Topic 12A Chemical Reactions
Updated 2008-08-19 R
Please send comments to: J.Elfick@uq.edu.au
See also: Interesting websites

---

Table of contents
12.10.0 Salts, sodium chloride "table salt"
12.10.7.0 pH value of buffer solutions
12.11.0 Tests for an unknown substance, qualitative analysis
12.11.4.0 Group tests to identify cations, prepare a solution for group analysis
12.11.5.0 Tests for anions in unknown solution, tests for acid radicals in solution
12.11.6.0 Tests for metallic radicals
12.12.0 Tests for gases
12.14.0 Activity series of metals as reducing agents
12.15.0 Reactions of metals with water
12.16.0 Carbonates
12.17.0 Oxides, acidic and basic oxides

---

12.10.0 Salts, sodium chloride "table salt"
12.10.1 Crystals of different salts
12.10.2 Prepare salts by different methods
12.10.2.1 Reactions of dilute acids with metals
12.10.2.2 Test the reactions of dilute acids with alkalis
12.10.2.3 Reactions of acids with metallic oxides
12.10.2.4 Reactions of acids with metal carbonates
12.10.2.5 Direct union of the elements, sodium and chlorine
12.10.2.6 Add solutions of a salt to another salt
12.10.2.6.1 Artificial gemstones, potassium sulfate, aluminium sulfate
12.10.3 Hydrolysis of sodium carbonate
12.10.3.1 Hydrolysis of sodium hydrogen carbonate
12.10.3.2 Hydrolysis of ammonium chloride
12.10.3.3 Reactions of salts with water
12.10.4 pH of different salts dissolved in water
12.10.5 Hydrolysis of iron (III) chloride
12.10.6 Prepare an acid salt, sodium hydrogen sulfate

---

12.10.7.0 pH value of buffer solutions
12.10.7 Buffer solutions
12.10.7.1 Dilution effect on buffer solutions
12.10.7.2 Natural buffers
12.10.7.3 Prepare a buffered solution
12.10.7.4 Salt effect on buffer solutions
12.10.8 Prepare solutions of pH values 3 to 11 with buffer solutions
12.10.9 Show the effect of a buffer salt
12.10.10 Change in pH near the equivalence point
12.10.11 pH values of solutions of salts
12.10.12 Fireproofing

---

12.11.0 Tests for an unknown substance, qualitative analysis
12.11.3.1 Borax bead tests to identify metals in metallic salts and minerals
12.11.3.2 Flame tests to identify metals and their compounds
12.11.3.3 Solubility tests, prepare a solution for group analysis
12.11.3.4 Test substances by action of heat in dry test-tube
12.11.3.5 Test substances by action with dilute hydrochloric acid and note gas produced
12.11.3.6 Test substances by action with hot concentrated sulfuric acid
12.11.3.7 Test substances by action of caustic soda solution, sodium hydroxide
12.11.3.8 Test substances by action of heat on charcoal with fusion mixture
12.11.3.9 Test for aluminium compounds in solution
12.11.3.10 Confirmatory tests with original solution or solid

---

12.11.4.0 Group tests to identify cations, prepare a solution for group analysis
12.11.4.1 Group 1 tests for Ag⁺, Pb²⁺
12.11.4.2 Group 2 tests for Bi³⁺, Cd²⁺, Cu²⁺, Sn²⁺
12.11.4.3 Group 3 tests for Al³⁺, Cr³⁺, Fe²⁺, Fe³⁺
12.11.4.4 Group 4 tests for Co²⁺, Mn²⁺, Ni²⁺, Zn²⁺
12.11.4.5 Group 5 tests for Ba²⁺, Ca²⁺, Sr²⁺
12.11.4.6 Group 6 tests for K⁺, Mg²⁺, Na⁺, NH₄⁺

---

12.11.5.0 Tests for anions in unknown solution, tests for acid radicals in solution
12.11.5.1 Tests for acetates, CH₃COO^- and (CO₃)^2-
12.11.5.2a Tests for antimonates, borates, oxalates
12.11.5.3 Tests for arsenates
12.11.5.4 Tests for bicarbonates
12.11.5.5 Tests for borates
12.11.5.6 Tests for bromides
12.11.5.7 Tests for carbonates
12.11.5.8 Tests for chlorides
12.11.5.9 Tests for chromates
12.11.5.10 Tests for halides, Cl^-, Br^-, I^-
12.11.5.11 Tests for hydroxides
12.11.5.12 Tests for iodides
12.11.5.13 Tests for nitrates
12.11.5.14 Tests for oxalates
12.11.5.15 Tests for phosphates
12.11.5.16 Tests for sulfates
12.11.5.17 Tests for sulfides
12.11.5.18 Tests for sulfites

---

12.11.6.0 Tests for metallic radicals
12.11.6.1 Chemistry of group separations.
12.11.6.2 Preliminary experiments before the separation of Group I metals, silver and lead.
12.11.6.3 Separation into groups.
12.11.7.1 Group I Insoluble chlorides, PbCl₂, AgCl [Hg₂Cl₂ omitted]
12.11.7.2 Group II Sulfides insoluble in dilute hydrochloric acid
12.11.7.2a Group IIa PbS, Bi₂S₃, CuS, CdS [HgS omitted]
12.11.7.2b Group IIb As₂S₃, Sb₂S₃, SnS, SnS₂
12.11.7.3 Group III Insoluble hydroxides, Fe(OH)₃, Cr(OH)₃, Al(OH)₃
12.11.7.4 Group IV Insoluble sulfides precipitated by hydrogen sulfide, ZnS, MnS, CoS, NiS
12.11.7.5 Group V Insoluble carbonates, CaCO₃, BaCO₃, SrCO₃
12.11.7.6 Group VI Magnesium, sodium and potassium, Mg, Na, K

---

12.12.0 Tests for gases
3.31.3 Test for water with cobalt (II) chloride
3.33.1 Tests for ammonia, ammonia fountain experiment
3.34.1 Tests for carbon dioxide
3.34.2 Test the breath for carbon dioxide
3.40.1 Tests for chlorine
3.41.1 Tests for hydrogen
3.42.1 Tests for hydrogen chloride
3.43.1 Tests for hydrogen sulfide solution
3.45.1 Test for dinitrogen oxide (nitrous oxide)
3.49.1 Tests for oxygen
3.51.1 Test for sulfur dioxide

---

12.14.0 Activity series of metals as reducing agents
12.14.1 Zinc displaces lead from lead nitrate solution
12.14.2 Metals in copper (II) sulfate solution
12.14.2.1 Magnesium, or zinc, in copper (II) sulfate solution
12.14.2.2 Iron in copper (II) sulfate solution
12.14.2.3 Iron and zinc with copper (II) sulfate solution
12.14.2.4 Zinc in lead nitrate solution and iron in copper (II) sulfate solution
12.14.2.5 Zinc and copper in sulfuric acid
12.14.2.6 Activity of metals and tendency to form ions
3.72 Displacement of copper from solution of copper ions
3.74 Displacement of hydrogen from acids by metals

---

12.15.0 Reactions of metals with water
3.73 Reactions of sodium with water
3.75 Reactions of salts with water
12.15.1 Reactions of metals with water, Cu, Zn, Fe, Mg, Al
12.15.2 Reactions of sodium with water
12.15.3 Reactions of metals with steam
12.15.4 Reactions of metals with water
12.15.5 Reactions of non-metals with water

---

12.16.0 Carbonates
12.16.1 Pass carbon dioxide through calcium carbonate suspension
12.16.1.1 Test for carbon dioxide, pass carbon dioxide through calcium hydroxide solution, limewater
12.16.2 Prepare sodium hydrogen carbonate with sodium carbonate
12.16.3 Heat different carbonates, carbonates of Cu, Mg, Na, Pb and Zn
12.16.3.1 Heat ammonium carbonate (smelling salts)
12.16.3.2 Reactions of ammonium carbonate with alkalis
12.16.3.3 Reactions of ammonium carbonate with acids
12.16.3.4 Ammonium carbonate solution precipitates metal carbonates
12.16.4 Heat a hydrogen carbonate, sodium hydrogen carbonate (sodium bicarbonate)
12.16.5 Put egg in bottle and tie knot in a bone!
12.16.6 Reactions of dilute acids with sodium hydrogen carbonate
12.16.6.01 Prepare an imitation volcano with baking soda
12.16.7 Solvay process, sodium carbonate
3.31.2 Expose sodium carbonate decahydrate, washing soda, to the air

