_4$ , $\text{PbSO}_4$
5	All carbonates ( $\text{CO}_3^{2-}$ ), phosphates ( $\text{PO}_4^{3-}$ ), sulfides ( $\text{S}^{2-}$ ), and hydroxides ( $\text{OH}^-$ ) are insoluble, except with Group 1 metals ( $\text{Li}^+$ , $\text{Na}^+$ , $\text{K}^+$ , etc.) and $\text{NH}_4^+$ .	Insoluble: $\text{CaCO}_3$ , $\text{FePO}_4$ , $\text{CuS}$ , $\text{Fe(OH)}_3$ Soluble: $\text{Na}_2\text{CO}_3$ , $\text{K}_3\text{PO}_4$ , $\text{NaOH}$ , $\text{NH}_4\text{OH}$
6	All compounds of Group 1 metals ( $\text{Li}^+$ , $\text{Na}^+$ , $\text{K}^+$ , $\text{Rb}^+$ , $\text{Cs}^+$ ) and ammonium ( $\text{NH}_4^+$ ) are soluble.	$\text{NaCl}$ , $\text{KOH}$ , $\text{NH}_4\text{NO}_3$

Watch out

A common mistake: students remember "all chlorides are soluble" and forget the exceptions. AgCl is insoluble — that's why it's used in photography and qualitative analysis. Always check the exceptions first.

How to Use the Rules (Step-by-Step)

Suppose you need to predict whether $\text{PbSO}_4$ dissolves in water.

Identify the ions: $\text{Pb}^{2+}$ and $\text{SO}_4^{2-}$ .
Check the rules in order:
- Rule 1 (nitrates)? No. …

Why this formula?

Solubility Rules: Why They Work (Not Just What They Say)

Solubility rules are not arbitrary — they emerge from thermodynamics and electrostatic interactions between ions in water. Let's break down the why behind the key patterns.

1. The Core Idea: "Like Dissolves Like" at the Ionic Level

Water dissolves ionic compounds because it is polar. The $\delta^+$ hydrogen ends attract anions, and the $\delta^-$ oxygen end attracts cations.

Driving force: The lattice energy (energy holding the solid together) vs. the hydration energy (energy released when ions are surrounded by water).
Net energy change:

\Delta H_{\text{solution}} = \text{Lattice Energy} - \text{Hydration Energy}

If $\Delta H_{\text{solution}}$ is negative (exothermic) or small positive, the compound tends to dissolve.

2. Why Some Salts Are Always Soluble (Group 1 & NH₄⁺)

Rule: All salts of $\ce{Na+}$ , $\ce{K+}$ , $\ce{NH4+}$ are soluble.

Why?

These cations are large and have low charge density (charge/size ratio is small).
Their lattice energies are relatively low because the ions are far apart in the crystal.
Their hydration energies are high enough to overcome the lattice energy.

Key formula: For a cation like $\ce{K+}$ , the hydration energy is roughly:

$\Delta H_{\text{hyd}} \propto \frac{q^2}{r}$

where $q$ is charge and $r$ is ionic radius.

Since $q=1$ and $r$ is large, $\Delta H_{\text{hyd}}$ is moderate — but lattice energy is even smaller.

Result: The net $\Delta H_{\text{solution}}$ is negative → spontaneous dissolution.

3. Why Nitrates, Acetates, and Chlorates Are Always Soluble

Rule: All nitrates ( $\ce{NO3-}$ ), acetates ( $\ce{CH3COO-}$ ), and chlorates ( $\ce{ClO3-}$ ) are soluble.

Why?

These anions are large and polyatomic — their charge is spread over many atoms.
This delocalization of charge means they have low charge density.
They form weak ionic bonds with cations (low lattice energy).
Water can easily hydrate them because the negative charge is not concentrated.

Key insight: The lattice energy for $\ce{NaNO3}$ is much smaller than for $\ce{NaCl}$ because $\ce{NO3-}$ is larger and more polarizable.

4. The "Exceptions" — Why Some Salts Are Insoluble

4.1. Carbonates, Phosphates, Sulfides (Except with Group 1 & NH₄⁺)

Rule: Most carbonates ( $\ce{CO3^{2-}}$ ), phosphates ( $\ce{PO4^{3-}}$ ), and sulfides ( $\ce{S^{2-}}$ ) are insoluble.

Why?

These anions have high charge ( $-2$ or $-3$ ) and are small (especially $\ce{S^{2-}}$ ).
This gives them very high charge density.
They form extremely strong ionic bonds with cations (very high lattice energy).
The hydration energy, though large, is not enough to overcome the lattice energy.

