School Science Lessons
Topic 12 Chemical equations, types of chemical reactions, acids
Updated 2009-11-14
Please send comments to: J.Elfick@uq.edu.au
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Table of contents
12.1.0 Chemical equations and ionic equations, conservation of mass
12.2.0 Types of chemical reactions
12.3.0 Properties of acids
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, sulfur trioxide
12.9.0 Phosphoric acid, ionization reaction
12.10.0 Boric acid, ionization reaction, prepare boric acid crystals
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 reactions, double exchange, metathesis
12.2.1.1 Sodium carbonate with calcium hydroxide, double decomposition
12.2.1.2 Cobalt chloride with calcium hydroxide, double decomposition
12.2.1.3 Sodium carbonate with zinc sulfate, double decomposition
12.2.1.4 Sodium carbonate with magnesium sulfate, double decomposition
12.2.1.5 Sodium hydroxide with copper sulfate
12.2.1.6 Sodium hydroxide with cobalt chloride
12.2.1.7 Sodium hydroxide with iron sulfate
12.2.1.8 Sodium hydroxide with alum (Al2(SO4)3.K2(SO4).24H2O, aluminium potassium sulfate)
12.2.1.9 Prepare sodium hydroxide solution
12.2.2 Synthesis reactions (combination reactions)
12.2.2.1 Heat iron with sulfur
12.2.2.2 Heat copper with iodine
12.2.2.3 Heat iron with iodine
12.2.2.4 Combination reaction to form a hydrate, CuSO4.5H2O
12.2.3 Decomposition reactions
12.2.3.1 Examples of decomposition reactions
12.2.3.2 Decomposition of copper carbonate to prepare copper oxide
12.2.3.3 Decomposition of zinc carbonate to prepare zinc oxide
12.2.4 Displacement reaction (substitution reaction)
12.2.4.1 Iron displaces copper in copper sulfate solution
12.2.4.2 Prepare zinc sulfate crystals
12.2.4.3 Magnesium displaces hydrogen in ethanoic acid (acetic acid, CH3COOH)
12.2.5 Acid-base reactions
12.2.6 Redox reactions (oxidation-reduction reactions, electron transfer reactions)
12.2.7 Conditions for chemical reactions to occur, sulfuric acid with sodium chloride
12.3.0 Properties of acids
12.3.0.1 Amphoteric substances
12.3.0.2 Polyprotic acids
12.3.0.3 Strong acids and weak acids
12.3.0.4 pH
12.3.0.5 Ionization reaction of carbonic acid
12.3.1 Taste of acids, solid acids in the home
12.3.2 Dilute acids with metals, hydrochloric acid
12.3.2.1 Dilute acids with metals, sulfuric acid, hydrochloric acid, ethanoic acid
12.3.3 Dilute acids with metals, sulfuric acid with iron
12.3.3.1 Dilute acids with metals, sulfuric acid with aluminium
12.3.3.2 Magnesium with sodium hydrogen sulfate
12.3.3.3 Iron with sodium hydrogen sulfate
12.3.4 Dilute acids with non-metals, carbon, sulfur
12.3.5 Dilute acids with basic oxides (metal oxides), copper (II) oxide
12.3.5.01 Copper sulfate from copper oxide
12.3.5.1 Dilute acids with basic oxides, (metal oxides) zinc oxide
12.3.6 Dilute acids with hydroxides, magnesium hydroxide
12.3.7 Dilute acids with hydroxides, sodium hydroxide
12.3.7.1 Dilute acids with sodium hydroxide
12.3.7.2 Dilute hydrochloric acid with hydroxides
12.3.8 Dilute acids with acidic oxides (non-metal oxides), carbon dioxide, sulfur dioxide
12.3.9 Dilute acids with carbonates, common carbonates
12.3.9.1 Dilute hydrochloric acid with calcium carbonate
12.3.9.2 Dilute hydrochloric acid with sodium carbonate
12.3.9.3 Dilute tartaric acid with sodium carbonate
12.3.9.4 Dilute tartaric acid with egg shell, soil, wood ash
12.3.9.5 Dilute hydrochloric acid with calcium carbonate
12.3.9.6 Dilute sulfuric acid with calcium carbonate
12.3.10 Dilute acids with sodium hydrogen carbonate
12.3.10.1 Dilute acids with calcium hydrogen carbonate
12.3.11.0 Dilute nitric acid with copper
12.3.11.1 Nitric acid with metals
12.3.12 Concentrated acids with metals, nitric acid with copper
12.3.13 Concentrated acids with metals, sulfuric acid with copper