---

12.17.0 Oxides, acidic and basic oxides
12.17.1 Examine different oxides
12.17.1.1 Different oxides and the periodic table
12.17.2 Basic oxides, copper (II) oxide (copper oxide)
12.17.2.1 Copper (II) oxide and zinc
12.17.3 Acidic oxides, carbon dioxide
12.17.3.1 Pass carbon dioxide through sodium hydroxide solution
12.17.3.2 Pass carbon dioxide through barium hydroxide solution
3.34 Prepare carbon dioxide, CO₂
3.34.1 Tests for carbon dioxide
3.34.2 Test the breath for carbon dioxide
3.34.3 Solubility of carbon dioxide in water
3.34.4 Reduce carbon dioxide with burning magnesium
3.34.5 Frozen carbon dioxide ("dry ice", "hot ice")
3.34.6 Soda-acid fire extinguisher
3.35 Carbon dioxide in the home
3.35.1 Washing soda
3.35.2 Baking soda
3.35.3 Baking powder
3.35.4 Yeast cells
3.36 Carbon dioxide and photosynthesis
3.37 Carbon dioxide and respiration
3.38 Carbon dioxide and fermentation for brewing
3.41.3 Reduce metal oxides to metals with hydrogen

---

12.10.0 Salts, sodium chloride "table salt"
See also 3.71.1: Solubility table and solubility rules
A salt is the product with water of the reaction of an acid with a base. A salt is a compound formed when the hydrogen ion of an acid is replaced by a metal ion or electropositive complex ion, e.g. NH₄.
An acid salt forms when an acid contains more than one replaceable hydrogen ion, e.g. H₂SO₄ and not all the hydrogen ions are replaced, e.g. NaH(SO₄)₂.
Salts are usually crystalline and are composed of positive and negative ions. You can prepare insoluble salt precipitates from pairs of solutions of salts by using the solubility rules. Sodium chloride is an ionic solid. Crystals of sodium chloride contain Na⁺ and Cl^- ions attracted to each other by strong ionic bonds in a crystal lattice. The crystals are hard and have high melting points and boiling points. When melted or in solution, sodium chloride conducts electricity, but the solid is a poor conductor of electricity.

---

12.10.1 Crystals of different salts
Dissolve different salts in water. Slowly evaporate the solution until salt crystals start to form. Add a crystal of salt to help crystallization. Describe the colour and shape of different salt crystals.

---

12.10.2 Prepare salts by different methods
See also 12.3.3: Reactions of dilute acids with metals, sulfuric acid and iron

---

12.10.2.1 Reactions of dilute acids with metals
Reactions with K and Na are too vigorous. No reaction for metals below hydrogen in the activity series.

---

12.10.2.2 Test the reactions of dilute acids with alkalis
See also 12.3.7: Reactions of dilute acids and hydroxides, sodium hydroxide
This method requires use of an indicator to prepare sure that no excess acid or alkali remains when the reaction is complete.

---

12.10.2.3 Reactions of acids with metallic oxides
See also 12.3.5: Reactions of dilute acids and basic oxides, copper (II) oxide
The reaction needs heat.

---

12.10.2.4 Reactions of acids with metal carbonates
See also 12.3.9: Reactions of dilute acids and common carbonates

---

12.10.2.5 Direct union of the elements sodium and chlorine

---

12.10.2.6 Add solutions of a salt to another salt
This is the only way to prepare an insoluble salt. In this type of reaction, the needed salt forms a precipitate. When solutions of two ionic substances are mixed and the ions of an insoluble salt are in this mixture, then a precipitate of the insoluble salt forms.
Prepare dilute solutions of different salts in separate test-tubes, e.g. barium nitrate, silver nitrate and lead nitrate. To each add a small quantity of dilute hydrochloric acid from a dropping tube. Note the colour and appearance of any precipitate that forms.
Repeat the procedure using 1. sodium chloride solution, 2. sodium sulfate solution, 3. dilute sulfuric acid.
silver ions(aq) + chloride ions(aq) --> silver chloride(s) [Silver chloride is insoluble in water]
lead ions(aq) + chloride ions(aq) --> lead chloride(s) [Lead chloride is insoluble in water]
sodium nitrate(aq) + copper (II) sulfate(aq) --> sodium ions(aq) + nitrate ions(aq) + copper ions(aq) + sulfate ions (aq) [No precipitate because both sodium sulfate and copper nitrate are soluble in water.]

---

12.10.2.6.1 Artificial gemstones, potassium sulfate, aluminium sulfate
Half fill a Petri dish with water. At one side, carefully pour a solution of potassium sulfate. At the other side carefully pour a solution of aluminium sulfate (swimming pool flocculent powder). Leave to allow crystals of potassium aluminium sulfate to form in the middle. Add a solution of lead chromate. The crystals will change colour like artificial gemstones.

---

12.10.3 Hydrolysis of sodium carbonate
Washing powders contain di-sodium tetraborate (III)-10-water (borax) + sodium carbonate and are alkaline in solution. Hydrolysis is a chemical reaction of a compound with water. Hydrolysis of salts is the reverse of neutralization. Salts of weak acids or weak bases hydrolyse when dissolved in water. Weak acids and weak alkalis dissociate very slightly. Solvation occurs when solvent molecules form bonds with a solute particle.
Dissolve sodium carbonate in water. Some hydrogen ions react to form the weak acid carbonic acid leaving excess hydroxyl ions in the solution. The solution turns red litmus blue.
salt + water --> acid + base
Na₂CO₃(aq) <--> 2Na⁺(aq) + CO₃^2-(aq)
H₂O(l) <--> H⁺(aq) + OH^-(aq)
2H⁺(aq) + CO₃^2-(aq) <--> H₂CO₃(aq)
Na₂CO₃(aq) + H₂O(l) <--> 2NaOH (s) + H₂CO₃(aq)

---

12.10.3.1 Hydrolysis of sodium hydrogen carbonate
Sodium hydrogen carbonate bicarbonate baking soda has a basic reaction and can be used to neutralize acids in fruit or neutralize bee stings.
Dissolve sodium hydrogen carbonate in water. A solution of sodium hydrogen carbonate in water turns red litmus blue.
NaHCO₃(aq) <--> Na⁺(aq) + HCO₃^-(aq)
H₂O(l) <--> H⁺(aq) + OH^-(aq)
HCO₃^- (aq) + H⁺(aq) <--> H₂CO₃(aq)

---

12.10.3.2 Hydrolysis of ammonium chloride
Dilute ammonia solution is only slightly dissociated because it is a very weak alkali. The ammonium ions react with hydroxyl ions to form undissociated dilute ammonia solution leaving excess of hydrogen ions. So the solution of ammonium chloride has pH value of about 6.
Dissolve ammonium chloride in water.
NH₄Cl (aq) <--> NH₄⁺(aq) + Cl^-(aq)
NH₄⁺(aq) + OH^-(aq) <--> NH₄OH(s)
NH₄Cl (aq) + H₂O (aq) <--> NH₄OH (s) + H⁺(aq) + Cl^-(aq)

---

12.10.3.3 Reactions of salts and water
Water and salts do not usually react but sometimes hydrolysis occurs and the solution becomes either acidic or alkaline.
Dissolve a small amount of the following salts in demineralized water and test each solution with red and with blue litmus paper: sodium chloride, sodium carbonate, copper (II) sulfate, sodium acetate, iron chloride. copper (II) sulfate and iron chloride give acidic solutions. Sodium carbonate and sodium acetate give alkaline solutions. Sodium chloride solution is neither acidic nor alkaline.

---

12.10.4 pH of different salts dissolved in water
Add three drops of Universal Indicator to 5 mL of 0.2 M salt solution
Salt Colour pH: NH₄Cl, orange-red pH 5, NaCl yellow-green pH 7, Na₂HPO₄ blue-green pH 9_, KNO₂ blue pH 9.5, Na₂CO₃ violet pH 10, Na₂S red-violet pH10.5

---

12.10.5 Hydrolysis of iron (III) chloride
Iron chloride exists as anhydrous iron (III) chloride (FeCl₃) and FeCl₃.6H₂O. Iron (III) chloride is rapidly hydrolysed in moist air and is partially hydrolysed in solution. Hydrolysis can be suppressed by addition of HCl. Iron (III) chloride-6-water. Fe(OH)₂, green, is oxidized to Fe(OH)₃, brown, in moist air.
Dissolve iron (III) chloride in boiling water. Add drops of dilute ammonia solution The reaction forms a brown-red precipitate of iron (III) hydroxide.
FeCl₃(aq) + NH₄OH(aq) --> Fe(OH)₃(s) + NH₄Cl(aq)
Heat to evaporate some solution. The reaction forms a red-brown precipitate of iron (III) hydroxide.
FeCl₃(aq) + 3H₂O(l) --> Fe(OH)₃(s) + 3HCl(l)
Pour the clear saturated solution into hot water. The reaction forms a red precipitate of hydrated iron (III) oxide.
2FeCl₃(aq) + 3H₂O(l) --> Fe₂O₃(s) + 6HCl(l)
Add drops of sodium hydroxide solution. The reaction forms a red-brown precipitate of iron (III) hydroxide.
FeCl₃(aq) + 3NaOH(aq) --> Fe(OH)₃(s) + NaCl(aq)

