topic12

Topic 12 Chemical reactions, pH, periodic table
Updated 2008-09-27
Please send comments to: J.Elfick@uq.edu.au
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Table of contents
12.1.0 Chemical equations and ionic equations
12.2.0 Types of chemical reactions
12.3.0 Properties of acids, ionization of carbonic acid
12.4.0 Hydrochloric acid
12.5.0 Nitric acid
12.5.01 Nitrous acid, ionization of nitrous acid
12.6.0 Sulfuric acid, ionization of hydrogen sulfate ion
12.9.0 Phosphoric acid, ionization reaction
12.20.0 Boric acid, ionization reaction
12.20.1 Prepare boric acid crystals

12.7.0 Bases, properties of bases
12.8.0 Acid-base neutralization
3.53 pH and acid-base indicators, acidity and alkalinity
3.53.1 Test different indicators
3.53.2 Prepare acid-base plant extract indicators

12.10.0 Salts
12.11.0 Tests for substances
12.12.0 Soaps and synthetic detergents (syndets)
12.13.0 Hardness in water
12.14.0 Activity series, activity 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

12.18.0 Periodic table
12.19.0 Group 17 of the periodic table, the halogens

12.1.0 Chemical equations and ionic equations, conservation of mass
12.1.1 Conservation of mass with effervescent tablets, health salts, sodium bicarbonate (baking soda)
12.1.2 Conservation of mass with burning steel wool
12.1.3 Conservation of mass in a cycle of copper reactions, the copper cycle experiment

12.2.0 Types of chemical reactions
12.2.1 Precipitation reactions, double decomposition
12.2.2 Synthesis reactions (combination reactions)
12.2.3 Decomposition reactions
12.2.4 Displacement reaction (substitution reactions)
12.2.5 Acid-base reactions
12.2.6 Redox reactions (oxidation-reduction reactions, electron transfer reactions) disproportionation
12.2.2.1 Synthesis reaction, Heat iron with sulfur
12.2.2.2 Synthesis reaction, Heat copper with iodine
12.2.2.3 Synthesis reaction, Heat iron with iodine
12.2.3.1 Decomposition reactions, Different decomposition reactions
12.2.7 Conditions for chemical reactions to occur, sulfuric acid with sodium chloride

12.3.0 Properties of acids
3.74 Displacement of hydrogen from acids by metals
12.3.1 Taste of acids, solid acids in the home
3.91.0 Rate of reaction
12.3.2 Reactions of dilute acids with metals, hydrochloric acid
12.3.2.1 Reactions of dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
12.3.3 Reactions of dilute acids with metals, sulfuric acid with iron
12.3.3.1 Reactions of dilute acids with metals, sulfuric acid with aluminium
12.3.4 Reactions of dilute acids with non-metals, carbon, sulfur
12.3.5 Reactions of dilute acids with basic oxides, copper (II) oxide
12.3.5.1 Reactions of dilute acids with basic oxides, zinc oxide
12.3.6 Reactions of dilute acids with hydroxides, magnesium hydroxide
12.3.7 Reactions of dilute acids with hydroxides, sodium hydroxide
12.3.7.1 Reactions of dilute acids with sodium hydroxide
12.3.7.2 Reactions of dilute hydrochloric acid with hydroxides
12.3.8 Reactions of dilute acids with acidic oxides, carbon dioxide, sulfur dioxide
12.3.9 Reactions of dilute acids with common carbonates
12.3.9.1 Reactions of dilute hydrochloric acid with calcium carbonate
12.3.9.2 Reactions of dilute hydrochloric acid with sodium carbonate
12.3.10 Reactions of dilute acids with sodium hydrogen carbonate
12.3.10.1 Reactions of dilute acids with calcium hydrogen carbonate
12.3.11 Reactions of dilute nitric acid with metals
12.3.11.1 Reactions of nitric acid with metals
12.3.12 Reactions of concentrated nitric acid with copper
12.3.13 Reactions of concentrated sulfuric acid with copper
12.3.14 Reactions of concentrated acids with non-metals, carbon, sulfur

12.4.0 Hydrochloric acid
2.42 Prepare hydrogen chloride, HCl
2.42.1 Tests for hydrogen chloride

12.5.0 Nitric acid
12.5.1 Prepare nitric acid, sulfuric acid with sodium nitrate

12.6.0 Sulfuric acid
12.6.1 Sulfuric acid acts as an oxidizing agent
12.6.2 Sulfuric acid dehydrating copper (II) sulfate crystals
12.6.3 Sulfuric acid dehydrating sucrose (cane sugar)
12.6.4 Sulfuric acid in water

12.9.0 Phosphoric acid
Ionization reaction
H₃PO₄ + H₂O <--> H₃O⁺ + H₂PO₄^-
H₂PO₄^- + H₂O <--> H₃O⁺ + HPO₄^2-
HPO₄^2-+ H₂O <--> H₃O⁺ + PO₄^3-

12.7.0 Bases, properties of bases
12.7.1 Feel of alkalis
12.7.2 Solubility of alkalis
12.7.3 Reactions of alkalis with metals, sodium hydroxide
12.7.3.1 Recycle an aluminium drink-can as potassium aluminium sulfate, alum.
12.7.4 Reactions of alkalis with salts, sodium hydroxide with copper salts
12.7.4.1 Reactions of alkalis with different salts, solubility of hydroxides
12.7.5 Reactions of alkalis with basic oxides, copper oxide
12.7.6 Reactions of alkalis with acidic oxides, carbon dioxide
12.7.7 Reactions of alkalis with amphoteric oxides and hydroxides
12.7.7.1 Action of sodium hydroxide with zinc chloride solution
12.7.8 Reactions of alkalis with sodium carbonate

12.8.0 Acid-base neutralization
3.78 Titration of acids with bases
3.82 Heat of a neutralization reaction
12.8.1 Ammonia with sulfuric acid
12.8.2 Sodium hydroxide with hydrochloric acid
12.8.3 Simple titration of acids with bases
12.8.4 Titrate dilute hydrochloric acid with sodium hydroxide solution, using a burette
12.8.4.1 Titrate dilute hydrochloric acid with sodium hydroxide solution, using a burette (second method)
12.8.5 Effect of carbon dioxide on acid-base titration
12.8.6 Heat of neutralization titration
12.8.7 Microscale titration, sodium hydroxide with dilute acids

pH and acid-base indicators
2.53 pH and acid-base indicators, acidity and alkalinity
2.53.1 Test different indicators
2.53.2 Prepare acid-base plant extract indicators
12.9.4 Rainbow reactions, t-Butyl chloride (2-chloro-2-methylpropane) with sodium hydroxide

12.1.0 Chemical equations and ionic equations, conservation of mass
Specific criteria can be used to classify chemical reactions. Redox reactions involve a transfer of electrons and a change in oxidation number. Precipitation reactions result in the appearance of a solid from reactants in aqueous solution. Acid-base reactions involve transfer of protons from donors to acceptors. Polymerization reactions produce large molecules with repeating units. Chemical reactions involve energy changes. All chemical reactions involve energy transformations. The spontaneous directions of chemical reactions are towards lower energy and greater randomness. Chemical reactions, but not nuclear reactions, obey the law of conservation of mass. The mass of the products is equal to the mass of the reactants. In a balanced chemical equation, the number of atoms of each element in the reactants is equal to the number of atoms in the product(s). Ionic equations show soluble ionic compounds in solution as separate ions and the sum of charges on the product side is equal to the sum of charges on the reactant side.
Chemical equation: HCl(aq) + NaOH(aq) --> NaCl(aq) + H₂O(l)
Ionic equation: H⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) --> H₂O(l) + Cl^-(aq) + Na⁺(aq)
The Cl^- and Na⁺ ions are called spectator ions because they are on both sides of the equation and do not react. Cancel the spectator ions to leave the net ionic equation:
H⁺ + OH^---> H₂O

12.1.1 Conservation of mass with effervescent tablets, health salts, sodium bicarbonate (baking soda)
See diagram 12.1.1: Measure conservation of matter | See also 13.7.7 Prepare carbon dioxide by heating hydrogen carbonates
When a chemical reaction occurs, matter is neither created nor destroyed. The mass of the reactants = the mass of the products. Effervescent tablets or "fruit salts" contain sodium hydrogen carbonate, and dry citric acid or tartaric acid
Put the tablet or fruit salts in water in a test-tube. Carbon dioxide forms as bubbles and any other substance in the tablet or fruit salts dissolves easily. Tests for carbon dioxide by holding a test-tube containing limewater at an angle near the mouth of the test-tube containing the effervescent tablet. The carbon dioxide given off by effervescence is heavier than air and will roll into the limewater test-tube where the limewater will turn milky because of the presence of carbon dioxide. Weigh an effervescent tablet or fruit salts and put into the bottom corner of a small plastic bag. Twist the bag above the corner and tie around the twist with thin string or wire. Weigh the plastic bag + string + effervescent substance. Pour a known amount of water into the open bag then tie string tightly to close the bag so that no liquid or gas can escape. The weight of the water + bag + string + effervescent substance is now known. Undo the string around the twisted part of the bag and untwist the bag. The acid and sodium hydrogen carbonate dissolves in the water and react to produce a salt and carbon dioxide. Weigh the bag and products of the reaction. The weight is the same.
2. Place a small amount of water in a plastic cup and place on the pan of the electronic scale.
Cover the top with a piece of paper or metal foil. Place two "Alka Seltzer" tablets on top of the cover. Record the initial mass. Tilt the cover so that the tablets drop into the water and immediately replace the cover so water droplets will not escape. Record the mass reading until it is constant. Repeat this experiment with the cup enclosed in a sealed container.
3. Place a small amount of water in a large 2 litre plastic drink bottle. Break two Alka Seltzer tablets in pieces that will fit in the bottle. Weigh the bottle. bottle cap, and Alka Seltzer tablets together. Drop the Alka Seltzer tablet pieces into the bottle and quickly replace the bottle cap tightly and place back on the scale. Record the mass reading until it is constant. The results differ from the experiment carried out in an open container. Remove the bottle from the scale, loosen the bottle cap, and measure again the weight of the bottle and its cap.