Example: For $\ce{CaCO3}$ :

$\text{Lattice energy} \approx -2800 \text{ kJ/mol}$

$\text{Hydration energy} \approx -2400 \text{ kJ/mol}$

Net $\Delta H_{\text{solution}} \approx +400 \text{ kJ/mol}$ → insoluble

4.2. Silver, Lead, Mercury Halides

Rule: $\ce{AgCl}$ , $\ce{PbCl2}$ , $\ce{Hg2Cl2}$ are insoluble (most other chlorides are soluble).

Why?

$\ce{Ag+}$ , $\ce{Pb^{2+}}$ , $\ce{Hg2^{2+}}$ are soft (polarizable) cations.
They form covalent character in their bonds with halides (especially $\ce{Cl-}$ , $\ce{Br-}$ , $\ce{I-}$ ).
This covalent contribution increases the effective lattice energy beyond what simple ionic models predict.
Water cannot break these partially covalent bonds.

Key formula: The polarizing power of a cation is:

$\phi = \frac{q}{r}$ …

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The solubility of organic compounds in water is primarily determined by their polarity and ability to form hydrogen bonds; compounds with strong polar groups and small non-polar parts are highly soluble, those whose one polar group fights a sizeable non-polar part are partially soluble, and those that are largely non-polar (or cannot donate a hydrogen bond) are insoluble. Formic acid and ethylene glycol are highly soluble, phenol and pentanol are partially soluble, and toluene and chloroform are insoluble.

The solubility of a substance in water is governed by the fundamental principle of "like dissolves like." Water is a highly polar solvent, capable of forming extensive hydrogen bonds with itself. For an organic compound to dissolve in water, it must be able to interact favorably with water molecules, either by forming hydrogen bonds or by having significant polarity that allows for strong dipole-dipole interactions.

Here's a breakdown of the factors influencing solubility in water:

Polarity: Water molecules are highly polar due to the electronegativity difference between oxygen and hydrogen, leading to partial positive charges on hydrogen atoms and a partial negative charge on the oxygen atom. Organic compounds with polar functional groups (like hydroxyl $(-\text{OH})$ , carboxyl $(-\text{COOH})$ , amino $(-\text{NH}_2)$ ) can interact with water through dipole-dipole forces.
Hydrogen Bonding: This is the most crucial factor. Functional groups that can act as hydrogen bond donors (e.g., $-\text{OH}$ , $-\text{NH}_2$ , $-\text{COOH}$ ) form strong hydrogen bonds with water molecules. These interactions help overcome the strong hydrogen bonds between water molecules themselves and allow the organic compound to integrate into the water structure. A group that can only accept a hydrogen bond (like the Cl in chloroform, or a carbonyl oxygen) gives a much weaker effect than a genuine $-\text{OH}$ / $-\text{NH}_2$ donor.
Size of the Non-Polar Hydrocarbon Part: Even if an organic compound has polar functional groups, a large non-polar hydrocarbon chain or ring will diminish its solubility. The non-polar part cannot form favorable interactions with water and disrupts water's hydrogen-bond network, making dissolution energetically unfavorable.

We classify solubility broadly:

Highly soluble: Dissolves readily, often miscible or dissolves to a high concentration (e.g., greater than 10 g/100 mL).
Partially soluble: Dissolves to a limited but still clearly measurable extent (roughly 1-10 g/100 mL).
Insoluble: Dissolves to a negligible extent (well under about 1 g/100 mL), effectively forming two layers or barely mixing.

Let's analyze each compound:

Phenol (C6H5OH) -- The $-\text{OH}$ group is polar and donates hydrogen bonds to water, but the benzene ring is a fairly large non-polar part. Conclusion: partially soluble (about 8 g/100 mL at 20 degrees C) -- enough $-\text{OH}$ character to dissolve appreciably, but the ring keeps it well short of "highly soluble."
Toluene (C6H5CH3) -- A hydrocarbon with no polar functional group capable of hydrogen bonding. Conclusion: insoluble (about 0.05 g/100 mL).
Formic acid (HCOOH) -- The $-\text{COOH}$ group is highly polar and hydrogen-bonds strongly (both as donor and acceptor); the non-polar part is negligible (just one H). Conclusion: highly soluble -- miscible with water in all proportions.
Ethylene glycol (HOCH2CH2OH) -- Two $-\text{OH}$ groups on a two-carbon chain give extensive hydrogen bonding with a tiny non-polar part. Conclusion: highly soluble -- miscible with water in all proportions. …

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Method: Polarity & Hydrogen Bonding Analysis

This method uses the principle "like dissolves like" -- polar solvents (water) dissolve polar solutes best. The key factors are:

Hydrogen bonding ability (most important for water solubility) -- specifically, can the solute donate an O-H/N-H bond to water?
Polarity of the functional group
Hydrocarbon chain length (the non-polar part that works against solubility)

Steps

Step 1: Identify the functional group in each compound and classify it as polar/non-polar, and note whether it can donate a hydrogen bond.