---

12.10.6 Prepare an acid salt, sodium hydrogen sulfate
An acid salt is the salt of an acid containing more than one acidic hydrogen, e.g., H₂SO₄, which has not all the hydrogens replaced by positive ions.
Add drops of 2 M sulfuric acid to 2 M sulfuric acid. Count the drops until the solution is neutral to litmus.
Repeat the experiment by adding half the number of drops of acid.
NaOH(aq) + H₂SO₄(aq) -->NaHSO₄(aq) + H₂O(l)

---

12.10.7 Buffer solutions
The pH value of buffer solutions changes very little when acids or alkalis are added or when diluted with water. Although the salts of weak acids are completely dissociated into ions, weak acids do not dissociate completely. A buffer solution contains a weak acid and the salt of the weak acid, e.g. H₂CO₃ / HCO₃^- (carbonic acid / sodium hydrogen carbonate). By mixing an acid with its conjugate base, definite hydrogen ion concentrations, within a certain range depending on the dissociation constant of the acid, are obtainable. Such solutions have the advantage that evaporation will not affect the value of [H+] for the ratio [acid] / [base] remains constant. Contamination by small quantities of acidic or basic impurities will not affect the pH. If an acid is added to a buffer solution, the H⁺ added react with the HCO₃^-. If a base is added to a buffer solution, the OH^- reacts with the undissociated H₂CO₃ to form the salt and water. Natural body fluids are buffered.

---

12.10.7.1 Dilution effect on buffer solutions
Add 1 mL of 0.01M HCl to 1 mL of water. The pH value changes from 7 to 5.

---

12.10.7.2 Natural buffers
Add 1 mL of 0.01M HCl to one cube of beef soup (beef cube infusion). Almost no pH change occurs because of buffering action.

---

12.10.7.3 Prepare a buffered solution
(Methyl orange. pH 2.5 - red, pH 3.5 - straw colour, pH 4.5 - orange). Add a drop of methyl orange to: 1. deionized water. It turns yellow. 2. deionized water + 5 drops ethanoic acid (acetic acid). It turns pink. 3. deionized water + 5 drops ethanoic acid + crystals of sodium acetate-3-water. It turns yellow. The 3. solution is buffered, so it does not turn pink as in the 2. solution.

---

12.10.7.4 Salt effect on buffer solutions
Add drops of methyl orange to: 1. deionized water. The solution turns yellow. 2. Dilute hydrochloric acid. The solution turns red. 3. Dilute ethanoic acid (acetic acid). The solution turns slightly red. 4. Very dilute acetic acid. The solution turns red. The very dilute acetic acid is red as with dilute hydrochloric acid. 5. Half the very dilute acetic acid solution + sodium chloride crystals. The solution turns pale red. The salt effect prevents reformation of molecular acetic acid.

---

12.10.8 Prepare solutions of pH values 3 to 11 with buffer solutions
The pH value of a buffer solution does not alter for small additions of acid or alkali, e.g. a mixture of highly ionized sodium acetate, CH₃COONa, and partly ionized acetic acid, CH₃COOH or HAc.
1. If add hydrogen ions to the solution, the HAc that forms is undissociated and so H⁺ are removed from the solution.
H⁺ + Ac^- --> HAc
2. If add alkali to the solution, more HAc dissociates to form hydrogen ions that combine with the hydroxyl ions to form H₂O that is undissociated and so OH^- ions are removed from the solution.
HAc -> H⁺ + Ac^-
H⁺ + OH^- --> H₂O
1. Prepare solutions with hydrogen ion concentrations of 10^-3 to 10^-6 g ions per litre
Use 1. 0 1M acetic acid solution 2. 0.1 M sodium acetate solution (13.6 g of crystalline sodium acetate, CH₃COONa.3H₂O per litre)
1.1 Hydrogen ion concentration 10^-3: 1 litre 0.1 M acetic acid and 18 mL 0.1 M sodium acetate
1.2 Hydrogen ion concentration 10^-4: 1 litre 0.1 M acetic acid and 180 mL 0.1 M sodium acetate
1.3 Hydrogen ion concentration 10^-5: 555 mL 0.1 M acetic acid and 1 litre 0.1 M sodium acetate
1.4 Hydrogen ion concentration 10^-6: 55 mL 0.1 M acetic acid and 1 litre 0.1 M sodium acetate

---

2. Prepare solutions with hydrogen ion concentrations of 10^-7 to 10^-11 g ions per litre
Use 1. Disodium phosphate solution (Na₂HPO₄): Dissolve 0.1 mole of the crystalline salt Na₂HPO₄, 35.8 g 2. 0.1 M hydrochloric acid 3. 0.1 M sodium hydroxide.
Na₂HPO₄ solution. 0.1 M HCl, 0.1 M NaOH
2.1 Hydrogen ion concentration 10^-7: 1 litre Na₂HPO₄ solution and 322 mL 0.1 M HCl solution
2.2 Hydrogen ion concentration 10^-8: 1 litre Na₂HPO₄ solution and 47 mL 0.1 M HCl solution
2.3 Hydrogen ion concentration 10^-9: 1 litre Na₂HPO₄ solution and 5 mL 0.1 M HCl solution
2.4 Hydrogen ion concentration 10^-10: 1 litre Na₂HPO₄ solution and 3.6 mL 0.1 M NaOH solution
2.5 Hydrogen ion concentration 10^-11: 1 litre Na₂HPO₄ solution and 3.6 mL 0.1 M NaOH solution

---

3. To show the changes in colour shown by the indicator in solutions of various hydrogen ion concentrations arrange three rows of nine test-tubes on white paper.

10-3	10-4|	10-5	10-6	10-7|	10-8	10-9	10-10	10-11
MO	MO	MO	MO	--	PH	PH	PH	PH
MR	MR	MR	MR	MR	TP	TP	TP	-
-	LI	LI	LI	LI	LI	LI	-	-

MO = methyl orange, PH = Phenolphthalein, PP. LI = Litmus, MR = Methyl red, Thymol phthalein = TP.

---

4. Add two drops of the indicators and note the colour changes observed. The solutions marked in bold show the hydrogen ion concentrations where colours are most noticeable.

---

5. Show that the solutions are comparatively stable with a solution of M / 1000 HCl (pH = 3, and M / 1000 NaOH (pH = 11). Add two drops of universal indicator to each solution. Pour out the same volume of buffer solutions of pH 3 and pH 11 above and add two drops of universal indicator to them. Add a drop of acid or alkali in turn to each of the four solutions and note the colour change and estimate the alteration in pH value. Note the rapid change with the M / 1000 HCl and NaOH but little change in the buffer solutions pH 3 and pH 11.

---

12.10.9 Show the effect of a buffer salt
A buffer salt is essentially a highly ionized salt of a weak acid.
1. Add two drops of universal indicator to 10 mL of 0.1 M sodium hydroxide solution. Titrate the mixture with 0.1 M hydrochloric acid. Note the colour changes that indicate the rapid change of pH about the equivalence point.
2. Add 5 g of sodium acetate to 10 mL of 0.1 M sodium hydroxide solution, then two drops of indicator. Titrate the mixture with 0.1 M hydrochloric acid. Note hydrogen ions are added but the green colour of the indicator persists because the pH remains constant over a long period of addition of hydrogen ions. The buffer salt, sodium acetate, is highly ionized and gives acetate ions. The hydrogen ions from the hydrochloric acid form molecular acetic acid instead of increasing the hydrogen ion concentration in the solution.
NaAc --> Na⁺ + Ac^-
H⁺ + Ac^- <--> HAc
When a large excess of hydrogen ions is added, the pH of the solution decreases. Adding a strong alkali to a highly ionized salt of a weak base does not at first increase the pH of the mixture. Ammonium chloride solution gives ammonium ions that react with the added hydroxide ions of a strong alkali to form molecular "ammonium hydroxide". [Not "ammonium hydroxide, NH₄OH". Ammonia solution. is shown as NH₃(aq) because "NH₄⁺" ions and "OH^-" ions can be detected,
but "NH₄OH" cannot be detected.] The pH of the solution rises only after an excess of alkali is added.