12.1.2 Conservation of mass with burning steel wool
Roll some steel wool into a ball about the size of an egg and weigh it. Hold it with tongs over a sheet of paper. Heat the steel wool until red-hot; remove the flame and blow gently on the red-hot steel until it stops burning. When cold, weigh the steel wool and any fragments fallen onto the sheet of paper on the balance. When iron burns the product formed, iron oxide, is heavier than the iron.

12.1.3 Conservation of mass in a cycle of copper reactions, the copper cycle experiment
See diagram: 1.13a Simple fume hood
Step 1. Convert copper metal to copper nitrate
1. Weigh 1.000 g of copper wire. It must be clean, bright and shiny. Twist the wire into a flat spiral and put it in a beaker in a fume hood. Slowly add 4 .0 mL of 16 M nitric acid. Be careful! Note the brown fumes of nitric oxide (NO), nitrogen dioxide (NO₂), and dinitrogen tetroxide (N₂O₄), i.e. NOx. When all the copper is dissolved add 100 mL of deionized water.
Cu(s) + 8HNO₃(aq) --> 3Cu(NO₃)₂(aq) + 4H₂O(l) +2NO(g) [dilute nitric acid]
Cu(s) + 4HNO₃(aq) --> Cu(NO₃)₂(aq) + 2H₂O(l) + 2NO₂(g) [concentrated nitric acid]
N₂O₄ <--> 2NO₂ [in equilibrium]
Step 2 Convert copper nitrate to copper hydroxide
Add 30.0 mL of 3.0 M sodium hydroxide to the solution while stirring. If red litmus paper does not turn blue in the solution, add more sodium hydroxide. Note the precipitate of copper hydroxide, an ionic solid.
Cu(NO₃)₂(aq) + 2NaOH(aq) --> Cu(OH)₂(s) + 2NaNO₃(aq)
Step 3 Convert copper hydroxide to copper oxide
Heat the solution on a hot plate while continually stirring to prevent bumping caused by steam bubbles. Note the precipitate changing to a black solid, copper oxide, CuO. Carefully decant the liquid, add deionized water and decant again.
Heat the precipitate until it becomes a firm mass.
Cu(OH)₂(s) + heat --> CuO(s) + H₂O(l)
Step 4 Convert copper oxide to copper sulfate
Reaction 4: Converting copper oxide to copper sulfate
Add 15 mL of 6.0 M sulfuric acid to the copper oxide while swirling, not stirring, the copper oxide to help it dissolve.
CuO(s) + H₂SO4(aq) --> CuSO4(aq) + H₂O(l)
Step 5 Convert copper sulfate to copper metal
CuSO4(aq) + Zn(s) --> ZnSO4(aq) + Cu(s)
In the fume hood add 2.0 g of zinc metal and keep stirring until the solution becomes colourless. The zinc is oxidized as it reduces the copper.
Add one drop of the solution to 1 mL of concentrated ammonia solution in a test-tube. If the ammonia turns deep blue, some unreduced copper is till in the solution so the reaction is not finished. When all the copper is reduced, decant the liquid and add 20.0 mL of 6M hydrochloric acid to dissolve excess zinc. Note the bubbles of hydrogen until the reaction is complete. Cool the beaker and observe the metallic copper settling on the bottom. Carefully decant the solution and wash the copper metal with deionized water and decant again. Transfer the copper to an evaporating dish using a small amount of deionized water. Decant excess water evaporating dish then wash the precipitate with methanol. Decant the liquid the gently heat the precipitate on a hot plate. If you heat the copper precipitate too strongly it will oxidize to copper oxide. Transfer the dried copper metal to a preweighed beaker and calculate the mass of recovered copper.
% recovery = mass of recovered copper / initial mass of copper X 100.

12.2.0 Types of chemical reactions
In a chemical reaction, a chemical change occurs where elements or compounds (reactants) form new substances (products)
The main types of chemical reactions are precipitation reactions, synthesis reactions, decomposition reactions, displacement reactions, acid-base reactions and redox reactions.

12.2.1 Precipitation reactions, double decomposition
When sodium chloride solution is added to silver nitrate solution and insoluble solution of of silver chloride forms.
AgNO₃(aq) + NaCl(aq) --> AgCl(s) + NaNO₃(aq)
silver nitrate + sodium chloride --> silver chloride + sodium nitrate
The silver chloride precipitate cabn be separated from the sodium nitrate solution.
Ionic equations
The ionic equation that shows all the substances
Ag⁺(aq) + NO₃^-(aq) + Na⁺(aq) + Cl^-(aq) --> AgCl(s) + Na⁺(aq) + NO₃^-(aq)
The net ionic equation that does not contain the "spectator ions" that appear on both sides of the equation but do not form a precipitate
Ag⁺(aq) + Cl^-(aq) --> AgCl(s)
The above reaction may also be called a double decomposition reaction (metathesis) because the positive and negative parts of two compounds swap partners, i.e. exchange radicals. In general: AB + CD --> AD + CB

12.2.2 Synthesis reactions (combination reactions)
Elements or simple molecules combine to form a new compound, e.g. reaction of zinc with sulfur
Zn(s) + S(s) --> ZnS (s)
zinc + sulfur --> zinc sulfide

12.2.3 Decomposition reactions
A compound breaks down into simpler compounds or into elements, usually caused by heat - the opposite of a synthesis reaction. All compounds decompose on heating to a high enough temperature to form elements or simple molecules. Compounds of metals higher in the activity series are harder to decompose by heating than compounds lower in the activity series, so copper compounds break up more readily on heating than sodium compounds.
Decomposition of calcium carbonate
CaCO3(s) --> CaO(s) + CO₂(g)
calcium carbonate --> calcium oxide + carbon dioxide

12.2.4 Displacement reactions (substitution reactions)
The reactants are an element and a compound. The element replaces part of the compound with the same valence and same sign, e.g. displacement of Cu2+ by zinc
Zn(s) + CuSO₄(aq) --> ZnSO₄(aq) + Cu(s)
zinc + copper sulfate --> zinc sulfate + copper
The copper precipitates as the element and the zinc metal goes into solution as zinc ions.
See also 3.71 Reactions of ions in solutions | See also12.14.1: Zinc displaces lead from lead nitrate solution

12.2.5 Acid-base reactions
An acid dissociates in water to produce positive hydrogen ions, H⁺, that is solvated to produce hydronium ions (hydroxonium ions), H3O+, by transferring a proton (H⁺) to a water molecule.
HCl(g) + H₂O(l) --> H₃O⁺(aq) + Cl^-(aq)
A base dissociates in water to produce negative hydrolide ions, OH^-.
NaOH --> Na⁺(aq) + OH^-(aq)
Acids reactwith bases to from salts and water. the producta are neither acidic nor basic so this reaction is called neutrlaization.
HCl + NaOH --> NaCl + H₂O
hydrochloric acid + sodium hydroxide --> sodium chloride + water
The ionic equation that shows all the substances
H₃O⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) --> 2H₂O + Na⁺(aq) + Cl^-(aq)
The net ionic equation
H₃O⁺(aq) + OH^-(aq) --> 2H₂O

12.2.6 Redox reactions (oxidation-reduction reactions, electron transfer reactions)
Electrons move from one atom to another. Oxidation is loss of electrons. Reduction is gain of electrons. The same number of electrons are gained in the reduction as are lost in the oxidation.
In the reaction of dilute hydrochloric acid on magnesium ribbon, each magnesium atom loses two electrons to two hydrogen atoms.
Mg(s) + HCl(aq) --> MgCl₂(aq) + H₂(g)
Mg(s) + 2H₃O⁺(aq) + 2Cl^-(aq) --> H₂(g) + Mg²⁺(aq) + 2Cl^-(aq) + 2H₂O
In reactions where no ions are formed, use the idea of oxidation number (oxidation state) to show the "apparent charge" on an atom. In the reaction between gases:
2SO₂(g) + O₂(g) --> 2SO₃(g)
Give the oxygen atom a net charge of -2, but give O₂ a net charge of zero because the oxygen atom is in the elemental form. Then the sulfur atom in SO₂ has an oxidation number +4 and the sulfur atom in SO₃ has an oxidation number +6. The sulfur atoms have been oxidized because the oxidation number has increased and the oxygen atoms in O₂ have been reduced because the oxidation number has decreased.
Similarly, in the the following equation:
NH₃ + CuO --> Cu + H₂O + N₂
The oxidation number of hydrogen atom in NH₃ is +1 and in H₂ is zero, because the hydrogen atom is in the elemental form. The oxidation number of the nitrogen atom has increased from -3 in NH₃ to 0 in N₂, because N in N₂ is in the elemental form. The oxidation number of the copper has decreased from +2 in CuO to zero in Cu, because the Cu atom is in elemental form. The nitrogen atom has been oxidized and the copper atom has been reduced.
However, when chlorine has dissolves in water a disproportionation occurs because the the chlorine becomes both oxidized, when HClO is formed, and reduced, when HCl is formed.
Cl₂ (g) + H₂O(l) <--> HClO(aq) + HCl(aq)

12.2.2.1 Heat iron with sulfur
See diagram 12.2.1
S₈(s) + 8Fe(s) --> 8FeS(s)
BE CAREFUL! THE REACTION OF IRON (II) SULFIDE WITH HYDROCHLORIC ACID WILL FORM THE POISONOUS GAS, HYDROGEN SULFIDE, WITH AN ODOUR OF ROTTEN EGGS.
1. Mix uniformly reduced iron powder and powdered sulfur in a weight ratio of seven to four. Carve the word "FeS" on a red coloured brick with a knife. Spread the iron sulfur mixture throughout the word groove and press the powdered mixture solid. Heat one tip of a glass rod until red-hot with an alcohol burner and then immediately dig the hot tip into the mixture at one end of the word groove. A chemical reaction is starts immediately. The reaction continues violently to release a large amount of heat and meanwhile to develop rapidly a red glow, which looks like a small "fiery dragon". The heat lost by the reaction is more than the heat needed to start the reaction. The reaction produces a new black solid substance, iron (II) sulfide, that has different properties from the two reactants, iron and sulfur. Compare iron powder, powdered sulfur and iron (II) sulfide. Note their appearance. Test them respectively with a magnet. Add in drops hydrochloric acid solution to them respectively.
2. Mix equal amounts of iron filings and powdered sulfur. Heat the mixture in a crucible or a small tin with sand in the bottom. The sand prevents the bottom of the tin from melting by spreading the heat. Heat the mixture strongly until you see a red glow spreading through the mass. The heat lost by the chemical reaction is more than the heat needed to start the reaction. The reaction forms a new substance iron (II) sulfide that has different properties from the two elements used to make it. Compare iron filings, powdered sulfur, and iron (II) sulfide. Note their appearance. Test with a magnet. Add drops of hydrochloric acid.
3. Repeat the experiment with sulfur in the bottom of a test-tube and a strip of zinc half way up the test-tube.