12.3.14 Concentrated acids with non-metals, carbon, sulfur
12.3.15 Acids with salts
12.4.0 Hydrochloric acid
3.42 Prepare hydrogen chloride, HCl
3.42.02.1 Prepare hydrochloric acid
3.42.1 Tests for hydrogen chloride
12.5.0 Nitric acid
12.5.01 Nitrous acid, ionization of nitrous acid
12.5.1 Prepare nitric acid, sulfuric acid with sodium nitrate
12.6.0 Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide
12.6.0.1 Formation of acid rain, SOx, by burning sulfur or sulfur compounds
12.6.0.2 The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite).
12.6.0.3.1 Sulfur dioxide to sulfuric acid 1.
12.6.0.3.2 Sulfur dioxide to sulfuric acid 2.
12.6.0.4 Hydrogen peroxide oxidizes sulfur dioxide to 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
12.1.0 Chemical equations and ionic equations, conservation of mass
Chemical reactions are influenced by the conditions under which they take place and, being reversible, may reach a state of equilibrium. Chemical reactions occur at different rates and changing the nature of the reactants, temperature, or concentration, or introducing a catalyst, may alter these. Chemical reactions may be reversible. Reversible chemical reactions may reach a state of dynamic balance known as equilibrium which, when disturbed, will be re-established.
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. 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: Conservation of matter | See 13.7.7 Prepare carbon dioxide by heating hydrogen carbonates | See 19.2.13: Prepare fruit salts
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 gas 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 Types of chemical reactions
Chemical reactions involve energy changes. All chemical reactions involve energy transformations. The spontaneous directions of chemical reactions are towards lower energy and greater randomness. In a chemical reaction, a chemical change occurs where elements or compounds (reactants) form new substances (products). Specific criteria can be used to classify chemical reactions. The main types of chemical reactions are precipitation reactions, synthesis reactions, decomposition reactions, displacement reactions, acid-base reactions,  redox reactions and polymerisation reactions that produce large molecules with repeating units.
12.2.1 Precipitation reactions, double decomposition reactions, double exchange, metathesis
Precipitation reactions result in the appearance of a solid from reactants in aqueous solution. Salts can be made by the action of acids with alkalis, carbonates, metals, metal oxides and by replacement and double decomposition reactions. A salt contains a metal and part of an acid, e. g. copper sulfate from sulfuric acid, sodium chloride from hydrochloric acid. A salt is a compound formed when the hydrogen of an acid is replaced by a metal. For example, when zinc reacts with hydrochloric acid it replaces the hydrogen and forms the salt, zinc chloride. The hydrogen comes away as hydrogen gas. In a double decomposition reaction the radicals are exchanged. A double decomposition reaction can occur if one of the products (substances formed) is insoluble or is a gas.
When sodium chloride solution is added to silver nitrate solution an 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 can be separated from the sodium nitrate solution.
AgNO3 (aq) + KCl (aq) --> KCl (s) + KNO3 (aq)
silver nitrate + potassium chloride --> silver chloride + potassium nitrate
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
12.2.1.1 Sodium carbonate with calcium hydroxide, double decomposition
To half a test-tube of sodium carbonate solution add limewater (calcium hydroxide solution). Describe the reaction which occurs. A white solid (precipitate) of calcium carbonate forms.
sodium carbonate + calcium hydroxide --> calcium carbonate + sodium hydroxide.
