Sodium Chloride and Tests for Halide Ions – Chlorine and Its Inorganic Compounds (Senior 3 Chemistry) | Ganzaa

Chlorine and Its Inorganic Compounds

Lessons

Senior 3 · Chemistry

Chlorine and Its Inorganic Compounds

4 lessons

Lesson 4 of 4

Sodium Chloride and Tests for Halide Ions

25 min

White crystals from a crater lake

At Lake Katwe, a salt lake in a volcanic crater in western Uganda, workers let the salty water evaporate in shallow pans under the sun and rake up the crystals that form. The main compound in those crystals, and in the salt on every kitchen table, is sodium chloride. It seasons and preserves food, it is a raw material for chlorine and sodium hydroxide, and a small amount of it in a drink can save the life of a child with diarrhoea. Too much of it, on the other hand, damages health. This lesson looks at sodium chloride, and at how chemists recognise chloride and the other halide ions.

Obtaining and preparing sodium chloride

From nature. Sodium chloride is obtained by evaporating sea water or the brine of salt lakes in pans, as at Lake Katwe, and by mining underground beds of rock salt left when ancient seas dried up.

In the laboratory. It is a soluble salt, so it is prepared by the titration method you used in S2: dilute hydrochloric acid is added to sodium hydroxide solution until the mixture is exactly neutral, then the solution is evaporated until crystals form:

$\mathrm{NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)}$

It also forms when sodium carbonate reacts with hydrochloric acid, and when sodium burns in chlorine (Lesson 2).

Properties of sodium chloride

A white, crystalline solid made of sodium ions and chloride ions in a giant ionic lattice, so it has a high melting point (801 °C).
Soluble in water (about 36 g in 100 g of water at room temperature); the solution is neutral, pH 7.
Conducts electricity when molten or dissolved, because its ions can then move; it is decomposed by electrolysis (a later unit).
Gives hydrogen chloride when warmed with concentrated sulfuric acid (Lesson 3), and a white precipitate with silver nitrate solution.

Definition

Oral rehydration solution (ORS): a drink of clean water containing measured amounts of sodium chloride, glucose and other salts, which replaces the water and salts lost in diarrhoea. The glucose helps the gut absorb the sodium, and water follows the sodium into the body. Example: the WHO and UNICEF formula contains 2.6 g of sodium chloride and 13.5 g of glucose in every litre.

Uses of sodium chloride

Food: seasoning, and preserving fish and meat, because concentrated salt draws water out of bacteria and stops them growing.
Medicine: oral rehydration solution, and saline drips in hospitals.
Industry: electrolysis of brine makes chlorine, hydrogen and sodium hydroxide; salt is also used to make sodium carbonate and to separate soap during soap-making.

Health hazard: too much salt

The body needs some sodium, but a diet high in salt raises blood pressure, which increases the risk of heart disease and stroke. The World Health Organization recommends that adults eat less than 5 g of salt a day, just under one teaspoon, including the salt already in bread, stock cubes, snacks and processed foods.

Tests for chloride and the other halide ions

Acidify the solution with dilute nitric acid, then add silver nitrate solution. Each halide gives a precipitate of a different colour; ammonia solution confirms which (Figure 4.1):

$\mathrm{Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)} \quad \text{white}$

$\mathrm{Ag^+(aq) + Br^-(aq) \rightarrow AgBr(s)} \quad \text{cream}$

$\mathrm{Ag^+(aq) + I^-(aq) \rightarrow AgI(s)} \quad \text{yellow}$

Why nitric acid? It removes carbonate and sulfite ions, which would also give precipitates with silver ions. Hydrochloric acid must not be used, because it adds chloride ions, and sulfuric acid can give a precipitate of silver sulfate.
The silver halide precipitates slowly turn grey or purple in light, because light decomposes them to silver. This is the chemistry of old photographic film.

Figure 4.1: Identifying chloride, bromide and iodide ions with silver nitrate and ammonia.

Warning

Safety. Silver nitrate stains skin and clothes black and is harmful to the eyes; concentrated ammonia solution gives off an irritating gas. Wear eye protection and gloves, and use concentrated ammonia in a fume cupboard.

Summary: the tests of this year's non-metal units

Table 4.1: Tests for gases

Gas	Test	Result
oxygen, $\mathrm{O_2}$	glowing splint	relights
hydrogen, $\mathrm{H_2}$	lighted splint	squeaky pop
carbon dioxide, $\mathrm{CO_2}$	bubble through limewater	limewater turns milky
carbon monoxide, $\mathrm{CO}$	light it; test the product with limewater	burns with a blue flame; product turns limewater milky
ammonia, $\mathrm{NH_3}$	damp red litmus; concentrated hydrochloric acid on a glass rod	litmus turns blue; dense white fumes
nitrogen dioxide, $\mathrm{NO_2}$	look; damp blue litmus	reddish-brown gas; litmus turns red
sulfur dioxide, $\mathrm{SO_2}$	acidified potassium dichromate(VI)	orange turns green
hydrogen sulfide, $\mathrm{H_2S}$	moist lead(II) ethanoate paper	turns black
chlorine, $\mathrm{Cl_2}$	look; damp blue litmus	yellow-green; litmus turns red, then white
hydrogen chloride, HCl	concentrated ammonia on a glass rod	dense white fumes