12.2.2.2 Heat copper with iodine
Mix equal amounts and heat gently in a test-tube.
BE CAREFUL! THE IODINE MAY STAIN THE SKIN. REMOVE STAINS WITH SODIUM THIOSULFATE SOLUTION.
Stop heating when you hear a hissing noise. Heat again to make sure all the copper reacts with the iodine Excess iodine sublimes and solidifies up the tube. Let the tube cool then scrape out the product of the reaction. Compare the crushed product with the reactants copper and iodine. The reaction forms a new substance, copper iodide.

12.2.2.3 Heat iron with iodine
Put iodine crystals in a test-tube and then push in a plug of steel wool. Clamp the test-tube at an angle and heat the steel wool with a Bunsen burner. The steel wool glows red and the iodine evaporates. A new substance forms.

12.2.3.1 Different decomposition reactions
1. Place small quantities of zinc oxide and copper oxide in separate small test-tubes and heat carefully. Identify any gases produced. Note any new substance formed in the test-tube.
2. Repeat using sodium or potassium nitrate, lead nitrate, copper nitrate. If a brown gas is produced, it is nitrogen dioxide.
3. Repeat using sodium carbonate, magnesium carbonate, calcium carbonate, lead carbonate.
4. Repeat using sodium sulfate, magnesium sulfate, zinc sulfate and copper (II) sulfate.
potassium nitrate(s) --> potassium nitrite(s) + oxygen(g)
magnesium carbonate(s) --> magnesium oxide(s) + carbon dioxide(g)
copper nitrate(s) --> copper oxide(s) + nitrogen dioxide(g) + oxygen(g)
zinc sulfate(s) --> zinc oxide(s) + sulfur trioxide(g)
copper (II) sulfate(s) --> copper oxide(s) + sulfur trioxide(g)

12.2.7 Conditions for chemical reactions to occur, sulfuric acid with sodium chloride
See also 3.71.1: Solubility table and solubility rules | See also 3.71: Reactions of ions in solutions
1. Reactions of concentrated sulfuric acid with solid sodium chloride BE CAREFUL!
The reactions contain no water. Two reactions occur and both go to completion if heated. The reactions occur because hydrogen chloride has a lower boiling point than sulfuric acid.
2NaCl(s) + H₂SO₄(l) --> Na₂SO_4(aq) + 2HCl(g)
NaCl(s) + H₂SO₄(l) --> NaHSO₄(aq) + HCl(g)
2. Reactions of dilute sulfuric acid with solid sodium chloride: The reaction does not go to completion because the hydrochloric acid dissolves in the water. One product of the reaction is a slightly ionized substance, e.g. water. In neutralization reactions HOH is forming, so the reaction can almost go to completion. One product of the reaction is a precipitate. An insoluble substance leaves the solution. The solubility rules state that all chlorides are soluble except Ag⁺, Hg²⁺ and Pb²⁺(slightly). Predict whether the following reaction occurs. The reaction occurs because insoluble silver chloride precipitates.
NaCl(aq) + AgNO₃(aq) --> NaNO₃(aq) + AgCl(s)
NaOH(aq) + HCl(l) --> NaCl(aq) + H₂O(l)

12.3.0 Properties of acids
Ionization reaction of carbonic acid
CO₂ + H₂O <--> H₃O⁺ + HCO₃^-, K1 = 4.4 X 10^-7
HCO₃^- + H₂O <--> H₃O⁺ + CO₃^2-, Ka = 4.7 X 10^-11
Acids are good electrolytes, react with active metals, turn blue litmus red, and have a sour taste. Dilute acids contain hydrogen ions in aqueous solution. You can represent the hydrogen ion, which is really a proton, in different ways to show how it is related to the water molecules in the solution. You can show it as the hydrated hydrogen ion, [proton, H⁺(aq)] or as the hydronium ion [oxonium ion, H₃O⁺(aq)] but, for convenience, use H⁺(aq). Concentrated sulfuric acid exists mainly as H₂SO₄molecules. Hydrochloric acid and nitric acid dissociate into ions even in concentrated solution. Weak acids, e.g. ethanoic acid (acetic acid, CH₃COOH) carbonic acid and sulfurous acid dissociate very little in aqueous solution, but their salts, e.g. potassium acetate (CH₃COOK) are completely dissociated into ions. You can think of an acid in a reaction as a proton donor. You can think of an acid in a reaction as an electron pair acceptor
2H⁺ + 2e^- --> H₂
An acid, e.g. HX, dissociates to form hydrogen ions (a proton)
HX <--> H⁺ + X^-
The hydrogen ion is solvated to form hydronium ion (hydroxonium ion)
HX + H₂O <--> H₃O⁺ + X^-12.3.1 Taste of acids, solid acids in the home
See also 19.1.0: Solid acids, solubility
BE CAREFUL! NEVER TASTE ACIDS IN THE LABORATORY!
Lemon juice (citric acid) and sour milk (lactic acid) are two examples of common foods that taste sour because they contain acids.
Moisten your finger with a very dilute solution of hydrochloric acid. Rub your fingers together and then lick them. Repeat the procedure with very dilute solutions of acetic acid, citric acid and tartaric acid. Do not taste any other acids because they may damage living tissues.

12.3.2 Reactions of dilute acids with metals, hydrochloric acid
Reactions of acids with metals are exothermic. The higher the metal is in the activity series the greater the heat liberated. Reactions of dilute hydrochloric acid with zinc:
Zn(s) + 2HCl(l) --> H₂(g) + ZnCl₂(aq)
The order of activity of metals with acids is similar to the order of activity with water.
Evolution of hydrogen occurs

Metal	2M Hydrochloric acid	2M Sulfuric acid
Magnesium	Very rapid	Rapid
Aluminium	Slight	None
Zinc	Moderate	Slight
Iron	Very slight	Very slight
Tin	None	None
Lead	None	None
Copper	None	None

1. Use different cleaned metals, e.g. calcium pieces, iron nail, lead sinker, magnesium ribbon, copper wire, aluminium sheet and zinc granules. Rub them with emery paper to make surfaces clean of oxides.
Put each metal into a separate test-tube. Add 10 mL of 2 M hydrochloric acid to test-tubes. Properties of any gas liberated and name it - test with moist pieces of red and of blue litmus paper, with a drop of limewater hanging from a glass rod and with a lighted splint. Compare the rate at which hydrogen evolves by noting the rate and size of the hydrogen bubbles from the reaction. Describe the rate of reaction as nil, very slow, slow, moderately fast, very fast, and whether energy, in the form of heat, is produced (exothermic) or absorbed (endothermic). List the acids in order of their activity towards metals and state whether the same gas was liberated during each reaction and whether a salt may be isolated when the acids react with a metal. Make up an activity series by listing the elements in approximate order of their activity with respect to acids, from the most active to the least active. Compare the results with the table of activity series of some metals. The order of activity of the metals used, from the most active to the least active, is: magnesium, aluminium, zinc, iron with lead and copper displaying no noticeable reaction. When reaction did occur, the gas liberated was hydrogen. The reactions of these acids with metals are exothermic. The order of activity of the acids is that dilute hydrochloric and dilute sulfuric acids are about equal in activity but that they are more reactive than acetic acid. The order of activity of the metals with respect to acids is similar to that with respect to water. Magnesium ribbon forms most rapid bubbles of hydrogen then zinc then iron. Tin forms few bubbles of hydrogen. Copper forms no bubbles of hydrogen Lead forms some lead chloride precipitate on the surface of the lead. Aluminium develops a layer of aluminium oxide that obstructs further chemical reactions.
2. Note the properties of any gas that forms. Test the gas with moist litmus paper a lighted splint and a hanging drop of limewater on a glass rod.
3. Feel the test-tube to note whether heat energy is released or absorbed. The reactions of these acids with metals are exothermic. 4. List the elements in approximate order of their activity with respect to hydrochloric acid from the most active to the least active. The order of activity is: magnesium, aluminium, zinc, iron, lead - no noticeable reaction, copper - no noticeable reaction.

12.3.2.1 Reactions of dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
Dilute hydrochloric and dilute sulfuric acids are about equal in activity, but that they are more reactive than ethanoic acid (acetic acid). Note the slower production of hydrogen with the weak acetic acid.
The reaction with sulfuric acid forms insoluble sulfates on the surface of calcium and lead that obstructs or stops reactions. List the acids in order of their activity on metals.

12.3.3 Reactions of dilute acids with metals, sulfuric acid with iron
Add dilute sulfuric acid to steel wool in a test-tube. Test the gas that forms with a lighted taper. BE CAREFUL! THE GAS IS HYDROGEN GAS!
Heat the mixture in a beaker of hot water until all the steel wool has dissolved. Add more acid when necessary. Filter the hot solution then leave it to cool. Crystals form on cooling. If no crystals form, add alcohol because the salt is less soluble in it. Dry the green crystals of iron (II) sulfate-7-water between absorbent paper.
Fe(s) + H₂SO₄(aq) --> H₂(g) + FeSO₄(aq)

12.3.3.1 Reactions of dilute acids with metals, sulfuric acid with aluminium
Heat dilute sulfuric acid with pieces of aluminium foil in a test-tube. Some effervescence occurs but sometimes not enough to test for hydrogen with a lighted taper. After heating for 5 minutes, decant the solution that contains aluminium sulfate into another test-tube and add ammonia solution. A white jelly-like precipitate of aluminium hydroxide forms.