---

12.10.10 Change in pH near the equivalence point
1. Add two drops of universal indicator to 10 mL of sodium hydroxide solution. Titrate the mixture with hydrochloric acid. Note the rapid change of colour from blue-green at pH about 8.5 to orange-red at pH about 4. When a strong alkali is titrated against a strong acid, the indicator indicates the equivalence point with negligible error.
Repeat the experiment with a low pH indicator, e.g. methyl orange, and a high pH indicator, e.g. phenolphthalein, and note the slight difference.
2. Add two drops of universal indicator to 10 mL of sodium hydroxide solution and titrate the mixture with acetic acid. Note that when the equivalence point is reached, the pH is about 8.5. Note also the considerable excess of acid necessary to approach the orange colour of pH 4, showing that only a high pH indicator is efficient in the titration of a strong alkali with a weak acid.
Repeat the experiment using phenolphthalein and methyl orange.

---

12.10.11 pH values of solutions of salts
A normal salt is one in which the replaceable hydrogen atoms of an acid have been completely replaced by a metal. However, a normal salt is not necessarily a neutral salt since hydrolysis may occur, e.g. sodium carbonate is alkaline in solution but ammonium chloride is acidic in solution.
Half fill seven test-tubes with water and add two drops of universal indicator. Add 1.25 mL of the following salts: sodium carbonate, sodium sulfite, sodium chloride, ammonium chloride, aluminium chloride, borax, iron (II) sulfate. Note the pH value according to the colour produced. Warm the solutions and note whether this increases the hydrolysis, in some cases producing greater divergence from neutrality.

---

12.10.12 Fireproofing
1. Fireproof with ammonium chloride solution. Soak cloth in the solution and let it dry. Hold the cloth over a flame. The cloth will only smoulder. When the flame is removed, the burning stops. Fabrics are often treated this way to prepare them fireproof.
2. Fireproof paper with alum. Soak half a piece of paper or cloth in an alum solution, potash alum Al₂(SO₄)₃.K₂(SO₄).24H₂O [also shown as KAl(SO₄)₂.12H₂O]. Set fire to the untreated half and observe the flames when they reach the treated half. Alum acts as a flame retardant. It forms a buffer between a flammable substance and a source of ignition, being incombustible with its many molecules of water that it can lose at 200^oC
3. Fireproof wood with sodium silicate solution. Hold a match by the head and dip it into a mixture of 1 part sodium silicate solution and 3 parts water. Let the match dry completely. Light the match. The wood burns, but when the flame reaches the coating of sodium silicate, it is extinguished.
4. Fireproof with sodium thiosulfate (hypo). Soak 2 wooden matches in a concentrated solution of sodium thiosulfate for 2 hours. Strike the matches and see whether the wood will burn

---

12.11.5.0 Tests for anions in unknown solution, tests for acid radicals in solution
Before testing a solution for acidic radicals remove heavy metals that may interfere with the tests, leaving only sodium, potassium or ammonium in solution, e.g. to test for a sulfate radical in solution, add dilute hydrochloric acid and barium chloride solution. A white precipitate of barium sulfate indicates the presence of a sulfate.
Ba²⁺ + SO₄^2- --> BaSO₄ (s)
However, if the solution already contains the silver ion, the white precipitate is silver chloride.
Ag⁺ + Cl^- --> AgCl (s)
Boil 1 g the finely divided unknown solid with sodium carbonate solution to precipitate heavy metals as carbonates, or as hydroxides by hydrolysis. Filter off the precipitates. Copper may rarely form a soluble double carbonate. The acidic radicals, originally combined with the heavy metals, are now in the filtrate as the sodium salts if double decomposition has occurred, e.g. a mixture containing barium chloride and calcium nitrate:
BaCl₂ + Na₂CO₃ --> BaCO₃ (precipitate) + 2NaCl (solution)
Ca(NO₃)₂ + Na₂CO₃ -->CaCO₃ (precipitate) + 2NaNO₃ (solution)
The filtrate is alkaline with excess sodium carbonate and now must be made acidic, e.g. barium chloride use hydrochloric acid, with silver nitrate use nitric acid so you do not add the radical you are testing for. If the solutions are not made acid, the sodium carbonate precipitates the metal of the testing reagent as a heavy metal carbonate.

---

12.11.5.1 Tests for acetates, CH₃COO^- and (CO₃)^2-
1. Add to 5 drops of original solution drops of dilute HCl, or HNO₃ if using a Pb salt. If effervescence occurs, pass the gas through lime water. A milky precipitate indicates (CO₃)^2-. If effervescence does not occur, heat the solution. The odour of vinegar indicates CH₃COO^-.
2. Neutralize with dilute nitric acid and ammonia, then add iron (III) chloride solution. A blood red colour, lost by adding hydrochloric acid, indicates an acetate.
3. Add an equal volume of alcohol and then drops of concentrated sulfuric acid. Heat gently and smell the vapour. The fruity smell of ethyl acetate indicates the presence of an acetate.
CH₃COONa + C₂H₅OH + H₂SO4 --> CH₃COOC₂H₅ + NaHSO₄ + H₂O

---

12.11.5.2a Tests for antimonates, borates, oxalates
Add excess dilute nitric acid, followed by silver nitrate solution. Filter off the precipitate. Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a white precipitate forms that indicates antimonate, borate, or oxalate in the filtrate.

---

12.11.5.3 Tests for arsenates
1. Add dilute nitric acid and excess ammonium molybdate solution. Heat to boiling. A yellow precipitate of ammonium arsenomolybdate (NH₄)₃AsO₄.12MoO_3, indicates arsenate.
2. Add excess dilute nitric acid, followed by silver nitrate solution. Filter off the precipitate. Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a brick-red precipitate forms that indicates arsenate in the filtrate.

---

12.11.5.4 Tests for bicarbonates
Add magnesium sulfate solution. A white precipitate in the cold indicates the presence of carbonate. No precipitate in the cold, but a white precipitate on boiling, confirms bicarbonate. If the original solid is insoluble in water, an aqueous suspension of it may be boiled. A solution that produces carbon dioxide indicates the presence of bicarbonate.

---

12.11.5.5 Tests for borates
1. Dissolve 1g of boric acid in 10 mL of ethanol. Use a trigger pump operated spray bottle, e.g. window cleaner spray bottle, to spray the solution onto a roaring Bunsen burner flame. A green flame indicates borates.
2. Add concentrated sulfuric acid to the unknown substance then pour into methylated spirit into an evaporating dish while stirring with a glass rod. Heat the evaporating dish and light the vapour rising it. A green colour in the flame produced by the volatile compound, ethyl borate, indicates borate radical.
Na₂B₄O₇ + H₂SO₄ + 5H₂O --> Na₂SO₄ + 4H₃BO₃
H₃BO₃ + 3C₂H₅OH --> B(OC₂H₅)₃ + 3H₂0
The test may not work for few minerals containing boron, e.g. borosilicates.
3. To confirm borate, acidify the solution and test with turmeric paper. Dry the paper over a small flame. The change of colour from yellow to brown, which becomes blue or blue-black in caustic soda solution indicates a borate.
4. Test for borate, oxalate, antimonate
Add excess dilute nitric acid, followed by silver nitrate solution. Filter off the precipitate. Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a white precipitate forms that indicates the presence of antimonate, borate, or oxalate in the filtrate.

---

12.11.5.6 Tests for bromides
1. Add excess dilute nitric acid, followed by silver nitrate solution. A pale yellow precipitate of silver bromide, sparingly soluble in ammonia, indicates the presence of the bromide radical.
Ag⁺ + Br^- --> AgBr (s)
2. To confirm the bromide radical, heat the solid with manganese dioxide and concentrated sulfuric acid and observe the dark red vapour of bromine.

---

12.11.5.7 Tests for carbonates
Add magnesium sulfate solution. A white precipitate in the cold confirms carbonate. No precipitate in the cold, but a white precipitate on boiling, confirms bicarbonate. If the original solid is insoluble in water, an aqueous suspension of it may be boiled. If the solution produces carbon dioxide, a bicarbonate is indicated.

---

Tests for chlorides
See also: 12.19.8.1: Reactions of chlorides
1. Add excess dilute nitric acid, followed by silver nitrate solution. A white precipitate of silver chloride, soluble in ammonia, indicates the presence of chloride radical.
Ag⁺ + Cl^- --> AgCl (s)
AgCl + 2NH₃ --> Ag(NH₃)₂Cl (soluble silver amine)
2. To confirm the presence of the chloride radical
3. Prepare chromyl chloride, CrO₂Cl₂

---

12.11.5.9 Tests for chromates
1. Acidify with dilute nitric acid, add ammonia solution, NH₃(aq) ("ammonium hydroxide") until just alkaline. Heat to boiling then divide intro 2 parts. To one part add the solution. silver nitrate solution. A crimson red precipitate, soluble in dilute nitric acid indicates chromate.
2Ag⁺ + CrO₄^2- --> Ag₂CrO₄ (s)
To the other part add barium chloride solution. A yellow precipitate soluble in hydrochloric acid confirms chromate.
Ba²⁺ + CrO₄^2- --> BaCrO_{4 (s)}
2. Acidify the sodium carbonate extract with dilute sulfuric acid. Add drops of amyl alcohol then hydrogen peroxide solution. Shake then leave to stand. A blue colour in the alcohol confirms chromate.
3. Add excess dilute nitric acid, followed by silver nitrate solution. Filter off the precipitate. Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a crimson red precipitate forms that indicates chromate in the filtrate.