Compound	Functional Group	H-bond donor to water?
Phenol	-OH (aromatic)	Yes
Toluene	-CH3 on benzene	No
Formic acid	-COOH	Yes (strong)
Ethylene glycol	Two -OH groups	Yes (strong)
Chloroform	C-Cl (3 Cl atoms)	No (acceptor only)
Pentanol	-OH (5-carbon chain)	Yes, but chain limits it

Step 2: Check hydrogen-bonding ability with water.

Highly soluble: Strong H-bond donors with a short carbon chain (3-4 carbons or fewer) -- e.g. formic acid, ethylene glycol.
Partially soluble: Can H-bond, but the non-polar part is large enough to noticeably limit solubility -- e.g. phenol (benzene ring vs one -OH) and pentanol (5-carbon chain vs one -OH).
Insoluble: Cannot donate a hydrogen bond at all -- toluene (no polar group) and chloroform (polar, but no H on O/N/F).

Step 3: Apply the rule -- For alcohols, solubility falls off sharply once the chain reaches about 4-5 carbons; a single -OH group can no longer fully out-compete the hydrocarbon tail. Pentanol therefore keeps only a limited solubility (about 2 g/100 mL) -- partially soluble, not insoluble; genuinely insoluble compounds are the ones with no H-bond donor at all.

Final Classification

| Compound | Solubility | Reason |

|----------|------------|--------| …

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Here is a breakdown of the common mistakes students make when applying solubility rules to organic compounds, specifically for this question, and how to avoid each.

Mistake 1: Confusing "Polar" with "Soluble"

The Error: Students see a polar bond (like C-Cl in chloroform or O-H in phenol) and immediately label the compound as "highly soluble." They forget that the size of the non-polar (hydrocarbon) part is the deciding factor.

Example from the list: Chloroform ( $CHCl_3$ ) is polar, but it is insoluble in water. Phenol ( $C_6H_5OH$ ) has an -OH group, but the large benzene ring makes it only partially soluble.

How to Avoid:

Apply the "Like Dissolves Like" rule strictly: Water is a small, highly polar molecule. For a compound to be highly soluble, it must be small and have a high density of polar groups (like -OH, -COOH).
Use the "Carbon Count" rule of thumb: For a molecule to be highly soluble in water, the number of carbon atoms should generally be ≤ 3 for every strong polar group (like -OH). If the carbon chain is longer, solubility drops sharply.
- Example: Formic acid (1 C) → Highly soluble. Pentanol (5 C) → only partially soluble (about 2 g/100 mL).

Mistake 2: Ignoring the "Hydrocarbon Bully" Effect

The Error: Students focus only on the functional group (e.g., "phenol has an -OH, so it's like alcohol, therefore soluble"). They ignore that the benzene ring (in phenol) or long alkyl chain (in pentanol) is hydrophobic and pulls solubility down.

Example from the list: Phenol is often wrongly marked as "highly soluble." It is actually partially soluble (about 8 g/100 mL). The benzene ring is a large, non-polar "bully" that resists water.

How to Avoid:

Visualize the molecule: Draw the structure. If the non-polar part (hydrocarbon ring or chain) looks significantly larger than the polar part, expect reduced solubility -- partial at best.
Memorize the exceptions: Phenol is a classic "partially soluble" example. It is more soluble than toluene but far less soluble than ethanol.

Mistake 3: Forgetting that "Insoluble" ≠ "Non-polar"

The Error: Students assume that only non-polar compounds can be insoluble, and conversely that any polar compound must dissolve.

Why it's wrong: Chloroform is polar, yet it is insoluble -- polarity without an H-bond donor is not enough. And a polar -OH compound like pentanol is dragged down to merely partial solubility by its long chain.

How to Avoid:

Classify by structure, not just polarity: Ask: "Can it donate a hydrogen bond, and is the molecule small enough to fit into water's hydrogen-bonding network?"
- Highly soluble: Small, strong polar groups (e.g., formic acid, ethylene glycol).
- Partially soluble: One good polar group fighting a sizeable non-polar part (e.g., phenol, pentanol).
- Insoluble: No H-bond donor at all, whether non-polar (toluene) or polar (chloroform).

Mistake 4: Misjudging "Partially Soluble" vs. "Insoluble"

The Error: Students treat "partially soluble" as a vague middle ground and often guess wrong -- most commonly by calling pentanol "insoluble" because its chain is long. …

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Q.Amongst the following compounds, identify which are insoluble, partially soluble and highly soluble in water?

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