12.3.4 Reactions of dilute acids with non-metals, carbon, sulfur
Add a piece carbon and sulfur to dilute hydrochloric acid, dilute sulfuric acid and dilute ethanoic acid (acetic acid) in separate test-tubes. Heat the test-tubes. No reaction occurs. Non-metals do not react with dilute acids.

12.3.5 Reactions of dilute acids with basic oxides, copper (II) oxide
1. Heated dilute acids react with metal oxides to form a salt and water: Pour dilute sulfuric acid into a Pyrex test-tube and heat in a beaker of boiling water until the sulfuric acid is nearly boiling. BE CAREFUL!
Add pieces of copper (II) oxide one by one while stirring until some remains unreacted with the acid. Filter the undissolved copper oxide form the hot solution. Leave the filtrate in a watch glass to cool and form crystals. Blue crystals of copper (II) sulfate-5-water form with water. Remove the crystals and dry them by pressing between absorbent paper.
H₂SO₄(aq) + CuO(s) --> CuSO₄(aq) + H₂O(l)
acid + basic oxide --- salt + water
2. Repeat the experiment with dilute nitric acid.
2HNO₃(aq) + CuO(s) --> Cu(NO₃)₂(aq) + H₂O(l)

12.3.5.1 Reactions of dilute acids with basic oxides, zinc oxide
Oxides of Sn, Al, Zn, Pb, and Sb are amphoteric. Amphoteric oxides react with bases to form a salt + water. Amphoteric oxides react with acids to form a salt + water.
Add dilute hydrochloric acid to zinc oxide.
2HCl(aq) + ZnO(s) --> ZnCl₂(aq) + H₂O(l)
2NaOH(aq) + ZnO(s) --> Na₂ZnO₂(aq) + H₂O(l)

12.3.6 Reactions of dilute acids with hydroxides, magnesium hydroxide
Basic hydroxides are insoluble in water and react with acids to form a salt and water. Many metallic hydroxides react with acids to form a salt and water.
Add magnesium hydroxide to dilute sulfuric acid until the reaction stops. Filter the mixture. Test the filtrate with litmus paper. Evaporate the filtrate to dryness so that crystals form.
Mg(OH)₂(s) + H₂SO₄(aq) --> MgSO₄(aq) + H₂O(l)

12.3.7 Reactions of dilute acids with hydroxides, sodium hydroxide
Acids react with (neutralize) alkalis to form a salt and water.
Pour 5 mL of dilute sodium hydroxide solution into a watch glass. Test with litmus paper. Red litmus turns blue. Add dilute hydrochloric acid drop by drop. Stir as each drop is added. Test the mixture with the litmus paper until the litmus paper is neither red nor blue, but between these colours. Evaporate the solution to dryness by heating the watch glass over a beaker of boiling water. Crystals of sodium chloride (common salt) form.

12.3.7.1 Reactions of dilute acids with sodium hydroxide
Repeat the previous experiment with: dilute sulfuric acid, dilute nitric acid, ethanoic acid (acetic acid).

12.3.7.2 Reactions of dilute hydrochloric acid with hydroxides
[NH₃(aq) is used because while "NH₄⁺" ions and "OH^-" ions can be detected, "NH₄OH" cannot be detected, so ammonia solution is shown as "NH₃(aq) + H₂O(l)"]
Repeat the experiment with dilute solutions of: potassium hydroxide, calcium hydroxide, aqueous ammonia solution.
acid + (base) alkali --> salt + water
HCl(aq) + NaOH(aq) --> NaCl(aq) + H₂O(l)
HNO₃(aq) + NaOH(aq) --> NaNO₃(aq) + H₂O(l)
HCl(aq) + KOH(aq) --> KCl(aq) + H₂O(l)
HCl(aq) + NH₃(aq) + H₂O(l) --> NH₄Cl(aq) + H₂O(l)

12.3.8 Reactions of dilute acids with acidic oxides, carbon dioxide, sulfur dioxide
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Note any reaction for five minutes then evaporate to dryness. In each case, no reaction occurs. In each experiment there is no precipitate. If you evaporate a sample of a remaining solution to dryness in a fume cupboard, no residue remains. Pass carbon dioxide through hydrochloric acid or ethanoic acid (acetic acid) solution. Pass sulfur dioxide through hydrochloric acid or ethanoic acid (acetic acid) solution.

12.3.9 Reactions of dilute acids with common carbonates
Dilute acids react with metal carbonates to form a salt, carbon dioxide and water. Geologists use this reaction to identify calcium carbonate in rock. Drops of hydrochloric acid cause bubbles to form.
1. Make a chemical egg peeler. Put an egg in vinegar (contains acetic acid, ethanoic acid). Note the bubbles forming on the outside of the egg. Leave overnight then, the next day, pick up the egg with your fingers. The egg has become soft. Leave to stand for a few days and the egg shell disappears completely. You can now see through the raw egg.
The acetic acid in the vinegar + calcium carbonate in the egg shell --> calcium acetate in solution + bubbles of carbon dioxide + water
2. Add 5 mL vinegar or dilute hydrochloric acid or dilute sulfuric or dilute nitric acid to pea size amounts of finely divided common carbonates: sodium hydrogen carbonate, sodium carbonate, calcium carbonate, magnesium carbonate, nickel carbonate, limestone, lime, oyster shells, egg shell, snail shell, coral. Continue to add the solid until no further reaction occurs. Filter and evaporate the filtrate to dryness. Note any visible changes. Test any gas liberated by inserting in the mouth of the tube first damp pieces of red and of blue litmus paper then a drop of limewater hanging on the tip of a glass rod and finally a burning splinter. In each case the gas is carbon dioxide.

12.3.9.1 Reactions of dilute hydrochloric acid with calcium carbonate
See diagram 3.34.1: Limewater test for carbon dioxide
Put calcium carbonate in a test-tube. Add 2 mL 1.0 M hydrochloric acid. Tilt the test-tube so that its mouth is touching a second test-tube containing 5 mL of limewater. The surface of the limewater turns milky. Shake the test-tube containing the limewater. The milky colour on the surface disappears.
CaCO₃(s) + 2HCl(aq) --> CO₂(g) + CaCl₂(aq) + H₂O(l)
carbonate + acid --> carbon dioxide + salt + water

12.3.9.2 Reactions of dilute hydrochloric acid with sodium carbonate
1. Put sodium carbonate in a test-tube and add drops of dilute hydrochloric acid.
Test any gases formed from the reaction with moist litmus paper, a lighted splint, and a drop of limewater on a glass rod. The reaction forms carbon dioxide. Add more carbonate until no more reaction occurs. Filter and evaporate the filtrate to dryness. Repeat the experiment with dilute nitric acid. Repeat the experiment with magnesium carbonate.
Na₂CO₃(s) + 2HCl(aq) --> 2NaCl(aq) + H₂O(l) + CO₂(g)
Na₂CO₃(s) + 2HNO₃(aq) --> 2NaNO₃(aq) + H₂O(l) + CO₂(g)
2. Shake different solid acids in separate test-tubes half filled with water. Divide the solutions in the test-tubes into three different test-tubes. Test-tube A: Add small pieces of red and of blue litmus paper. Test-tube B: Add three drops of methyl orange solution. Test-tube C: Add three drops of phenolphthalein solution. Observe any changes in the solutions. Add solid sodium carbonate to each acid solution. Observe any changes in the solutions. Pass some gas given off into a test-tube containing limewater. Shake the test-tube for thorough mixing. Note the milkiness of the solution because carbon dioxide was produced when the acids reacted with sodium carbonate.

12.3.10 Reactions of dilute acids with sodium hydrogen carbonate
The only stable hydrogen carbonates are KHCO₃and NaHCO₃. Sodium hydrogen carbonate, bicarbonate of soda, is used in baking soda, baking powder, self-raising flour, effervescent fruit salts and soda-acid fire extinguishers and treatment for acid burns. Some people swallow sodium hydrogen carbonate to counteract excess acid in the stomach but using magnesium oxide or magnesium hydroxide that does not react with acids to produce carbon dioxide is better.
1. Add sodium hydrogen carbonate, or other hydrogen carbonates, to acids to form carbon dioxide, water and a salt.
NaHCO₃ + HCl --> CO₂+ H₂O + NaCl
hydrogen carbonate + acid --> carbon dioxide + water + salt
2. Mix vinegar with bicarbonate of soda in a glass jar. Drop some naphthalene mothballs into the solution. The carbon dioxide formed by the reaction of the vinegar (acetic acid) with the sodium hydrogen carbonate forms bubbles of carbon dioxide on the mothballs in the bottom of the jars. The mothballs rise to the surface, lose the bubbles and sink again.

12.3.10.1 Reactions of dilute acids with calcium hydrogen carbonate
Put powdered calcium carbonate into a test-tube containing about 10 mL of water. Bubble carbon dioxide through the suspension until no further change takes place. Soluble calcium hydrogen carbonate forms. Boil the mixture for 10 minutes. Add acids to form carbon dioxide, water and a salt.