12.2.1.2 Cobalt chloride with calcium hydroxide, double decomposition
To half a test-tube of cobalt chloride solution add limewater (calcium hydroxide solution). Describe the reaction which occurs. A pink / mauve precipitate of cobalt hydroxide forms.
cobalt chloride + calcium hydroxide --> cobalt hydroxide + calcium chloride
12.2.1.3 Sodium carbonate with zinc sulfate, double decomposition
Add sodium carbonate solution to the zinc sulfate solution. Filter the mixture to obtain the precipitate of zinc carbonate. Allow the filter paper, unfolded, to dry, and scrape off the white powder.
sodium carbonate + zinc sulfate --> sodium sulfate + zinc carbonate
12.2.1.4 Sodium carbonate with magnesium sulfate, double decomposition
Add magnesium sulfate solution to sodium carbonate solution. Filter the solution to obtain the white precipitate of insoluble magnesium carbonate.
magnesium sulfate + sodium carbonate --> magnesium carbonate + sodium sulfate
12.2.1.5 Sodium hydroxide with copper sulfate
Add 5 mL of sodium hydroxide solution to the same volume of copper sulfate solution. Describe what you see. A blue precipitate of copper hydroxide forms. sodium hydroxide + copper sulfate –> copper hydroxide + sodium sulfate.
12.2.1.6 Sodium hydroxide with cobalt chloride
Add 5 mL of sodium hydroxide solution to the same volume of cobalt chloride solution. Describe what you see. A blue-green precipitate of cobalt hydroxide forms.
Sodium hydroxide + cobalt chloride –> cobalt hydroxide + sodium chloride.
Repeat experiment 153, substituting a dilute solution of cobalt chloride for the copper sulfate solution.
12.2.1.7 Sodium hydroxide with iron sulfate
Add 5 mL of sodium hydroxide solution to a dilute solution of iron sulfate. Stopper the test-tube, shake well, and leave to stand. Describe what you see. A green precipitate forms that turns brown. The green iron hydroxide first forms, but it soon reacts with oxygen gas to form a different type of iron hydroxide which is brown.
12.2.1.8 Sodium hydroxide with alum (aluminium potassium sulfate, Al2(SO4)3.K2(SO4).24H2O)
Add 5 mL of sodium hydroxide solution to a dilute solution of alum.. Describe what you see. A faint white precipitate forms. The part of alum that reacts with the alkali is aluminium sulfate.
12.2.1.9 Prepare sodium hydroxide solution
Make a concentrated sodium carbonate solution and add to it calcium hydroxide solution. Shake well for three minutes. Describe what you see. A white precipitate forms. Filter into a clean bottle and label it. This double decomposition reaction between solutions of sodium carbonate and calcium hydroxide occurs because one of the products, calcium carbonate, is insoluble. The filtrate is sodium hydroxide solution.
sodium carbonate + calcium hydroxide --> sodium hydroxide + calcium carbonate (precipitate)
Test the sodium hydroxide solution with universal indicator paper. The universal indicator paper turns blue or violet.
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
12.2.2.1 Heat iron with sulfur
See diagram 12.2.1: FeS
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.
12.2.2.4 Combination reaction to form a hydrate, CuSO4.5H2O
Prepare white copper sulfate by heating a finger width of blue copper sulfate crystals in an evaporating basin, while stirring with a glass rod. Do not overheat, and stop heating as soon as the chemical has turned white. Leave to cool. Add drops of water until the powder is blue but still dry. The white compound has combined chemically with the water to form the blue compound, because all the added water has disappeared. This reaction is the combination of an anhydrous salt with water to form a hydrate.