---

12.11.5.10 Tests for halides, Cl^-, Br^-, I^-
Mix 1 g of unknown solid with 1 g of MnO₂ add concentrated H₂SO₄ then heat. Orange-red gas indicates Br^-. Violet layer of gas indicates I^-. Yellow-green gas that turns KI / starch paper blue to indicate Cl^-.

---

12.11.5.11 Tests for hydroxides
Add one drop of sodium hydroxide solution to ten drops of the unknown solution. 1. A white or glassy precipitate indicates Al³⁺, Bi³⁺, Cd²⁺, Mg²⁺, Mn²⁺, Pb²⁺, Zn²⁺, Sn²⁺. 2. A green precipitate indicates Fe(OH)₂, Ni²⁺, Cr³⁺. 3. A brown precipitate indicates Ag⁺ and Fe(OH)₃. 4. A blue precipitate indicates Cu²⁺ and Co²⁺. 5. The reaction with Ca²⁺ forms a slightly soluble white precipitate. If the reaction forms no precipitate, heat the solution to identify the presence of NH₄⁺ from the odour of ammonia

---

12.11.5.12 Tests for iodides.
1. Add excess dilute nitric acid, followed by silver nitrate solution. A yellow precipitate of silver iodide, insoluble in ammonia, indicates the presence of the iodide radical.
Ag⁺ + I^- --> AgI (s)
2. To confirm the iodide radical, heat the solid with manganese dioxide and concentrated sulfuric acid and observe the violet vapour of iodine.

---

12.11.5.13 Tests for nitrates
1. First test: When the cation is not a salt of Na⁺, NH₄⁺ or K⁺, remove it as insoluble carbonate. Add 10 mL Na₂CO₃ solution to 1 g of the solid salt, boil, filter and prepare up to 2 mL with deionized water. Add to 5 drops of unknown solution, 5 drops of water, 5 drops concentrated H₂SO₄ and Cu foil. Brown fumes of nitrogen dioxide and a blue-green solution indicate NO₃^-.
2. Second test: Add to 5 drops of unknown solution in an evaporating basin, 3 drops of concentrated sulfuric acid and a crystal of iron (II) sulfate. A purple colour on the crystal indicates NO₃^-.
3. This test is called the brown ring test. Add excess of cold dilute sulfuric acid to the unknown solution then add excess freshly prepared iron (II) sulfate solution. Transfer the solution to a boiling tube to a depth of 2 cm. Fix the boiling tube in a sloping position then very carefully pour concentrated sulfuric acid down the sloping side of the tube to form a separate 2 cm layer beneath the solution. Observe a brown ring at the junction of the acid and unknown solution. The nitrate and the concentrated sulfuric acid first form nitric acid to be reduced by iron (II) sulfate to nitric oxide. The nitric oxide reacts with more iron (II) sulfate to form the brown compound, NO.2FeSO₄. Carefully shake the boiling tube to spread the brown colour. The solution becomes warm as the acid and water mix and the brown colour disappears as the unstable brown compound decomposes.
2FeSO₄ + 2NaNO₃ + 5H₂SO₄ --> 2NaHSO₄ + 3Fe₂(SO₄)₃ + 4H₂O + 2NO(g)
NO + 2FeSO₄ --> NO.2FeSO₄ (brown colour)
NO.2FeSO₄ --> NO + 2FeSO₄ (brown colour disappears)
4. If a bromide or iodide is in the unknown solution, a ring due either to free bromine or to free iodine forms and the iron (II) sulfate is not part of this reaction. However, if bromide or iodide is already known to be in the unknown solution, add silver sulfate solution to precipitate the bromide or iodide as a silver salt and then test the filtrate for the nitrate ion.
5. If a nitrite is in the unknown solution, a diffuse brown ring forms. To eliminate nitrite, add a concentrated solution of urea, then dilute sulfuric acid and warm until effervescence of nitrogen stops. Then test for nitrate.
6. Heat a mixture of the original solid with copper and drops of concentrated sulfuric acid. The nitrate radical reacts with concentrated sulfuric acid to form nitric acid which reacts with copper to produce brown nitrogen dioxide gas. The brown gas indicates the nitrate radical.
Cu + 4HNO₃ ---> Cu(NO₃)₂ + 2H₂O + 2NO₂ (g)

---

12.11.5.14 Tests for oxalates
1. Dissolve the unknown substance in water, add excess calcium chloride solution and heat to boiling. Decant and wash the remaining precipitate of calcium oxalate with warm dilute sulfuric acid. Add a drops of potassium permanganate solution which is decolorized.
2KMnO4 + 3H₂SO₄ + 5H₂C₂O₄ --> K₂SO₄ + 2MnSO₄ + 8H₂O + 10CO₂
2. Test for oxalate, antimonate, borate. Add excess dilute nitric acid, followed by silver nitrate solution. Filter off the precipitate. Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a white precipitate forms that indicates antimonate, borate, or oxalate in the filtrate.

---

12.11.5.15 Tests for phosphates
1. Add dilute nitric acid and excess ammonium molybdate solution. Heat but do not boil. A yellow coloration, with precipitate of ammonium phosphomolybdate on standing (NH₄)₃PO₄.12MoO_3, indicates phosphate.
2. Add excess dilute nitric acid, followed by silver nitrate solution. Filter off the precipitate. Add ammonia to the filtrate solution drop by drop If the filtrate contains excess silver nitrate, a yellow precipitate forms that indicates phosphate in the filtrate.

---

12.11.5.16 Tests for sulfates
1. Add to 5 drops of unknown solution 2 drops of HCl, heat then add 3 drops of BaCl₂ solution. A white precipitate indicates SO₄^2-.
Ba²⁺ + SO₄^2- --> BaSO₄ (s)
2. Add excess dilute hydrochloric acid, and then barium chloride solution. A white precipitate of barium sulfate proves sulfate radical present.
3. To confirm the presence of sulfates, heat the unknown with fusion mixture on a charcoal block and test the residue on a wet silver surface. A black stain of silver sulfide indicates a sulfide formed by partial reduction of the sulfate. This test is not applicable if sulfide is in the unknown substance.

---

12.11.5.17 Tests for sulfides
Add a solution of lead acetate. A black precipitate indicates sulfide.

---

12.11.5.18 Tests for sulfites
Add barium chloride solution. A white precipitate, soluble in hydrochloric acid, indicates sulfite.

---

12.14.0 Activity series of metals as reducing agents
The activity series is also called reactivity series or electrochemical series.
Decreasing activity from left to right: potassium, sodium, barium, calcium, magnesium, aluminium, zinc, iron, tin, lead (hydrogen) copper, mercury, silver, platinum, gold
Metals above lead in the activity series react with acids and liberate hydrogen. However, nitric acid and concentrated sulfuric acid react with metals above platinum but do not produce hydrogen. Reactions of acids with metals are exothermic and the higher the metal in the activity series, the greater the heat liberated in its reaction with an acid.
1a = reaction with cold water to give the oxide and hydrogen
1b = reaction with hot water to give the oxide and hydrogen
1c = reaction with steam to give the oxide and hydrogen
2a = reaction with air - when heated form peroxides
2b = reaction with air - when heated as powders form oxides
3a = react with dilute hydrochloric acid or sulfuric acid to form hydrogen gas and metal ions and react with concentrated nitric acid or sulfuric acid to produce metal ions and nitrogen dioxide or sulfur dioxide
3b = react with concentrated nitric acid or sulfuric acid to produce metal ions and nitrogen dioxide or sulfur dioxide
3c = react with aqua regia (concentrated nitric acid and hydrochloric acid)