12.3.11 Reactions of dilute nitric acid with copper
Very dilute nitric acid may react with very active metals, e.g. magnesium to form hydrogen. When nitric acid reacts with most metals, it oxidizes the hydrogen to water.
Add drops of dilute nitric acid to copper. Nitrogen monoxide forms which immediately reacts with oxygen in the air to form nitrogen dioxide.
3Cu(s) + 8HNO₃(aq) --> 3Cu(NO₃)₂(aq) + 4H₂O(l) + 2NO(g)
2NO(g) + O₂(g) --> 2NO₂(g)

12.3.11.1 Reactions of nitric acid with metals
Add slowly small pieces of copper, magnesium and zinc to small amounts of dilute nitric acid in separate test-tubes. If no change is taking place, gently heat the mixture. Now repeat the procedure 1. with concentrated nitric acid 2. with concentrated sulfuric acid and 3. with concentrated hydrochloric acid. The reactions of metals with nitric acid and concentrated sulfuric acid are different from reactions of metals with hydrochloric acid, dilute sulfuric acid and dilute acetic acid. Although copper does not react with dilute acids or with concentrated hydrochloric acid, it does react with both dilute and concentrated nitric acids and with hot concentrated sulfuric acid but does not produce hydrogen in reaction with them. The residual mixtures contain solutions of salts but writing equations for the reactions is difficult because more than one reaction can occur simultaneously between copper or magnesium or zinc and nitric acid. For example when zinc reacts with nitric acid the reaction may produce nitrogen dioxide, nitric oxide, nitrous oxide, zinc nitrate and ammonium nitrate.

12.3.12 Reactions of concentrated nitric acid with copper
Nitric acid reacts with metals above platinum in the activity series, but does not form hydrogen.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Pour drops of concentrated nitric acid on pieces of copper in a test-tube. Put a stopper on the test-tube immediately because brown nitrogen dioxide gas forms. The nitric acid acts as an oxidizing agent and is reduced to nitrogen dioxide and water. The reaction is exothermic.
Cu(s) + 4HNO₃(aq) --> Cu(NO₃)₂(aq) + 2H₂O(l) + 2NO₂(g)

12.3.13 Reactions of concentrated sulfuric acid with copper
Concentrated acids should be handled only by experienced science teachers. Concentrated sulfuric acid reacts with metals above platinum in the activity series, but do not form hydrogen.
BE CAREFUL! DO THIS EXPERIMENT IN A FUME CUPBOARD.
Add hot concentrated sulfuric acid to a piece of copper foil. Brown nitrogen dioxide gas forms. The sulfuric acid acts as an oxidizing agent.
Cu(s) + 2H₂SO₄(aq) --> CuSO₄(aq) + 2H₂O(l) + SO₂(g)

12.3.14 Reactions of concentrated acids with non-metals, carbon, sulfur
DO NOT DEMONSTRATE THIS EXPERIMENT!
Hot sulfuric acid and nitric acid can react as oxidizing agents with carbon and sulfur. Carbon is oxidized to carbon dioxide and nitric acid is reduced to nitrogen dioxide and water.
C(s) + 4HNO₃(aq) --> CO₂(g) + 4NO₂(g) + 2H₂O(l)

12.3.15 Reactions of acids and with salts
Add small quantities of sodium chloride, sodium nitrate, sodium acetate, sodium sulfite and iron sulfide to about 5 mL of dilute hydrochloric acid in separate test-tubes. Observe what happens when the mixtures are cold and when they are warmed. Repeat the procedure using dilute sulfuric acid and then concentrated sulfuric acid. Dilute acids do not react with chlorides, nitrates, sulfates, or acetates unless the metal ions in the salt can form an insoluble salt with the ions in the acid. Acids react with sulfites to produce sulfur dioxide, water and a salt. Acids react with sulfides to produce hydrogen sulfide (rotten egg gas) and a salt. Concentrated sulfuric acid reacts with chlorides to produce hydrogen chloride and a sulfate. Concentrated sulfuric acid reacts with nitrates to produce nitric acid and a sulfate. Concentrated sulfuric acid reacts with acetates to produce acetic acid and a sulfate.

12.4.0 Hydrochloric acid
Hydrochloric acid is an aqueous solution of hydrogen chloride gas. Hydrochloric acid dissolves most metals to form chlorides and hydrogen gas. Hydrochloric acid is available as: 1. 5.0 M, 4.0 M, 2.0 M, 1.0 M and 0.5 M volumetric solutions 2. minimum assay 36% solution density 1.17 g cm^-3 at 20^oC 3. 36% "ANALAR" solution (d) commercial solution called muriatic acid for use in the building trades.

12.5.0 Nitric acid
See also 12.3.13: Reactions of concentrated sulfuric acid with copper | See also 12.3.5: Reactions of dilute acids with basic oxides, copper (II) oxide
Nitric acid is a yellow fuming corrosive liquid that dissociates into ions even in concentrated solution. Concentrated nitric acid is a strong oxidizing agent. Dilute nitric acid is a strong acid that reacts with metals as an acid or an oxidizing agent. Nitric acid is available as: 1. 2.0 M and 1.0 M volumetric solutions 2. minimum assay 69% density 1.41 g cm^-3 w / w HNO₃ solution 3. 70% "ANALAR" solution (d) commercial solution for use by tradesman.

12.5.01 Nitrous acid, ionization of nitrous acid
Weak acid prepared by acids on nitrites.
Ba(NO₂)₂ + H₂SO₄ --> BaSO₄ + 2HNO₂Heated nitrous acid decomposes to form nitric acid and nitrogen monoxide (nitric oxide). 2HNO₂ --> HNO₃ + NOIonization of nitrous acid, Ka = 4.5 X 10^-4
HNO₂ + H₂O <--> H₃O⁺ + NO₂^-

12.5.1 Prepare nitric acid with sulfuric acid and sodium nitrate
Add concentrated sulfuric acid to sodium nitrate.
BE CAREFUL! HEAT GENTLY. NITRIC ACID VAPOUR FORMS.
NaNO₃(s) + H₂SO₄(aq) --> HNO₃(l) + NaHSO₄(aq)

12.6.0 Sulfuric acid, ionization of hydrogen sulfate ion
See diagram: 12.6.0
Sulfuric acid is a colourless oily liquid available as: 1. 2.0 M (4.0 N) 1.0 M (2.0 N) and 0.5 M (1.0 N) volumetric solutions 2. Minimum assay 97% solution density 1.83 g cm^-3 3. 98% "ANALAR" solution (d) "Battery acid" solution for lead cell accumulators minimum assay 30% density 1.25 g cm^-3 at 20^oC (battery acid). Sulfuric acid is a strong dibasic acid that forms sulfates and hydrogen sulfates a strong oxidizing agent that dissolves copper and a strong dehydrating agent that can remove water from organic compounds. Sulfuric acid is made by the contact process. Sulfur is burned or the ores zinc sulfide or iron sulfide (pyrites) are heated to form sulfur dioxide. The gases pass over vanadium (V) oxide or platinum catalyst at 450^oC to form sulfur trioxide that combines with water to form sulfuric acid.
When coal is burnt, the compounds that contain sulfur can form sulfuric acid, as in the three equations below, to become components of acid rain (rainwater pH = 5.6, acid rain pH < 5)
S(s) + O₂(g) --> SO₂(g) sulfur dioxide
2SO₂(g) + O₂(g) <--> 2SO₃(g) sulfur trioxide
SO₃(g) + H₂O(l) --> H₂SO₄(aq) sulfuric acid
Ionization of hydrogen sulfate ion
HSO₄^- + H₂O <--> H₃O⁺ + SO₄^2-

12.6.1 Sulfuric acid acts as an oxidizing agent
See also 12.3.13: Reactions of concentrated sulfuric acid with copper
BE CAREFUL! YOU ARE USING HOT CONCENTRATED SULFURIC ACID!
Add hot concentrated sulfuric acid to carbon. The reaction forms carbon dioxide and sulfur dioxide.
C(s) + 2H₂SO₄(l) --> CO₂(g) + 2SO₂(g) + 2H₂O(l)
Add hot concentrated sulfuric acid to sulfur. The reaction forms sulfur dioxide and water.
S(s) + 2H₂SO₄(l) --> 3SO₂(g) + 2H₂O(l)
Add hot concentrated sulfuric acid to carbohydrates. The reaction forms carbon dioxide or carbon and water.

12.6.2 Sulfuric acid dehydrating copper (II) sulfate crystals (copper (II) sulfate)
Add drops of sulfuric acid to blue copper (II) sulfate crystals. The crystals turn white as they lose water. Concentrated sulfuric acid combines so readily with water that it can be used as a dehydrating agent, e.g. removing water from hydrated copper (II) sulfate crystals and from other hydrated salts.
CuSO₄.5H₂O(s) <--> CuSO₄(s) + 5H₂O(l)

12.6.3 Sulfuric acid dehydrating sucrose (cane sugar)
Put some sucrose (cane sugar) in a tall beaker. Add drops of concentrated acid to the sugar. BE CAREFUL!
The sugar turns yellow then brown then black and rises in the beaker. It reacts with carbohydrates like sugar and cellulose charring them by removing the elements of water from them and leaving a mass of black carbon behind.
C₁₂H₂₂O₁₁(s) --> 12C(s) + 11H₂O(l)

12.6.3.1 Sulfuric acid dehydrating sucrose (cane sugar)
BE CAREFUL!
Roll paper into a tube and hold it in the middle of a soft plastic container, e.g. ice cream tub. Do not use a glass jar. Fill the container with sugar. Pour just enough water to dampen the sugar down the tube to reach the bottom. Leave to stand for five minutes to allow the water to spread throughout the sugar. Remove the paper tube to leave a hole in the damp sugar. Pour 30 mL of concentrated (98%) sulfuric acid down the hole and onto the top of the sugar. The sugar starts to turn brown, and black in patches. After some minutes bubbles of steam form. The reaction became more vigorous as the material in the container expands. A black cylinder rises out of the jar. Jets of steam spurt out. Heat is given out as the cylinder keeps rising. The black steaming cylinder is spongy carbon. Tap with a spatula to show it is hard, like expanded polystyrene packaging. If the carbon solidifies to make a seal over the top of the jar and the reaction continues deeper in the container, below the seal, pressure may build up to cause an explosion and a shower of black crumbling carbon.

12.6.4 Sulfuric acid in water
BE CAREFUL! WHEN DILUTING STRONG ACIDS ALWAYS ADD THE ACID SLOWLY TO WATER WITH GREAT CARE. NEVER ADD WATER TO THE ACID.
Add concentrated sulfuric acid very slowly to water. Stir the mixture thoroughly each time a small amount of acid is added. Note any change in temperature. Pass hydrogen chloride gas into water. Add acetic acid to water. Acetic acid, a weak acid, produces less heat than the strong acids sulfuric acid and hydrochloric acid.