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
12.2.3.1 Examples of 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.
potassium nitrate (s) --> potassium nitrite (s) + oxygen (g)
copper nitrate (s) --> copper oxide (s) + nitrogen dioxide (g) + oxygen (g)
3. Repeat using sodium carbonate, magnesium carbonate, calcium carbonate, lead carbonate.
4. Repeat using sodium sulfate, magnesium sulfate, zinc sulfate and copper (II) sulfate.
zinc sulfate (s) --> zinc oxide (s) + sulfur trioxide (g)
copper (II) sulfate (s) --> copper oxide (s) + sulfur trioxide (g)
5. Repeat using magnesium carbonate
magnesium carbonate (s) --> magnesium oxide (s) + carbon dioxide (g)
12.2.3.2 Decomposition of copper carbonate to prepare copper oxide
1. Prepare copper carbonate by mixing sodium carbonate solution and copper sulfate solution. Pour off the liquid when the copper carbonate has settled in the test-tube. Heat to evaporate remaining liquid and heat more strongly to form the oxide. The oxide could be purified by washing with water, using a filtration apparatus. Most carbonates decompose to form a metal oxide and carbon dioxide, e.g. copper carbonate.
2. Prepare copper carbonate by adding half a test-tube of copper sulfate solution to the same quantity of sodium carbonate solution. Leave the insoluble copper carbonate to settle then pour off the liquid above the precipitate. Gently heat the carbonate just enough to drive off the remaining water as steam. Fit a stopper and bent tube to the test-tube. With the end of the bent glass delivery tube dipping into the limewater solution, heat the copper carbonate more strongly. The copper carbonate turns black and the limewater turns milky. The copper carbonate has been decomposed by the heat into black copper oxide and carbon dioxide gas which turns limewater milky.
12.2.3.3 Decomposition of zinc carbonate to prepare zinc oxide
Heat part of the zinc carbonate, from experiment 163, in an evaporating basin, or stand fairly strongly, and until the white powder turns yellow. You now have zinc oxide. Allow it to cool. Heat again. Note the colour changes. Zinc oxide is yellow when hot, white when cold. The carbonate decomposed on heating to form zinc oxide and carbon dioxide. Most carbonates decompose to form a metal oxide and carbon dioxide, e.g. zinc carbonate.
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 3.71 Reactions of ions in solutions | See12.14.1: Zinc displaces lead from lead nitrate solution
12.2.4.1 Iron displaces copper in copper sulfate solution
1. Prepare iron sulfate crystals. Add iron filings to sodium hydrogen sulfate solution in a test-tube, and heat. When there is no further reaction, filter the mixture and pour a finger width of sulfate solution of the filtrate into a evaporating basin. Describe what you see. Pale green crystals of iron sulfate form.
2. Add iron filings to half a test-tube of copper sulfate solution. Leave until the colour of the solution changes from blue to pale green. Filter the solution. Pour part of the filtrate into an evaporating basin. Describe what you see. Pale green crystals of iron sulfate form no different from the crystals formed in the previous experiment.
3. Place copper sulfate crystals in the test-tube and add a quarter of a test-tube of water. Shake to get a blue solution. Drop in the iron nail, which must not be rusty. Leave for ten minutes. Take out the nail. Describe what you see. The iron nail has turned a pinkish colour, due to a deposit of copper on it.
Iron + copper sulfate –> copper + iron sulfate
Repeat the experiment but use a finger width of iron filings instead of the nail. Leave the iron nail test-tube and contents for a few hours. Describe what you see. The blue liquid turns a pale green sulfate colour. The iron has completely displaced the copper in the copper sulfate, forming a solution of iron sulfate that is pale green.
Repeat the experiment using a centimetre of magnesium ribbon instead of the iron nail. Observe the metal a few moments after it has been in the copper sulfate solution. Describe what you see. The metal turns pink or copper coloured. Note whether iron or magnesium metal reacts more quickly. Magnesium is a more reactive metal than iron.
12.2.4.2 Prepare zinc sulfate crystals
Add pieces of granulated zinc to sodium hydrogen sulfate solution in a test-tube, and heat. When there is no further reaction all the blue colour has disappeared, filter the mixture. The filtrate is zinc sulfate solution.