K

1a

2a

3a

-

Zn

1c

2b

3a

-

Hg

-

2b

3b

-

Ba

1a

2a

3a

-

Fe

1c

2b

3a

-

Ag

-

-

3b

-

Sr

1a

2a

3a

-

Ni

1c

2b

3a

-

Pt

-

-

-

3c

Na

1a

2a

3a

-

Sn

-

2b

3a

-

Au

-

-

-

3c

Ca

1a

2a

3a

-

Pb

-

2b

-

3b

-

-

-

-

-

Mg

1b

2b

3a

-

H

-

-

-

-

-

-

-

-

Al

1c

2b

3a

-

Cu

-

2b

-

3b

-

-

-

-

-

Reaction with water: All metallic elements except Sn, Pb, Cu, Hg, Ag, Au and Pt react with cold water or hot water or steam.
Reaction with air or oxygen: All elements except Ag, Au and Pt react with air. K, Na and Ca form peroxides. The other elements form oxides, when heated as powders.
Reaction with dilute acids: Pb, Cu, Hg, Ag, AU and Pt do not react with dilute HCl or HNO₃. Pt and Au react with aqua regia. Metals react with dilute acids to form hydrogen and the metal ion.
Reaction with concentrated oxidizing acids: Au and Pt do not react with concentrated HNO₃ or H₂SO₄. Reactions form the metal ions of high oxidation number and sulfur dioxide if H₂SO₄
Reactions form nitrogen dioxide if HNO₃, e.g. copper has two oxidation numbers, number 1 (Cu⁺¹) and number 2 (Cu²⁺)

---

12.14.1 Zinc displaces lead from lead nitrate solution
A metal displaces a metal lower in the activity series from its salt solutions. The more active metal atoms lose electrons more easily to go into solution as ions. The less active metal ions attract electrons more easily to leave the solution as metal atoms. The position of the metal in the activity series represents its relative ease of oxidation, i.e. ease of losing electrons to form ions. The most active metals replace hydrogen from water. Metals that replace hydrogen from dilute acids are placed above hydrogen. Metals that do not replace hydrogen from such acids are placed below hydrogen. These metals may be oxidized by the oxidizing acids nitric acid and hot concentrated sulfuric acid. Gold and platinum do not react with the oxidizing acids, but do react with aqua regia (a mixture of concentrated hydrochloric acid and concentrated nitric acid in ratio 3:1 by volume).
Put a piece of granulated zinc in a test-tube containing lead (II) nitrate solution. The zinc becomes covered with a solution of metallic lead. The zinc granule becomes corroded. Zinc displaces lead from a solution of its salt.

---

12.14.2 Metals in copper (II) sulfate solution
A metal higher in the activity order is needed to displace copper metal from a solution of copper ions.
12.14.2.1 Magnesium, or zinc, in copper (II) sulfate solution
Magnesium or zinc displaces copper that is lower in the activity series from its salt copper (II) sulfate.
Use magnesium ribbon or zinc dust in a test-tube of copper (II) sulfate solution. The reaction can be vigorous with the magnesium. Copper metal deposits and the blue colour gradually disappear as the copper ion is displaced by the more reactive metal that is higher in the activity series. The reaction loses heat. When the solution is colourless, decant the solution leaving red copper powder at the bottom of the test-tube.
Mg(s) + CuSO₄(aq) -->MgSO₄(aq) + Cu(s)
Mg loses electrons: Mg --> Mg²⁺ + 2e^- (oxidation)
Cu gains electrons: Cu²⁺ + 2e^- --> Cu (reduction)

---

12.14.2.2 Iron in copper (II) sulfate solution
Clean a large iron nail with emery cloth. Put it in a test-tube containing copper (II) sulfate solution. The reaction forms a coating of copper on the iron nail as copper leaves the solution. The colour of the solution changes from blue to green iron enters the solution as ions. The iron nail is corroded. Iron displaces copper from a solution of its salt.

---

12.14.2.3 Iron and zinc with copper (II) sulfate solution
1. Add 10 g of copper (II) sulfate solution to 50 mL of water in two beakers. Add shiny iron nails to beaker 1. Add shiny pieces of zinc metal to beaker 2. Leave to stand and after 2 hours note any change in colour of the solution and any precipitate.
2. Add iron nails to the solution containing the zinc and add shiny pieces of zinc to the solution containing the iron nails. Notice any further reactions that take place.
CuSO₄ + Zn --> ZnSO₄(aq) + Cu(s)
CuSO₄ + Fe --> FeSO₄(aq) + Cu(s)
FeSO₄ + Zn --> ZnSO₄(aq) + Fe(s)
ZnSO₄ + Fe --> no reaction

---

12.14.2.4 Zinc in lead nitrate solution and iron in copper (II) sulfate solution
Clean a small strip of zinc and an iron nail with emery cloth. Prepare up separate solutions of lead (II) nitrate and copper (II) sulfate. Put the zinc in the lead nitrate solution and put the iron in the copper (II) sulfate solution. After a few minutes remove the metal strips and observe the appearance of each. Note a copper coating on the iron nail. Note the crystals of metallic lead on the zinc. After leaving the metals in the solution for a longer time you will notice that the original metal has corroded. The copper (II) sulfate solution the blue colour will be gradually replaced by a dirty green colour.

---

12.14.2.5 Zinc and copper in sulfuric acid
1. Hold a clean strip of zinc in dilute sulfuric acid. If the zinc is very pure, few bubbles of hydrogen gas will be evolved from its surface. Remove the zinc and hold a strip of copper in the acid. No gas is evolved.
2. Put both metal strips in the acid so that an edge of the zinc is in contact with the copper. Copious bubbles of gas are given off from the copper plate and practically none from the zinc.

---

12.14.2.6 Activity of metals and tendency to form ions
Dip pairs of metal into sodium chloride solution. Connect the metals to a voltmeter and note the direction of current flow. The more reactive metal forms the negative pole and so electrons flow from it. Test zinc with copper, lead, iron and magnesium. Test copper with lead, magnesium and iron. Test lead with iron and magnesium. Test iron with magnesium. For each pair of metals, note which metal forms the positive terminal, which metal forms the negative terminal and the voltage for each combination.

---

12.15.0 Reactions of metals with water
1. Metals act as reducing agents in displacing hydrogen from water.
2. K, Ba and Na displace hydrogen from cold water.
3. K reacts violently and forms hydrogen gas that catches alight and burns with a pink flame.
4. Ca reacts slowly and the solution turns milky because of the formation of calcium hydroxide.
5. Mg reacts slowly with cold water and quickly with hot water.
6. Al, Zn, Fe and Ni react with steam to produce oxide and hydrogen.
7. Sn, Pb, Cu, Hg, Ag, Au and Pt do not react with water.

---

12.15.1 Reactions of metals with water, Cu, Zn, Fe, Mg, Al
See diagram 12.15.1B
If metals are not pure, some reactions may be caused by the impurity.
Boil deionized water for 5 minutes to remove dissolved air leave it to cool then pour into test-tubes. Put in the test-tubes pieces of freshly polished: copper, zinc, iron, magnesium, aluminium. Leave for 10 minutes. Observe any changes in the metal or water. When you see bubbles on the metals, put the metal under an inverted test-tube of water and leave for two days to collect the gas. Test the collected gas with litmus paper, limewater, and a lighted splint. The bubbles are hydrogen. Calcium reacts slowly then sinks. Magnesium reacts very slowly in cold water, but reacts vigorously in steam.
Ca(s) + 2H₂O(l) -->Ca(OH)₂(aq) + H₂(g)

---

12.15.2 Reactions of sodium with water
See diagram 12.15.2A
BE CAREFUL! USE A PIECE OF SODIUM THE SIZE OF A WHEAT OR RICE GRAIN.
DO NOT POINT THE OPEN END OF A test-tube CONTAINING SODIUM AND WATER AT ANYBODY!
The experiment can be done on an overhead projector or drop the sodium in a swimming pool.
1. Put water into a Petri dish and add a few drops of phenolphthalein. Put the Petri dish on an overhead projector. Use forceps to take a small lump of metallic sodium stored under kerosene and wipe with a piece of filter paper. Cut a piece of sodium the size of a red bean from the lump and then put it in the Petri dish. Observe from the screen that the sodium grain floats on the water to melt into a small sphere, which moves very rapidly in all directions, becoming smaller and smaller. At last, the sodium sphere wholly disappears and the solution in the dish also changes its colour from colourless to red.
2. Put a piece of sodium in 2 cm of water in a test-tube.
BE CAREFUL! WATCH THE REACTION THROUGH THE SIDE OF THE test-tube, BUT DO NOT LOOK DOWN THE test-tube!
Do not point the open end of the test-tube at anybody! Test for hydrogen with a lighted splint. Test the contents with litmus paper or phenolphthalein. The reaction forms sodium hydroxide.
3. Pour a 3 mm layer of kerosene (paraffin oil) on to the surface of water in a test-tube.
BE CAREFUL! WATCH THE REACTION THROUGH THE SIDE OF THE test-tube, BUT DO NOT LOOK DOWN THE test-tube!
DO NOT POINT THE OPEN END OF THE test-tube AT ANYBODY!
Drop a small piece of sodium into the test-tube. The sodium sinks in the kerosene and floats in the water. The layer of kerosene should be shallow enough to let the top of the sodium protrude above the surface. The reaction of sodium with water is much slower than if the sodium had been dropped directly on to the water. A small area of the sodium suddenly reacts which causes a stream of bubbles to appear. The irregular shape of the sodium changes to that of a sphere. The sodium melts because the reaction forms heat. Something dissolving in the water below the sodium can be seen. Some smoke where the hot sodium is above the kerosene level suggests a reaction with air. Test the gas bubbles for hydrogen.
4. Support a short length of glass tubing vertically with one end 2 cm below the surface of water. Test the water with moist litmus paper. Add litmus paper to the water. The upper part of the tube must be quite dry. Drop a piece of sodium down the tube. Note the reaction. Test for hydrogen with a lighted splint. Test the water with litmus paper. The water becomes alkaline.
2Na(s) + 2H₂O(l) -->H₂(g) + 2NaOH(aq)