12.7.0 Properties of bases
Most bases dissolve in water releasing hydroxide ions (OH^-) and react with acids to form salts. In the reaction with dilute sulfuric acid, the base copper (II) oxide the oxide accepts hydrogen ions. So you can say that a base is a proton acceptor.
CuO(s) + H₂SO₄(aq) --> CuSO₄(aq) + H₂O(l)
O²⁺(g) + 2H⁺(aq) --> H₂O(l)
NaOH(s) --> Na⁺(aq) + OH- (aq)
Weak alkalis do not completely ionize in water, e.g. pass ammonia gas through water to form dilute ammonia solution.
NH₄OH(aq) <-- NH₄⁺(aq) + OH^-(aq)
or, using the more modern way of representing this reaction:
NH₄OH(aq) <-- NH₃(aq) + H₂O(l)
A basic oxide is a metal oxide, e.g. CuO. A basic hydroxide is a metal hydroxide that is insoluble in water, e.g. Mg(OH)₂. A base can dissolve in water to form hydroxyl ions and react with acids to form salts. The term "base" includes the alkalis basic oxides and basic hydroxides. Alkalis are bases that are easily soluble in water. The most commonly used alkalis are sodium hydroxide (caustic soda) calcium hydroxide and dilute ammonia solution. An alkali is a hydroxide that dissolves in water to form a solution with pH > 7 and contains hydroxyl ions (OH^-), e.g. NaOH. Alkalis are good electrolytes, turn red litmus blue, and feel slippery. When strong alkalis dissolve in water, they completely ionize.
Be careful! Strong alkalis may burn the skin and cause blindness if splashed in the eyes! Use safety glasses and nitrile chemical-resistant gloves.
NaOH(s) --> Na⁺(aq) + OH^-(aq)
Weak alkalis do not completely ionize in water, e.g. pass ammonia gas through water to form dilute ammonia solution. This solution is shown as NH₃(aq) + H₂O(l) because while "NH₄⁺" ions and" OH^-" ions can be detected, "NH₄OH" cannot be detected.

1. Add a little solid sodium hydroxide, potassium hydroxide, calcium hydroxide and barium hydroxide separately to a little water in separate test-tubes. Shake the test-tubes. Which of the substances are soluble? Place your fingers around each test-tube to see if heat is being produced. You will have noticed that the alkalis are not equally soluble; the order of decreasing solubility is sodium hydroxide, potassium hydroxide, barium hydroxide, calcium hydroxide.
2. Prepare solutions of sodium hydroxide, potassium hydroxide, calcium hydroxide and barium hydroxide. Test each solution as follows: 1. Pour a small amount into a test-tube and place in it a piece of red and a piece of blue litmus paper. 2. Add a few drops of methyl orange solution. 3. Add a few drops of phenolphthalein solution. Record what happens in each case. The alkalis turn red litmus paper to blue, colourless phenolphthalein solution to red and methyl orange solution to yellow. (d) To a solution of each alkali add a little solid sodium carbonate. Observe any changes in the solutions. No gas forms when sodium carbonate is added to solutions of the alkalis so they do not behave like acids. The white solids formed in the calcium hydroxide and barium hydroxide solutions are calcium carbonate and barium carbonate. A solid formed because of a chemical reaction in solution is called a precipitate.

12.7.1 Feel of alkalis
Feel some soap. It feels slippery because it contains alkalis. Prepare very dilute solutions of each sodium hydroxide, potassium hydroxide, calcium hydroxide and barium hydroxide. Moisten your finger tips with each solution and rub your fingers together. What do you feel?

12.7.2 Test the solubility of alkalis
Add to the same amount of water in test-tubes solutions of: sodium hydroxide, potassium hydroxide, barium hydroxide, calcium hydroxide. The relative solubility is in that order. Feel the heat of reaction.

12.7.3 Reactions of alkalis with metals, sodium hydroxide
Most metals do not react with alkalis. However, zinc or aluminium reacts with alkalis to form soluble zincate and hydrogen or soluble aluminate and hydrogen. Strong alkalis should not be stored in aluminium or zinc containers.
Add 5 mL of concentrated sodium hydroxide solution to test-tubes containing: copper, iron, aluminium, zinc. Heat gently if no reaction is observed. Test any gas from the reaction.

12.7.3.1 Recycle an aluminium drink-can as potassium aluminium sulfate, alum.
Waste material can be converted to new substances but full recovery is seldom attained because of incomplete reactions and loss of partially soluble materials. In this experiment, conversion of aluminium scrap metal to alum crystals requires large quantities of sulfuric acid and potassium hydroxide.
Be careful! Use safety glasses and nitrile chemical-resistant gloves. Use 4 g of pieces of an aluminium drink-can cut into thin shavings. Put 150 mL deionized water in a 400 mL beaker. Slowly add 50 ML of 6 M potassium hydroxide solution. Add the shavings, put the beaker on a tripod stand and mat inside a fume cupboard, fume hood. Heat the mixture gently over a small flame for 30 minutes to dissolve most of the aluminium metal. Adjust the flame to keep a controlled bubbling rate in the beaker. Use a glass stirring rod to prevent the metal shavings floating on top of the froth.
BE CAREFUL! KEEP YOUR FACE AWAY FROM THE CAUSTIC SPRAY!
When most of the aluminium has dissolved, turn off the Bunsen burner and filter the hot mixture through a funnel and glass wool plug to remove suspended paint, varnish and unreacted aluminium pieces. Collect the filtrate in a 400 mL, beaker. Cool the filtered solution to room temperature. Transfer the liquid to a 250 mL graduated cylinder to record the total volume of solution. Pour one quarter of this total volume into separate 150 mL beakers. With continuous stirring, acidify each portion of the solution by slowly pouring 20 mL, of 9 M sulfuric acid into the beaker. Be careful! Considerable heat is produced! If any lumps of aluminium hydroxide precipitate are present after adding the sulfuric acid, heat the mixture gently with stirring. Remove the heat when the mixture becomes clear. Cool the solution m an ice bath for about 20 minutes with frequent stirring. Crystals of alum, KAl(SO₄)₂.12H₂O, form in the beaker. Set up a Buchner funnel with all the holes are covered. Clamp the flask to a ring stand, and connect the flask to a water aspirator. After the alum crystals have fully formed in the ice bath, turn on the aspirator and transfer the alum crystals from the beaker to the Buchner funnel. Wash the beaker with 20 mL of 60% ethyl alcohol / water solution to transfer any crystals remaining in the beaker. Add these washings to the Buchner funnel. Run the aspirator for several minutes, allowing the crystals to become moderately dry. Remove the filter paper and alum crystals from the Buchner funnel and put on a watch glass. Divide the mass of alum crystals recovered by the original mass of aluminium present in the solution to show the mass of alum crystals obtained per gram of aluminium metal. Theoretically, 17.5 g alum crystals can be obtained from 1.00 g of aluminium metal. Calculate the percentage yield: = Actual mass alum recovered per g Aluminium / Theoretical mass alum recovered per g Aluminium
1. Dissolving the aluminium: 2Al(s) + 2KOH(aq) + 6H₂O(l) --> 2KAl(OH)₄(aq) + 3H₂(g)
2. Acidifying with sulfuric acid: 2KAl(OH)₄(aq) + H₂SO₄(aq) --> 2Al(OH)₃(s) + K₂SO₄(aq) + 2H₂O(l) 2Al(OH)₃(s) + 3H₂SO₄(aq) --> Al₂(SO₄)₃(aq) + 6H₂O(l)
3. Forming alum crystals: K⁺(aq) + Al³⁺(aq) + 2SO₄^2-(aq) + 12H₂O(l) --> KAl(SO₄)₂.12H₂O(s)

12.7.4 Reactions of alkalis with salts, sodium hydroxide with copper salts
Add drops of 2 M sodium hydroxide solution to 2 M solutions: copper (II) sulfate, copper (II) chloride, copper (II) nitrate. In each case, a blue precipitate results. These solutions contain only the copper (II) ion in common so this ion is the cause of the blue precipitate. Add drops of barium chloride solution to: sulfuric acid, sodium sulfate. In each case the white precipitate is caused by the sulfate ion.
Ba²⁺ + SO₄^2---> BaSO₄(s)

12.7.4.1 Reactions of alkalis with different salts, solubility of hydroxides
Prepare solutions of collection of salts in separate test-tubes, e.g. magnesium sulfate, copper (II) sulfate, iron sulfate, potassium nitrate, calcium chloride. To each solution slowly add a small quantity of sodium hydroxide solution. Note the colour of any precipitate formed and any other change you may observe.
copper (II) sulfate (aq) + sodium hydroxide(aq) --> sodium sulfate(aq) + copper hydroxide(s)
copper ions(aq) + hydroxide ions(aq) --> copper hydroxide(s)
Hydroxide ions (hydroxyl ions) form precipitates with most metal ions. All metallic hydroxides are insoluble in water except sodium hydroxide, potassium hydroxide and ammonia solution. [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.] Calcium hydroxide and barium hydroxides are only slightly soluble. However, all substances dissolve in water to some extent so there is no sharp distinction between soluble and insoluble substances.

12.7.5 Reactions of alkalis with basic oxides, copper oxide
Basic oxides do not react with alkalis.
Add a small quantity, about the size of a split pea of sodium hydroxide solutions to: copper (I) oxide, calcium oxide, magnesium oxide, iron oxide. In each case there is no reaction.