12.2.4.3 Magnesium displaces hydrogen in ethanoic acid (acetic acid, CH3COOH)
All acids contain hydrogen, and many metals can displace it, thus setting the hydrogen gas free. The acid in vinegar is ethanoic acid. Add 3 cm of magnesium to half a test-tube of vinegar. As soon as bubbles of hydrogen gas are coming well, hold your thumb or finger over the mouth of the test-tube for half a minute to trap a quantity of hydrogen gas. Quickly hold the open test-tube to the spirit burner flame. Describe what happens. A small pop or squeak occurs. It is a minor explosion. If you got no result, repeat the procedure of trapping the gas and igniting it in the flame. When hydrogen gas mixes with air it explodes, i.e. combines extremely rapidly with the oxygen gas in the air.
Repeat the experiment using tartaric acid (CHOHCOOH)2 ) instead of ethanoic acid.
Magnesium displaces hydrogen in tartaric acid
Repeat the experiment with iron filings instead of magnesium. You may have to heat the mixture.
12.2.5 Acid-base reactions
Acid-base reactions involve transfer of protons from donors to acceptors. An acid dissociates in water to produce positive hydrogen ions, H+, that is solvated to produce hydronium ions (hydroxonium ions, oxonium 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 hydroxide ions, OH-.
NaOH --> Na+ (aq) + OH- (aq)
Acids react with bases to from salts and water. The products are neither acidic nor basic so this reaction is called neutralization.
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)
Redox reactions involve a transfer of electrons and a change in oxidation number. 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 form, 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 gas 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)
12.3.0 Properties of acids
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 H2SO4 molecules. 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.
Using the Bronsted-Lowry definition of acids and bases an acid donates a proton (H+) to another substance and a base accepts a proton from another substance. When sulfuric acid dissociates in water it donates a proton (H+) to the water molecule. So in this reaction the water molecule acts as a base.
H2SO4 + H2O --> HSO4- + H3O+
When ammonia dissolves in water, ammonia accepts a proton and so it is the base. So in this reaction the water molecule acts as an acid.
NH3 + H2O <--> NH4+ + OH-
12.2.7 Conditions for chemical reactions to occur, sulfuric acid with sodium chloride
See 3.71.1: Solubility table and solubility rules | See also 3.71: Reactions of ions in solutions
1. 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) --> Na2SO4 (aq) + 2HCl (g)
NaCl (s) + H2SO4 (l) --> NaHSO4 (aq) + HCl (g)
2. 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)
12.3.0.1 Amphoteric substances can act as an acid or a base. In the above reactions water is acting as a base with sulfuric acid and is acting as an acid with ammonia. Similarly, bicarbonate ion can act as an acid to donate a proton to form carbonate ion: HCO3- + H2O <--> CO32- + H3O+
Also, bicarbonate ion can act as a base to accept a proton to form carbonic acid: HCO3- + H2O <--> H2CO3 + OH-
12.3.0.2 Polyprotic acids can donate more than one proton, e.g. carbonic acid.
H2CO3 + H2O <--> HCO3- + H3O+ (The first proton to be donated to a water molecule.)
HCO3- + H2O <--> CO32- + H3O+ (The second proton to be donated to a water molecule.)
12.3.0.3 Strong acids and weak acids
A strong acid completely dissociates into ions, e.g. nitric acid.
HNO3 (aq) + H2O --> NO3- (aq) + H3O+ (aq)
A weak acid only partly dissociates into ions, e.g. acetic acid.
CH3COOH + H2O <--> CH3COO- + H3O+
So describing acids and bases as strong or weak only refers to their reaction with water and has nothing to do with concentration or the number of moles in a volume. The strong acids are perchloric acid (HClO4), hydrochloric acid (HCl), hydrobromic acid( Hbr), hydroiodic acid( HI), nitric acid (HNO3), and sulfuric acid (H2SO4). Any other acid is a weak acid because it does not completely dissociate in water.