Table 4.2: Tests for anions

Anion	Test	Result
carbonate, $\mathrm{CO_3^{2-}}$	add dilute acid	fizzes; the gas turns limewater milky
sulfate, $\mathrm{SO_4^{2-}}$	dilute hydrochloric acid, then barium chloride	white precipitate that stays
sulfite, $\mathrm{SO_3^{2-}}$	warm with dilute hydrochloric acid	gas turns acidified dichromate(VI) green
nitrate, $\mathrm{NO_3^-}$	warm with sodium hydroxide and aluminium	ammonia given off (damp red litmus turns blue)
chloride, $\mathrm{Cl^-}$	dilute nitric acid, then silver nitrate	white precipitate, soluble in dilute ammonia
bromide, $\mathrm{Br^-}$	dilute nitric acid, then silver nitrate	cream precipitate, soluble in concentrated ammonia
iodide, $\mathrm{I^-}$	dilute nitric acid, then silver nitrate	yellow precipitate, insoluble in ammonia

Worked example

Problem. Excess silver nitrate solution is added to a solution containing 1.17 g of sodium chloride. What mass of silver chloride precipitates? (Na = 23, Cl = 35.5, Ag = 108)

Solution. $M(\mathrm{NaCl}) = 58.5\,\mathrm{g\,mol^{-1}}$ , so $n = \frac{1.17}{58.5} = 0.0200\,\mathrm{mol}$ .

$\mathrm{NaCl(aq) + AgNO_3(aq) \rightarrow AgCl(s) + NaNO_3(aq)}$ : 1 mol NaCl gives 1 mol AgCl. $M(\mathrm{AgCl}) = 108 + 35.5 = 143.5\,\mathrm{g\,mol^{-1}}$ , so the mass is $0.0200 \times 143.5 = 2.87\,\mathrm{g}$ .

Worked example

Problem. A white solid gives a bright yellow-orange flame test. Its solution, acidified with dilute nitric acid, gives a cream precipitate with silver nitrate; the precipitate does not dissolve in dilute ammonia but dissolves in concentrated ammonia. Identify the solid, and explain why the student acidified with nitric acid, not hydrochloric acid.

Solution. The yellow-orange flame shows sodium ions (S2). A cream precipitate soluble only in concentrated ammonia is silver bromide, so the anion is bromide. The solid is sodium bromide, NaBr. Hydrochloric acid would have added chloride ions, which give their own white precipitate of silver chloride and would confuse the result.

Tip

White, cream, yellow: Cl, Br, I. The precipitates get darker, and harder to dissolve in ammonia, as you go down the group. Always acidify with nitric acid first.

Halides in solution: test using acidified silver nitrate (4 min). Watch the colour of each precipitate as silver nitrate is added, and match it to Figure 4.1.

✦

Practice

Describe how sodium chloride crystals are prepared in the laboratory from hydrochloric acid and sodium hydroxide.
Explain why sodium chloride has a high melting point and conducts electricity only when molten or dissolved.
Give three uses of sodium chloride.
Why does the WHO advise adults to eat less than 5 g of salt a day?
Describe the test for chloride ions and give the ionic equation.
A solution gives a yellow precipitate with acidified silver nitrate that does not dissolve in concentrated ammonia. Which ion is present?
Why must a solution not be acidified with hydrochloric acid before testing it for chloride ions?
What mass of silver bromide forms when excess silver nitrate is added to a solution of 2.06 g of sodium bromide? (Na = 23, Br = 80, Ag = 108)

Show answersHide answers

Titrate the acid into the alkali with an indicator to find the exact volume needed; repeat with the same volumes and no indicator; evaporate the solution until crystals begin to form, cool, filter and dry the crystals.
It is a giant ionic lattice with strong attractions between oppositely charged ions, which need a lot of energy to overcome. In the solid the ions cannot move; when molten or dissolved they are free to move and carry charge.
Any three: seasoning and preserving food; oral rehydration solution; saline drips; making chlorine and sodium hydroxide by electrolysis; making sodium carbonate; soap-making.
Too much salt raises blood pressure, which increases the risk of heart disease and stroke.
Add dilute nitric acid, then silver nitrate solution: a white precipitate (soluble in dilute ammonia) shows chloride. $\mathrm{Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)}$ .
Iodide, $\mathrm{I^-}$ .
Hydrochloric acid contains chloride ions, which would give a white precipitate whether or not the solution contained chloride.
$M(\mathrm{NaBr}) = 103$ ; $n = \frac{2.06}{103} = 0.0200\,\mathrm{mol}$ ; $M(\mathrm{AgBr}) = 188$ ; mass $= 0.0200 \times 188 = 3.76\,\mathrm{g}$ .

1White crystals from a crater lake

2Obtaining and preparing sodium chloride

3Properties of sodium chloride

4Uses of sodium chloride

5Health hazard: too much salt

6Tests for chloride and the other halide ions

7Summary: the tests of this year's non-metal units