10th Class Chemistry Chapter 5 Notes PDF 2026 – Salts

Definition

A neutralization reaction is one in which an acid reacts with a base to produce a salt and water — e.g. $H C l + N a O H \to N a C l + H_{2} O$ .

When marble (calcium carbonate) reacts with dilute nitric acid, the salt formed is calcium nitrate, $C a (N O_{3})_{2}$ (along with water and carbon dioxide, since a carbonate reacting with an acid also releases $C O_{2}$ ).

Definition

General solubility patterns: all common sodium, potassium, and ammonium salts are soluble; all nitrates are soluble.

Most common carbonates are insoluble (except sodium, potassium, ammonium carbonates); most common sulphates are soluble, except a few key insoluble ones (like barium sulphate and lead(II) sulphate); most common chlorides are soluble, except a few key insoluble ones (like silver chloride and lead(II) chloride).

Lead(II) chloride is only slightly soluble in cold water, but its solubility increases noticeably in hot water — so it can appear to dissolve on heating and then recrystallize as the solution cools, which is a useful and recognizable property in the lab.

Insoluble lead salts include lead(II) chloride, lead(II) sulphate, and lead(II) carbonate (among others) — this is why lead salts are commonly prepared by precipitation, since mixing two soluble lead-free and lead-containing solutions immediately forms the insoluble lead salt as a solid.

Most metal hydroxides (bases) are insoluble in water, except Group I hydroxides (like NaOH, KOH) which are freely soluble; calcium hydroxide is only partially (slightly) soluble.

Definition

A crystal lattice is the regular, repeating three-dimensional arrangement of ions (or atoms/molecules) in a solid crystal.

Sodium chloride crystallizes with each $N a^{+}$ ion surrounded by $6$ $C l^{-}$ ions, and each $C l^{-}$ ion surrounded by $6$ $N a^{+}$ ions, in a repeating cubic pattern — this specific ratio and arrangement gives NaCl crystals their characteristic cubic shape.

The strength of the electrostatic (ionic) forces holding a crystal lattice together depends on the size of the ionic charges (higher charges attract more strongly) and the distance between ion centres (smaller ions attract more strongly, since they can get closer together).

Example 1 — comparing melting points

Compare the melting points of $N a C l$ and $M g C l_{2}$ .

$M g^{2 +}$ has a $+ 2$ charge compared to $N a^{+}$ 's $+ 1$ charge, and $M g^{2 +}$ is also a smaller ion than $N a^{+}$ .
Both a higher charge and a smaller ion size mean stronger electrostatic attraction in the lattice, requiring more energy to break apart.

Answer: $M g C l_{2}$ has a higher melting point than $N a C l$

Rule

Ionic salts generally have very high melting points because breaking apart their crystal lattice means overcoming strong electrostatic attractions between oppositely charged ions extending throughout the whole structure, not just a single bond — a large amount of energy is needed to separate the tightly-held, oppositely-charged ions from their fixed lattice positions.

Ionic salts exist as solids at room temperature (rather than liquids or gases) for the same reason: the strong electrostatic attractions between ions throughout the lattice hold them firmly in fixed positions, and only very high temperatures provide enough energy to overcome this and let the ions move freely (melt).

Definition

Ionic compounds conduct electricity only when molten or dissolved in water (aqueous solution) — not when solid.

In the solid state, ions are locked into fixed positions in the rigid crystal lattice and cannot move to carry charge, so no current flows. When molten or dissolved, the lattice breaks down and ions become free to move throughout the liquid, allowing them to carry electric current towards the oppositely charged electrode.

Definition

Soluble salts are usually prepared by reacting an acid with an excess of an insoluble base or an insoluble metal — the leftover unreacted solid is then simply filtered off, leaving a solution of only the pure salt, which is then evaporated and crystallized to obtain solid crystals.

The insoluble base or metal is always added in excess when preparing a soluble salt this way, to ensure all of the acid reacts completely — this guarantees the final filtered solution contains only salt and water, with no leftover acid to contaminate the crystals; any excess solid is easily removed by filtration.

Soluble salts made from a soluble acid and a soluble base (e.g. $K_{2} S O_{4}$ from $K O H$ and $H_{2} S O_{4}$ ) instead need titration, since there's no excess solid to filter off both reactants are already dissolved.

In this titration method, phenolphthalein indicator is used to precisely detect the exact neutralization point — it's colourless in acid and pink in alkali, so a sharp, clearly visible colour change signals exactly when just enough acid has been added to fully neutralize the base, letting the correct volumes be measured for a pure, excess-free salt solution.

Definition

Insoluble salts are usually prepared by precipitation: mixing two suitable soluble solutions whose ions combine to form the desired insoluble salt, which immediately precipitates out of solution as a solid and can be filtered, washed, and dried.

Calcium sulphate (insoluble) can be prepared by mixing solutions of a soluble calcium salt (like calcium chloride) with a soluble sulphate (like sodium sulphate/dilute sulphuric acid), precipitating solid calcium sulphate directly.

Pure copper sulphate crystals (soluble) are prepared by reacting dilute sulphuric acid with excess copper(II) oxide or copper(II) carbonate (an insoluble base), filtering off the unreacted excess solid, then evaporating the filtrate to concentrate it and allowing it to cool and crystallize slowly into pure blue crystals.

Example 2 — preparation methods for given salts

How are the following salts prepared: $K_{2} S O_{4}$ , $C u C l_{2}$ , $C a C l_{2}$ ?

$K_{2} S O_{4}$ : both $K O H$ and $H_{2} S O_{4}$ are soluble, so titration is used (with phenolphthalein), followed by evaporation and crystallization.
$C u C l_{2}$ : $C u O$ or $C u C O_{3}$ (insoluble base) is added in excess to dilute hydrochloric acid, filtered to remove excess solid, then the filtrate is evaporated and crystallized.
$C a C l_{2}$ : excess calcium carbonate (insoluble) is added to dilute hydrochloric acid, filtered to remove excess solid, then the filtrate is evaporated and crystallized.

Answer: $K_{2} S O_{4}$ by titration; $C u C l_{2}$ and $C a C l_{2}$ by reacting the acid with an excess insoluble base/carbonate, then filtering and crystallizing

Quick Recap

●Neutralization: acid $+$ base $\to$ salt $+$ water
●Solubility: Na/K/NH $_{4}$ /nitrate salts always soluble; most carbonates insoluble; most sulphates/chlorides soluble except key exceptions (e.g. PbCl $_{2}$ , PbSO $_{4}$ , AgCl); most hydroxides insoluble except Group I
●Ionic lattice: strong electrostatic forces $\to$ high melting point, solid at room temperature; strength $↑$ with charge, $↓$ with ion size
●Ionic compounds conduct only when molten/dissolved (free-moving ions), not as solids (fixed ions)
●Soluble salt from insoluble base/acid: excess solid $+$ filter $+$ evaporate/crystallize. Soluble + soluble: titration (phenolphthalein)
●Insoluble salt: precipitation — mix two soluble solutions, filter the precipitate

10th Class Chemistry Chapter 5 Notes

Neutralization and Salt Formation

Solubility of Salts

Ionic Structure and Crystal Lattices

Electrical Conductivity of Ionic Compounds

Preparing Salts in the Laboratory

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