Strongest acid
 HClO4 + H2O --> H3O+ + ClO4- Weakest base
.
 HCl + H2O --> H3O+ + Cl-
 .
.
 H3PO4 + H2O <--> H3O+ + H2PO4-
 .
Weakest acid
 CH3COOH + H2O <--> H3O+ + CH3COO-
 Strongest base
.
 NH4+ + H2O <--> H3O+ + NH3
 .
.
 H2O + H2O <--> H3O+ + OH-
 .
12.3.0.4 pH
Water can transfer a proton from one molecule to another, autionization.
2H2O <--> H3O+ + OH-
and
H2O <--> H+ + OH-
The product of hydrogen ion concentration, [H+] and hydroxide ion concentration, [OH-] = the constant, Kw
Kw = [H+] X [OH-] = 1.00 X 10-14
So [H+] = 10-7 and [OH-] = 10-7
The hydrogen ion concentration is very small in pure water so the concentration is describes in terms of its negative log
pH is the negative log of the hydrogen ion concentration, pH = -log[H+], so hydrogen ion concentration, [H+] = 10-pH.
So acidic solutions have a high [H+} and low pH values . Basic solutions have low [H+] and high pH values.
A solution that is neither acidic nor basic, a neutral solution, has [H+] = [OH-], so pH = 7. A more acid solution has pH approaching 1. A more basic solution has pH approaching 14.
12.3.0.5 Ionization reaction of carbonic acid
H2O (l) <--> H+ (aq) + OH- (aq)
2H+ (aq) + CO32- (aq) <--> H2CO3 (aq) carbonic acid
CO2 + H2O <--> H3O+ + HCO3-, K1 = 4.4 X 10-7
HCO3- + H2O <--> H3O+ + CO32-, Ka = 4.7 X 10-11
12.3.1 Taste of acids, solid acids in the home
BE CAREFUL! NEVER TASTE ACIDS IN THE LABORATORY!
See appendix: Citric acid | See appendix: Acetic acid
Do NOT taste these acids in the laboratory. Each acid has a sour taste that is a characteristic of acids.
Lemon juice contains the white crystalline citric acid. Vinegar contains ethanoic acid (acetic acid, CH3COOH).
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 and citric acid. Do not taste any other acids because they may damage living tissues.
12.3.2 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. 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. Observe the properties of any gas liberated and name it. Test it 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 gas evolves by noting the rate and size of the hydrogen gas 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.
2. 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 gas. 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 gas then zinc then iron. Tin forms few bubbles of hydrogen gas. Copper forms no bubbles of hydrogen gas. Lead forms some lead chloride precipitate on the surface of the lead. Aluminium develops a layer of aluminium oxide that obstructs further chemical reactions.
3. 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.
4. 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 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 gas 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 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 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 gas 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.
12.3.3.2 Magnesium with sodium hydrogen sulfate
Add 3 cm of magnesium ribbon to 3 cm of sodium hydrogen sulfate solution in a test-tube. The metal reacts with the sulfuric acid in the solution. Describe what you see. Tests for hydrogen gas. Remove any magnesium which has not reacted from the solution, pour part of the liquid into an evaporating basin and leave for magnesium sulfate crystals to form.
12.3.3.3 Iron with sodium hydrogen sulfate
Add a finger width of iron filings to a finger width of sodium hydrogen sulfate solution in a test-tube. Heat the mixture to speed up the reaction. The metal reacts with the sulfuric acid in the solution. Describe what you see. Tests for hydrogen gas. Leave to stand until all bubbles have ceased to appear. Pour part of the liquid into an evaporating basin and leave for magnesium sulfate crystals to form.