12.7.6 Reactions of alkalis with acidic oxides, carbon dioxide
Acidic oxides react with alkalis to form salt and water.
1. Pass carbon dioxide bubbles through sodium hydroxide solution in a test-tube. Note the size of the bubbles. If the bubbles decrease in size as they rise through the solution, carbon dioxide is being used in a chemical reaction. Carbon dioxide combines with water to form carbonic acid, so a reaction of this acidic oxide with the alkali occurs.
NaOH(aq) + CO₂(g) --> Na₂CO₃(aq) + H₂O(l)
Add dilute hydrochloric acid. Test gases that form from the reaction with: moist litmus paper, a lighted splint, limewater. The production of carbon dioxide confirms that the reaction forms a carbonate.
HCl(aq) + Na₂CO₃(aq) --> CO₂(g) + 2NaCl(aq) + H₂O(l)
2. Pass a slow stream of carbon dioxide bubbles into the bottom of a measuring cylinder containing sodium hydroxide solution. Note any alteration in the size of the bubbles as they rise through the solution. After five minutes stop the flow of carbon dioxide and add 5 mL dilute hydrochloric acid. Test any gas liberated with a lighted splint, pieces of damp red and blue litmus paper, and limewater. The gradual decrease in the size of the ascending carbon dioxide bubbles shows that a reaction involving carbon dioxide occurs. Carbon dioxide and sulfur dioxide combine with water to form acids so you expect a reaction of these acidic oxides with dilute solutions of alkalis to occur. The production of carbon dioxide after adding hydrochloric acid to the solution in the measuring cylinder confirms this view because no gas is produced when hydrochloric acid reacts with sodium hydroxide, the only substance other than water originally present in the cylinder.
sodium hydroxide(aq) + carbon dioxide(g) --> water(l) + sodium carbonate(aq)
3. Pass carbon dioxide through barium hydroxide solution in a test-tube. Filter off the precipitate. Add dilute hydrochloric acid to the precipitate. Identify the gas liberated. The reaction forms barium carbonate.

12.7.7 Reactions of alkalis with amphoteric oxides and hydroxides
Oxides and hydroxides of Al, Pb, Sb, Sn, and Zn are amphoteric. When amphoteric substances react with alkalis, they behave as acidic oxides. When amphoteric substances react with acids, they behave as basic oxides.
1. Add dilute sodium hydroxide solution to: aluminium oxide, zinc oxide, aluminium hydroxide, zinc hydroxide. Heat gently. Note any reactions.
ZnO(s) + 2NaOH(aq) --> Na₂ZnO₂(aq) + H₂O(l)
zinc oxide + sodium hydroxide --> sodium zincate + water
2. Repeat the procedure but using dilute hydrochloric acid instead of dilute sodium hydroxide solution.

12.7.7.1 Action of sodium hydroxide on zinc chloride solution
The addition of aqueous sodium hydroxide to a test-tube containing zinc chloride solution will result in the formation of zinc hydroxide, a white precipitate. On addition of excess aqueous sodium hydroxide, the precipitate reacts with the sodium hydroxide to form a complex salt, sodium zincate
Zn(OH)₂ + 2NaOH --> Na₂Zn(OH)₄
insoluble -->. soluble
The resulting zincate is soluble and so dissolves to give the colourless solution. However, many secondary school chemistry teachers, chemistry textbooks, and handbooks on qualitative analysis, would describe the above reaction of zinc hydroxide with excess sodium hydroxide as "a white precipitate forms that dissolves in excess sodium hydroxide to give a colourless solution". Students are taught that dissolving is a physical change as the solute can be recovered easily and that no new sub stances are formed. Thus the use of the word "dissolve" in the above situation may give students the notion that the disappearance of the precipitate is a physical change, when, in fact, it is a chemical reaction that has taken place. The word "dissolve" has also been erroneously used to describe the following phenomena: 1. Reactions of excess aqueous sodium hydroxide with aluminium hydroxide, resulting in the formation of soluble sodium aluminate. 2. Reactions of aqueous ammonia with silver chloride, zinc hydroxide or copper (II) hydroxide resulting in the formation of soluble complex amines. 3. Reactions of acids with metals, insoluble bases, carbonates and sulfate (IV).
Chemists use the term "dissolve" loosely to mean the disappearance of a solid in a liquid, but they are aware of what is happening, that is, whether a reaction has taken place or whether it is just solvation. However, secondary new substances are formed. The use of the word "dissolve" above may give students the notion that the disappearance of the precipitate is a physical change, when, in fact, it is a chemical reaction that has taken place

12.7.8 Reactions of alkalis with sodium carbonate
Add solid sodium carbonate to a solution of each alkali. No gas forms. White precipitates of carbonates form in the barium hydroxide and calcium hydroxide solutions only.
Na₂CO₃(s) + Ca(OH)₂(aq) --> CaCO₃(s) + NaOH(aq)

12.8.0 Acid-base neutralization
In neutralization reactions, an acid and a base react in such proportions as to form a neutral solution of a salt and water. In the home, wool dresses spotted with another colour from acids or bases can be restored to original colour by neutralization. The reaction is between the hydrogen ions and the hydroxide ions.
H⁺(aq) + OH^-(aq) --> H₂O

1. Add 2 mL nitric acid or tartaric acid to 4 cm water in a test-tube. Add blue litmus paper which then turns red. Add drops of dilute ammonia solution or washing soda solution, Na₂CO₃.10H₂O, but stop when the solution just turns blue again. The acid is now neutralized by the alkali.
2. Put 3 mL dilute sodium hydroxide solution on a watch glass. Use a dropper to add dilute hydrochloric acid drop by drop while stirring continuously and testing the mixture with a fresh piece of litmus paper after each drop is added. You can get a mixture where the litmus paper is neither red nor blue but a tint midway between these two colours. This mixture does not have the taste of an acid or the feel of an alkali. The solution has the properties neither of an acid nor of an alkali. The acid and alkali have neutralized each other. Evaporate the solution to dryness by heating the watch glass over a beaker of boiling water. A small quantity of solid appears on the watch glass. This crystalline solid is sodium chloride, common salt. Water is also a product of the reaction of sodium hydroxide with hydrochloric acid.
3. Pour about 5 mL of dilute solutions of sodium hydroxide, potassium hydroxide, calcium hydroxide and ammonia solution into separate test-tubes. Neutralize each alkali solution with dilute hydrochloric acid then evaporate the resulting solutions to dryness. Do not taste the residues. Repeat the procedure with dilute sulfuric acid, dilute nitric acid or dilute acetic acid.
4. magnesium hydroxide in small amounts to dilute sulfuric acid until excess solid is present. Filter the mixture and test the filtrate with pieces of red and of blue litmus paper. Evaporate the filtrate to dryness. Many metallic hydroxides react with acids to produce water and a salt in the same way as alkalis do.
magnesium hydroxide(s) + sulfuric acid(aq) --> water(l) + magnesium sulfate(aq)
Metallic hydroxides that behave in this way with acids and are insoluble in water are called basic hydroxides, The three classes of compounds - alkalis, basic oxides and basic hydroxides - are representatives of a group of substances called bases.

12.8.1 Ammonia with sulfuric acid
Pass ammonia through sulfuric acid. The common fertilizer ammonium sulfate or sulfate of ammonia forms.
2NH₃(g) + H₂SO₄(aq) --> (NH₄)₂SO₄(aq)

12.8.2 Sodium hydroxide with hydrochloric acid
Put 10 drops of dilute sodium hydroxide solution on a watch glass. Add drops of dilute hydrochloric acid and stir. Test the mixture with litmus paper after adding each drop of hydrochloric acid.
When the litmus is neither red nor blue, but between the two colours, stop adding drops of acid. Wet the tip of the finger with the mixture. Rub the mixture between the fingers - it does not feel slippery, so the solution is not alkaline. When the correct quantities of hydrochloric acid and sodium hydroxide are mixed a solution forms that has the properties neither of the acid nor of the alkali. The acid and alkali have neutralized each other. Evaporate the neutralized solution to dryness by heating the watch glass over a beaker of boiling water. Crystals of sodium chloride appear on the watch glass.
HCl(aq) + NaOH(aq) NaCl(s) + H₂O(l)
acid + alkali --> salt + water
Add dilute hydrochloric acid to dilute solutions of: sodium hydroxide, potassium hydroxide, calcium hydroxide aqueous ammonia solution. Evaporate to dryness. Describe the salt formed. Repeat the experiment with: dilute sulfuric acid, dilute nitric acid, dilute ethanoic acid (acetic acid).

12.8.3 Simple titration of acids and bases
See also: Commercial pH soil test kit
Titration is an experimental method for measuring the concentration of a solution. Measure the volume of the solution "A" needed to react with a given volume of solution "B". For HCl and NaOH titration, molarity "A" X volume "A" = molarity "B" X Volume "B". The end point in a titration occurs when an indicator changes colour.
Use a medicine dropper or a teat pipette as a simple burette. The drops must always be the same size. Within experimental error, when the same dropper is used, the same number of drops of alkali is needed to neutralize the same number of drops of acid. When the concentration of the acid is known, the concentration of the base can be estimated by comparing the numbers of drops of acid and drops of base that just react. Drop 100 drops of water from a medicine dropper into a measuring cylinder. Calculate the volume of one drop. Measure 25 mL of 2 M sodium hydroxide solution in the measuring cylinder and pour into an evaporating dish. Add 2 drops of phenolphthalein solution and note the red colour of the indicator. Wash the medicine dropper with the 2 M hydrochloric acid to get rid of remaining sodium hydroxide. Add 2 M hydrochloric acid a drop at a time to the solution in the evaporating dish. Stir as each drop is added. Note the number of drops added until the colour just disappears completely. Calculate the volume of added acid. Heat the solution until almost dry. Use gentle heat to avoid spattering. Describe the appearance of the residue.
NaOH(aq) + HCl(aq) NaCl(aq) + H₂O(l)
40 + 36.5 --> 58.5 + 18
Weight of NaOH in 25 mL of 2M solution = (25 X 2 X 40 / 1000) g. The weight of sodium chloride expected in the evaporating dish = 58.5 X (25 X 2 X 40 / 1000) / 40 = 2.925 g.