---

12.15.3 Reactions of metals with steam
See diagram 12.15.3
Put wet cotton wool or glass wool at the bottom of a test-tube. Put another small piece of cotton wool or glass wool half way up the tube. Clean and polish a piece of magnesium ribbon and put it on the upper plug. Insert a 1-hole stopper fitted with a glass tube. Use a Bunsen burner to heat the lower cotton wool or glass wool until steam comes off. Use a second Bunsen burner to heat the magnesium ribbon.
Test for hydrogen with a lighted splint. Repeat the experiment with cleaned aluminium, copper wire, and iron wire. When heated in steam, magnesium, aluminium and iron react, but not copper.
Mg(s) + H₂O(g) -->MgO(s) + H₂(g)Al(s) + H₂O(g) -->Al₂O₃(s) + H₂(s)
Fe(s) + H₂O(g) -->Fe₂O₃(s) + H₂(s)

---

12.15.4 Reactions of metals with water
Prepare test-tubes containing deionized water or demineralized water. Boil the water then allow to cool. Put small pieces of freshly polished copper, zinc, iron, magnesium and aluminium in the boiled water and leave for 10 minutes. Note any change in the metal or water. Boil the water + metals for 5 minutes. Note any changes. If you see any bubbles on the metals, put the metal in a small basin of water and invert a test-tube of water over it. Leave for a few days to see whether larger quantities of the gas in the bubbles may be collected. Test any gas collected with litmus, limewater and a lighted splinter. The purpose of boiling the water before placing the metal into it is to remove any dissolved air that might react with the metal.

---

12.15.5 Reactions of non-metals with water
1. Shake small quantities of sulfur, carbon and iodine separately with water. Are there any indications of solution or chemical reaction? Filter each mixture. Test a little of the filtrate from the mixture containing iodine by pouring a little of it on to a piece of starch. Evaporate each filtrate to dryness and residue remains. The slight blue colour with iodine shows that iodine is slightly soluble in water. sulfur and carbon are insoluble in water. However, carbon heated to 1000^oC reacts with steam to produce the fuel water gas that can be added to coal gas.
carbon (s) + water (g) --> carbon monoxide (g) + hydrogen (g)
2. Pass some chlorine into water in a test-tube and shake the test -tube. Drop small pieces of red and blue litmus paper into the chorine and water. The blue litmus paper turns red then white as the chlorine and water bleaches it. Chlorine dissolves in water to produce hydrochloric acid and hypochlorous acid.

---

12.16.0 Carbonates
K, Na, Ca, Mg, Zn, and Pb carbonates are white. Fe carbonate is brown. Cu carbonate is blue-green. Only K and Na carbonates are soluble in water and are not decomposed by heat. Ammonium carbonate is a white powder.

---

12.16.1 Pass carbon dioxide through calcium carbonate suspension
Pass carbon dioxide through a suspension of calcium carbonate then boil the mixture. The calcium carbonate suspension disappears because the reaction forms soluble calcium hydrogen carbonate. Note that the reaction is reversible. Calcium hydrogen carbonate easily decomposes when heated.
CaCO₃(s) + CO₂(g) <--> CaHCO₃(aq)

---

12.16.1.1 Test for carbon dioxide, pass carbon dioxide through calcium hydroxide solution, limewater
See diagram 6.6.0: Test for carbon dioxide
Add water to cool freshly made calcium oxide (quicklime) in an evaporating basis to form calcium hydroxide. The reaction is exothermic and forms steam.
CaO(s) + H₂O(l) -->Ca(OH)₂(s)
Mix 1 mL of the solid calcium hydroxide with 10 mL of water. Test this with an indicator to show that it is a base. Leave the solution to stand. Decant the clear liquid that is limewater. Pass carbon dioxide through the clear liquid. The reaction forms a white precipitate of calcium carbonate. This reaction occurs when the mortar used in bricklaying sets hard to hold the bricks together. The water evaporates leaving the solid calcium carbonate.
Ca(OH)₂(aq) + CO₂(g) -->CaCO₃(s) + H₂O(l)
Continue to pass carbon dioxide through the solution. Soluble calcium hydrogen carbonate forms and the solution becomes clear again.
CaCO₃(s) + CO₂(g) + H₂O(l) -->Ca(HCO₃)₂(aq)
CO₂ (g) + H₂O(l) --> H₂CO₃ (aq) carbonic acid
H₂CO₃ (aq) + 2OH^-(aq) --> CO₃^2- (aq) + 4H₂O
Ca²⁺(aq) + CO₃^2- (aq) --> CaCO₃(s)
CaCO₃(s) + H₂CO₃ (aq) --> Ca²⁺(aq) + 2HCO₃^- (aq) bicarbonate ion

---

12.16.2 Prepare sodium hydrogen carbonate with sodium carbonate
Pass carbon dioxide through sodium carbonate solution to form sodium hydrogen carbonate sodium bicarbonate. If you heat dry sodium hydrogen carbonate the reverse reaction occurs.
Na₂CO₃(aq) + CO₂(g) + H₂O(l) -->2(NaHCO₃)(aq)

---

12.16.3 Heat different carbonates, carbonates of Cu, Mg, Na, Pb and Zn
Carbonates, except Na₂CO₃ and K₂CO₃, decompose on heating to form carbon dioxide and the oxide.
Heat powdered calcium carbonate with a strong burner. The calcium carbonate decomposes to form calcium oxide (quicklime) and carbon dioxide.
CaCO₃(s) -->CaO(s) + CO₂(g)
Heat different carbonates in a test-tube, e.g. carbonates of Cu, Mg, Na, Pb and Zn. Test the gases that form with: moist litmus paper, a drop of limewater on a glass rod, a lighted splint. The reaction forms carbon dioxide.
PbCO₃(s) -->PbO(s) + CO₂(g)

---

12.16.3.1 Heat ammonium carbonate (smelling salts)
Formerly this chemical was used to revive young ladies who had fainted by heating the container by hand to give off ammonia. Ammonium carbonate is a white powder fairly soluble in water forming a weak alkali.
Heat ammonium carbonate. Heat ammonium carbonate in a dry test-tube held sloping downwards. Observe the steam and condensed water on the cooler rim. Test for ammonia gas by smell and hold damp red litmus at the mouth of the test-tube. It turns blue. Ammonium carbonate decomposes to form three gases or vapours 1. steam 2. ammonia 3. carbon dioxide, leaving no residues. Smell the ammonia given off.

---

12.16.3.2 Reactions of ammonium carbonate with alkalis
Add ammonium carbonate to 2 cm depth of sodium carbonate (washing soda) solution or limewater solution. A vapour forms with an ammonia smell that turns red litmus blue.

---

12.16.3.3 Reactions of ammonium carbonate with acids
Add dilute hydrochloric acid or vinegar or citric acid t ammonium carbonate solution in a test-tube. Note the effervescence. test for carbon dioxide with limewater.

---

12.16.3.4 Ammonium carbonate solution precipitates metal carbonates
Add ammonium carbonate solution to solutions of copper (II) sulfate, iron (II) sulfate, magnesium sulfate, zinc sulfate and limewater. Note the colours of the precipitated metal carbonates.

---

12.16.4 Heat a hydrogen carbonate, sodium hydrogen carbonate (sodium bicarbonate)
This reaction is used in baking powder.
1. Heat a hydrogen carbonate in a test-tube. Test gases that form with: moist litmus paper, a drop of limewater on a glass rod, a lighted splint. The reaction forms carbon dioxide.
2. Heat sodium hydrogen carbonate (baking soda). Solid sodium hydrogen carbonate begins to decompose at 100^oC and is completely decomposed at 200^oC. The solution in water starts to decompose at room temperature.
2NaHCO₃(s) --> CO₂(g) + H₂0(g) + Na₂CO₃(s)

---

12.16.5 Put an egg in a bottle and tie a knot in a bone!
See also 12.3.9: Reactions of dilute acids and common carbonates
Cover a fresh egg with vinegar or dilute acid. Change the solutions each day for 2 days. The dilute acid dissolves most of the egg shell or bone The egg can now be squeezed through the narrow opening of a bottle. Show students the egg in the bottle and ask them to explain how the egg got into the bottle! Repeat the experiment using a thin chicken bone from a recently killed chicken, e.g. the wishbone (sternum). After some days tie a knot in the bone! Ask the students to explain how the chicken tied its bone into a knot!