Test the liquid with universal indicator paper (The indicator changes colour to red, orange, or yellow for acids and green, or violet for alkalis.) Pale green is the colour for neutral substances. Before testing, make the paper this colour by dipping it into neutral tap water for a few moments. The Universal indicator turns yellow indicating the presence of an acid. Filter the liquid, and pour part of the clear solution into the evaporating basin and leave for pale green crystals of iron sulfate to form (FeSO4.7H2O, green vitriol). A solution of a salt is not necessarily neutral because some salts, like iron sulfate, form acids when dissolved in water.
12.3.4 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 Dilute acids with basic oxides, metal 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 from 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)
12.3.5.01 Copper sulfate from copper oxide
Add half a test-tube of sodium hydrogen sulfate solution to copper oxide in a test-tube. Heat the solution slowly until it turns blue. Be careful of spurting from the test-tube. Some copper oxide may remain after the reaction. Filter the solution obtain the filtrate of copper sulfate solution.
12.3.5.1 Dilute acids with basic oxides, (metal 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 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 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 Dilute acids with sodium hydroxide
Repeat the previous experiment with: dilute sulfuric acid, dilute nitric acid, ethanoic acid (acetic acid).
HCl (aq) + NaOH (aq) --> NaCl (aq) + H2O (l)
hydrochloric acid + sodium hydroxide --> sodium chloride + water
12.3.7.2 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 Dilute acids with acidic oxides (non-metal 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 Dilute acids with carbonates, 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 Dilute hydrochloric acid with calcium carbonate
See diagram 3.34.1: Limewater tests 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 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.
12.3.9.3 Dilute tartaric acid with sodium carbonate
Put 5 g of sodium carbonate and the same quantity of tartaric acid in a test-tube fitted with a one-hole stopper and delivery tube With the end of the delivery tube dipping into a second test-tube of limewater add water to the first test-tube and quickly replace the stopper. Describe what you see. The limewater turns milky. The acid reacts with sodium carbonate to form a salt, carbon dioxide, and water.
tartaric acid + sodium carbonate --> sodium tartrate + carbon dioxide + water
12.3.9.4 Dilute tartaric acid with egg shell, soil, wood ash
Many common substances, such as mortar, egg shell, most soils, contain calcium carbonate. and wood ashes contain potassium carbonate. Observe the action of tartaric acid on these substances in a test-tube. Tests for carbon dioxide by holding a drop of limewater, at the end of a glass tube, in the mouth of the test-tube.
12.3.9.5 Dilute hydrochloric acid with calcium carbonate
Put 5 g of marble chips (calcium carbonate) and the same quantity of dilute hydrochloric acid in a test-tube fitted with a one-hole stopper and delivery tube With the end of the delivery tube dipping into a second test-tube of limewater add water to the first test-tube and quickly replace the stopper. Describe what you see. The limewater turns milky. The acid reacts with calcium carbonate to form a salt, carbon dioxide, and water.
hydrochloric acid + calcium carbonate --> calcium chloride + carbon dioxide + water
12.3.9.6 Dilute sulfuric acid with calcium carbonate
Put 5 g of marble chips (calcium carbonate) and the same quantity of dilute sulfuric acid in a test-tube fitted with a one-hole stopper and delivery tube With the end of the delivery tube dipping into a second test-tube of limewater add water to the first test-tube and quickly replace the stopper. Describe what you see. The limewater turns milky. The acid reacts with calcium carbonate to form a salt, carbon dioxide, and water. The reaction of sulfuric acid with calcium carbonate proceeds only for a few moments because the salt formed, calcium sulfate, is only slightly soluble and deposits on the carbonate, preventing this compound from reacting with the acid. So the reaction with hydrochloric acid above is much better.
sulfuric acid + calcium carbonate --> calcium sulfate + carbon dioxide + water
12.3.10 Dilute acids with sodium hydrogen carbonate
The only stable hydrogen carbonates are KHCO3 and 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 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.
12.3.11.0 Dilute nitric acid with copper
Very dilute nitric acid may react with very active metals, e.g. magnesium to form hydrogen gas. 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 gas 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 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 gas 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 Concentrated acids with metals, nitric acid with copper
Nitric acid reacts with metals above platinum in the activity series, but does not form hydrogen gas.