12.8.4 Titrate dilute hydrochloric acid with sodium hydroxide solution, using a burette
See diagram. 12.8.4
Pour a hydrochloric acid solution of known concentration (such as 0.11 mol / L) into a clear, dry burette until the liquid level is above the "0" line. Fix the burette vertically with a burette clamp. Rotate the stopcock carefully to set the lowest point of the liquid meniscus exactly to "0" and to make simultaneously the tapered portion of the burette full of the acid solution without any air bubbles in it. Use a pipette to transfer 20 mL of the sodium hydroxide solution to a conical flask. Add two drops of phenolphthalein to the flask. The solution immediately turns red. Stand the flask on a piece of white paper under the burette. While adding drop by drop the acid solution from the burette, swirl the flask constantly so that mixing of the base and acid solutions is rapid and thorough. Note any change in the solution colour. The neutralization is exactly completed and the end point occurs when half a drop or one drop of the acid solution turns the pale red solution colourless in the flask immediately after swirling. Stop the titration and record the burette meniscus reading. Read the volume of the used hydrochloric acid solution. Calculate the concentration of the sodium hydroxide solution according to the related chemical equation.

12.8.4.1 Titrate dilute hydrochloric acid with sodium hydroxide solution, using a burette (second method)
Use a burette containing 50 mL of 0.5 M sodium hydroxide. Use a pipette to put 10 mL of 0.5 M hydrochloric acid in a beaker under the burette. Add 2 drops of phenolphthalein to the beaker. Stand the beaker on white paper under the burette containing sodium hydroxide. Add a drop at a time of the sodium hydroxide from the burette and stir the beaker with a swirling motion. Note the colour change when a drop of acid disappears after the solution is swirled. The end point occurs when the drop does not change colour after swirling. The solution is now neutral. Test the neutral solution with litmus paper. Pour 5 mL of the neutral solution into an evaporating dish. Heat to dryness and weigh when cool.

12.8.5 Effect of carbon dioxide on acid-base titration
Add 2 drops phenolphthalein to 100 mL of deionized water. Add 2 drops of 0.1 M sodium hydroxide. The reaction forms a red colour. Swirl vigorously for one minute. The red colour fades because of absorption of carbon dioxide from the air.

12.8.6 Heat of neutralization titration
The end point occurs at maximum temperature.
Use 25 mL of dilute sodium hydroxide solution. Note the original temperature. Add 1 mL of 2 M hydrochloric acid, stir with a thermometer and note the temperature. Continue to add 1 mL of the acid and note the temperature. Use graph paper to plot temperature rise against volume of acid added. Read from the graph the maximum temperature rise and volume of acid that neutralized the sodium hydroxide solution.
Calculation: 25 X 1 / 1000 X concentration of NaOH = volume of HCl X 1 / 1000 X 2

12.8.7 Microscale titration, sodium hydroxide with dilute acids
See diagram 12.8.7
Use a sodium hydroxide solution to titrate a standardized acid solution. Use the sodium hydroxide solution to titrate an unknown acid. Calculate the concentration of the unknown acid
Conventional titration vs microscale titration: Conventional titration requires burettes, bulb pipettes and litres of solutions. Ten microscale titrations will use fewer solutions than used for one conventional titration Burettes are large and fragile so spillages and breakages of glassware occur sometimes. When an operator using conventional titration does three measurements, a microscale operator can do six to ten measurements. Microscale titration does require different hand and finger skills than conventional titration and involves some differences in the calculation methods.
1. Prepare 0.1 M monoprotic acid solution, 0.1 M sodium hydroxide solution, an acid of unknown molarity. Use two 2 mL graduated glass pipettes, graduated to 0.01 mL. Attach disposable pipette tips. If they fall off the ends of the pipettes, seal with silicone putty. Attach a 5 mL plastic syringe to each pipette with silicone tubing. Lubricate the syringe with glycerol. Transfer 15 mL of acid solution into a wide neck bottle and one drop of phenolphthalein indicator. Transfer 15 mL sodium hydroxide into another wide neck bottle. Clamp the two pipettes on a stand with a double clamp. Rinse then fill the pipettes by drawing solution up into the pipette with the syringe. Record pipette volumes to 0.001 mL. Use a 10 mL Erlenmeyer flask containing water for comparison when detecting the faint pink of the endpoint. Use pipette volumes in excess of 1.000 mL to provide four significant figures in the volume measurements
2. Deliver 1.25 mL of 0.1 M monoprotic acid into a 10 mL Erlenmeyer flask. Add 0.1 M sodium hydroxide until the colour changes to a faint pink. Record the final pipette volumes. Drop volume is less than 0.02 mL. Deliver 1.25 mL of the unknown acid into a 10 mL Erlenmeyer flask. Add 0.1 M sodium hydroxide until the colour changes to a faint pink. Record the final pipette volumes For each titration, calculate the ratio of acid volume to alkali volume to allow concordance between titrations. Discard discordant values and calculate the mean for accepted values.
Ratio 1 = volume 0.1 M monoprotic acid / volume 0.1 M sodium hydroxide
Ratio 2 = volume unknown acid / volume 0.1 M sodium hydroxide
Concentration of unknown acid = (Ratio 1 / Ratio 2) X concentration 0.1 M monoprotic acid, e.g. If Ratio 1 = 1.158, Ratio 2 = 1.079, molarity of unknown acid = (1.158 / 1.079) X 0.1 = 0.1073 M.
Addition of a drop of the indicator to the acid solutions being measured dilutes the acids, e.g. 0.02 mL of indicator solution added to 15 mL of acid solution dilutes the acid to affect the measured concentration at the fourth significant figure. However, this titration is done twice, with the "known" standard acid and with the "unknown" acid, so the errors will cancel at the fourth significant figure when the acids do not differ greatly in concentration.

12.9.4 Rainbow reactions, t-Butyl chloride (2-chloro-2-methylpropane) with sodium hydroxide
Prepare a pH 12 solution by adding 10 drops of 0.1 M NaOH to 100 mL water, in a 250 mL beaker. Add universal indicator to produce a distinct colour. Start with universal indicator. Use a second 250 mL beaker to mix by pouring the solution back and forth between the two beakers or put a magnetic bar into the solution and start the stirrer motor at a fast rate. Add 15 drops of t-butyl chloride (2-chloro-2-methylpropane) to the solution and begin mixing. Observe any colour changes. After 40 seconds add universal indicator and observe any colour changes. The full range of colour changes (purple, blue, cyan, emerald-green, lime-green, yellow, orange, orange-red, take about two minutes. The changes in the middle are more rapid than the changes at either extreme. Use different indicators to show different colour changes and different induction times:

Indicator	Colour Change	Induction Time	Indicator	Colour Change	Induction Time
Methyl red	Yellow	Initial	Phenolphthalein	Pink	Initial
"	Orange	40 seconds	"	Pale pink	40 seconds
"	Green	45 seconds	"	Colourless	45 seconds
Bromothymol blue	Blue	Initial	Bromophenol blue	Blue	Initial
"	Green	40 seconds	"	Green	75 seconds
"	Yellow	45 seconds	"	Yellow	80 seconds
Thymol blue	Blue	Initial	m-Cresol purple	Violet	Initial
"	Green	50 seconds	"	Red	52 seconds
"	Yellow	52 seconds	"	Yellow	54 seconds

Different formulations of universal indicator may give differing times and colour changes. For a wide range universal indicator, double the amounts of each reactant. The reaction is an "S_N1" reaction, i.e. a nucleophilic substitution reaction, in which the chlorine radical is replaced by an hydroxyl radical. As H⁺ions are produced in solution in the reaction, the OH^- are gradually neutralized as the reaction proceeds, and eventually excess H+ are produced. Thus, the pH of the solution progressively falls because of reaction.
(CH₃)₃-C-Cl + H₂O --> (CH₃)₃-C-OH + HCl

2. Prepare two solutions: 0.1 M 2-chloro-2 methylpropane (t-butyl chloride) in ethanol (1 g per 100 mL) and 0.01 M sodium hydroxide. Put 5 mL 0.1 M C₄H₉Cl in a test-tube. In another test-tube put 5 mL 0.1 M NaOH, 10 mL water and two drops of any one of the following indicators. Mix the solutions back and forth once and observe for the colour change that occurs after an induction period. With equal volumes 0.01 M sodium hydroxide and 0.1 M C₄H₉Cl the colour changes with universal indicator were: Purple to blue (on mixing) blue to green (after 12 seconds) green to yellow (after 15 seconds) yellow to orange (after 25 seconds total). Cooling the solutions greatly slows the reaction, increasing the induction period, e.g. with iced water, the methyl red change took more than 50 seconds.

Indicator	Colour Change	Induction Time
Methyl red	Yellow to red	6 seconds
Bromothymol blue	Blue to yellow	5 seconds
Thymol blue	Blue to yellow	10 seconds
Phenolphthalein	Red to colourless	11 seconds
Universal indicator	Purple to blue Blue to green Green to pink	Instantly 8 seconds 10 seconds

12.20.0 Boric acid, ionization reaction
Orthoboric acid, trioxoboric acid (III) acid, boracic acid, sassolite, H₃BO₃ is a weak acid. White to colourless triclinic crystals, m.p. 169^oC, occurs in volcanic steam vents, slightly soluble in cold water, used to make borosilicate glass, used in buffer solutions, detergents and in pharmacy, e.g. "boracic powder" for eye infections. Action of continuous heat: boric acid, H₃BO₃--> metaboric acid + water, H₂B₄O₄ --> tetraboric acid (pyroboric acid) H₂B₄O₇ --> boric oxide (anhydrous boron (III) oxide) B₂O₃. Boric oxide is an intermediate oxide, as is aluminium oxide, with weak acidic and basic properties. Borax is hydrated sodium borate. When heated it fuses to form clear glass that can dissolve metal oxides to give characteristic colours of the borax bead test.
Ionization reaction, Ka = 6.0 X 10^-10
H₃BO₃+ H₂O <--> H₃O⁺ + H₂BO₃^-

12.20.1 Prepare boric acid crystals
Use about 5 g of boric acid crystals. Pour some into 2 cm boiling water in a test-tube and allow to dissolve. Continue adding crystals and heat to boiling until all crystals dissolved. Allow to cool to see fine white crystals form.