---

12.16.6 Reactions of dilute acids with sodium hydrogen carbonate
See also 3.34.6: Soda-acid fire extinguisher | See also 12.3.10: Reactions of dilute acids and sodium hydrogen carbonate

---

12.16.6.01 Prepare an imitation volcano with baking soda
Prepare a heap of sand to represent the volcano and push a test-tube of long thin jar down into the heap of sand. Put baking soda or baking powder, food colouring, detergent and even glitter into the glass container. Carefully pour vinegar into the glass container.
BE CAREFUL! DO NOT LOOK DOWN INTO THE GLASS CONTAINER!

---

12.16.7 Solvay process, sodium carbonate
Soda ash is used to produce glass, detergents for metal refining, and for water purification.
In nature sodium carbonate decahydrate can be formed by the action of concentrated salt solutions on limestone.
2NaCl(aq) + CaCO₃(s) --> Na₂CO₃.10H₂O(s) + CaCl₂(aq)
In the laboratory sodium carbonate solution precipitates calcium carbonate from an aqueous solution of calcium chloride but in nature the reaction may be very slowly reversed in evaporating deposits because of the very high concentration of sodium chloride.
In the Solvay process, soda ash is produced by the reaction:
CaCO₃ + 2NaCl --> Na₂CO₃ + CaCl₂
The natural direction of this reaction is backwards but the reaction can be moved forward by various reactions including forcing carbon dioxide is forced under pressure into a concentrated cold brine solution saturated with ammonia adding ammonium ions and bicarbonate ions to the sodium and chloride ions already present.
NH₃(g) + CO₂(g) + NaCl(aq) + H₂O(l) -> NaHCO₃(s) + NH₄Cl (aq)
The least soluble combination of ions is sodium bicarbonate which precipitates. This anhydrous product is called light soda. The liquor is fed to the ammonia recovery plant where it is liberated with lime to leave calcium chloride. Lime kilns produce both lime and carbon dioxide for the process. Sodium bicarbonate is decomposed to sodium carbonate and the carbon dioxide released is recycled. The ammonia is regenerated and recycled by decomposing the ammonium chloride formed. Sodium carbonate solid is hydrated to monohydrate crystals for easier handling. Washing soda is produced by recrystallization, using the monohydrate from water to form the decahydrate. Washing soda will dehydrate spontaneously by efflorescence back to the monohydrate under dry conditions. Some of the waste concentrated calcium chloride liquor is used as a drilling mud for the oil industry and as an ice and snow melting salt in cold climates.

---

12.17.0 Oxides, acidic oxides and basic oxides
Acidic oxides are oxides of non-metals that react with bases to form salts + water, e.g. CO₂, NO₂, SO₂, SO₃. 2. Basic oxides are oxides of metals that react with acids to form salts + water. 3. Amphoteric oxides behave as acidic oxides and basic oxides, e.g. Al₂O₃, ZnO. 4. Carbon monoxide, dinitrogen oxide (nitrous oxide) nitrogen monoxide (nitric oxide) and water are neutral oxides. Hydrogen peroxide is an example of a higher oxide that forms oxygen when heated. 5. Mixed oxides contain more than one oxide, e.g. The anticorrosive pigment red lead oxide, dilead (II) lead (IV) oxide, Pb₃O₄(2PbO.PbO₂) The iron ore mineral magnetite, iron (II) iron (III) oxide, Fe₃O₄(FeO.Fe₂O₃). 6. Hydroxides refers to "hydrated oxides", OH.

---

12.17.1 Examine different oxides
All elements except the noble gases (inert gases) form oxides.
1. Different oxides, e.g. magnesium oxide, calcium oxide, aluminium oxide, carbon dioxide, sulfur dioxide, and nitrogen dioxide. 1. Describe the appearance. 2. Describe the odour. BE CAREFUL! DO NOT INHALE GASES DIRECTLY FROM THE TEST-TUBE! Fan the gas towards the nose with the hand and sniff cautiously. If no odour is detected, move closer and try again.
2. Add different oxides to water and shake. Note the relative solubility.
3. Test the acidity where solution has occurred.
4. Add drops of dilute sulfuric acid to each oxide. Note any reactions. Heat if no reaction occurs.
5. Add drops of sodium hydroxide solution to each oxide. Heat if no reaction occurs.
6. List the oxides in order of increasing acidic character.

---

12.17.1.1 Different oxides and the periodic table
All elements except inert gases form oxides. The oxides of metals in Group II were thought to be "like earth" and they form alkaline solutions, so the metals were called "alkaline earth" metals. Their oxides and hydroxides react with acids but not with alkalis. The oxide ion reacts with water to form the hydroxide (hydroxyl) ion.
O^2- + H₂O --> 2OH^-
With acids, the oxide ion reacts with the hydroxonium ion
O^2- + 2H₃O⁺ --> 3H₂O
The metallic properties become less to the right of the periodic table, e.g. aluminium oxide is insoluble in water, and reacts with both acids and alkalis to form water and salts so is called an amphoteric oxide. Farther to the right of the periodic table, the elements are non-metals.
They may react with water to form acid solutions.
Example 1. Carbon dioxide dissolves in water to form carbonic acid
CO₂(aq) + H₂O(l) --> H₂CO₃(aq)
Example 2. Phosphorus pentoxide (phosphorus (V) oxide) reacts violently with water to form phosphoric acid.
P₄O₁₀(s) --> H₂O(l) + 4H₃PO₄(aq)

---

12.17.2 Basic oxides, copper (II) oxide (copper oxide)
A basic oxide reacts with hydrogen ion to give water and a salt
CuO(s) + 2H⁺(aq) --> H₂O(l) + Cu²⁺(aq)
copper (II) oxide + hydrogen ion --> water + copper ion
CuO (s) + 2HCl(aq) --> CuCl₂ + H₂O(l)
copper oxide + hydrochloric acid --> copper (II) chloride + water
Basic oxides do not usually react with alkalis.
Put copper (II) oxide, calcium oxide, magnesium oxide and iron oxide in separate test-tubes. Add drops of alkali solution to each. Heat the mixture.

---

12.17.2.1 Copper (II) oxide and zinc
Weigh 2 g (0.025 mol) copper (II) oxide powder and 1.6 g (0.025 mol) zinc powder, zinc dust. Mix the powders to a uniform grey colour. Pour the mixture in the shape of a horizontal cylinder on a coffee tin lid. Heat one end of the mixture cylinder with a Bunsen burner until the mixture begins to glow. Stop heating and let the glow move along the cylinder of powder to the end leaving a white-grey mixture. Heat the coffee tin lid over a Bunsen burner to show that the white powder, zinc oxide, is yellow when hot and white when cool (because of change in the crystal structure of zinc oxide). Put the cooled residue in a beaker and add dilute hydrochloric acid to dissolve the zinc oxide and any remaining copper oxide and zinc, leaving red-brown copper. Heat the red-brown powder with concentrated nitric acid to give a blue solution of copper nitrate.
Repeat the experiment using course magnesium powder instead of zinc powder.

---

12.17.3 Acidic oxides, carbon dioxide
Acidic oxides dissolve in water to form an acid
CO₂(aq) + H₂O(l) <--> H₂CO₃(aq)
carbon dioxide + water <--> carbonic acid, that dissociates:
H₂CO₃^- <--> CO₃^2- + H⁺
carbonic acid <--> carbonate ion + hydrogen ion

---

12.17.3.1 Pass carbon dioxide through sodium hydroxide solution
Alkalis react with acidic oxides to form salt and water.
Pass carbon dioxide through sodium hydroxide solution. Note the reduction in the size of the bubbles, which shows that a reaction with carbon dioxide probably occurs. Stop the flow of carbon dioxide. Add drops of dilute hydrochloric acid. Test gases that form from the reaction with: moist litmus paper, a lighted splint, . The gas is carbon dioxide.
NaOH(aq) + CO₂(g) --> H₂O(l) + Na₂CO₃(aq)
Na₂CO₃(aq) + HCl(aq) -->NaCl(aq) + CO₂(g)

---

12.17.3.2 Pass carbon dioxide through barium hydroxide solution
Pass carbon dioxide through barium hydroxide solution. The reaction forms a white precipitate. Filter off the precipitate. Add dilute hydrochloric acid to the precipitate. Test the gas that forms with a lighted splint and moist litmus paper. The gas is carbon dioxide.
Ba(OH)₂(aq) + CO₂(g) --> BaCO₃(s) + H₂O(l)

---