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 Concentrated acids with metals, 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 gas.
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 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 Acids with salt
1. 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.
2. Repeat the procedure using dilute sulfuric acid and then concentrated sulfuric acid.
3. 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. 4. Acids react with sulfites to produce sulfur dioxide, water and a salt.
5. Acids react with sulfides to produce hydrogen sulfide (rotten egg gas) and a salt.
6. Concentrated sulfuric acid reacts with chlorides to produce hydrogen chloride and a sulfate.
7. Concentrated sulfuric acid reacts with nitrates to produce nitric acid and a sulfate.
8. 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 20oC 3. 36% "ANALAR" solution (d) commercial solution called muriatic acid for use in the building trades.
12.5.0 Nitric acid
See 3: Concentrated sulfuric acid with copper | See 12.3.5: 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 + 2HNO2
Heated nitrous acid decomposes to form nitric acid and nitrogen monoxide (nitric oxide).
2HNO2 --> HNO3 + NO
Ionization of nitrous acid, Ka = 4.5 X 10-4
HNO2 + H2O <--> H3O+ + NO2-
Nitrosamines, produced by nitrous acid with secondary amines, can be formed in the gut when nitrites react with amino acids.
12.5.1 Prepare nitric acid, sulfuric acid with 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)
Ionization of hydrogen sulfate ion
HSO4- + H2O <--> H3O+ + SO42-
12.6.0 Sulfuric acid, ionization of hydrogen sulfate ion, sulfur trioxide
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 20oC (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 450oC to form sulfur trioxide that combines with water to form sulfuric acid.
Sulfur trioxide is produced by the action of oxygen on sulfur dioxide in the presence of a catalyst, e.g. iron oxide.
12.6.0.1 Formation of acid rain, SOx, by burning sulfur or sulfur compounds
When coal is burnt, the compounds that contain sulfur can form sulfuric acid, as in the equations below, to become components of acid rain (rainwater pH = 5.6, acid rain pH < 5). There may be more than one pathway for the formation of sulfuric acid from sulfur dioxide.
12.6.0.2 The 1 - 3% "sulfur" in coal is usually iron (II) sulfide (pyrite).
4 FeS2 (s) + 11O2 (g) --> 2Fe2O3 (s) + 8SO2 (g)
S (s) + O2 (g) --> SO2 (g) sulfur dioxide
Also, other sulfide ores may produce sulfur dioxide in the atmosphere during smelting to obtain the pure metal.
Lead (II) sulfide (galena)
3PbS + 3O2 (g) --> 3Pb (s) + 3SO2 (g)
Copper (II) sulfide, chalcocite (bornite)
Cu2S (s) + O2 (g) --> 2Cu (s) + SO2 (g)
12.6.0.3.1 Sulfur dioxide to sulfuric acid 1.
Sulfur dioxide is oxidized to sulfur trioxide by oxygen gas.
2SO2 (g) + O2 (g) <--> 2SO3 (g)
Sulfur dioxide is oxidized to sulfur trioxide by nitrogen dioxide.
SO2 (g) + NO2 (g) --> SO3 (g) + NO (g)
Sulfur trioxide dissolves in water to form sulfuric acid.
SO3 (g) + H2O (l) --> H2SO4 (aq) sulfuric acid
12.6.0.3.2 Sulfur dioxide to sulfuric acid 2.
Sulfur dioxide dissolves in water to form sulfurous acid
SO3 (g) + H2O (l) --> H2SO3 (aq)
Sulfurous acid is oxidized to to sulfuric acid by ozone
H2SO3 (aq) + O3 (g) --> H2SO4 (aq) + O2 (g)
12.6.0.4 Hydrogen peroxide oxidizes sulfur dioxide to sulfuric acid
SO2 (g) + H2O2 (l) --> H2SO4 (aq)
12.6.1 Sulfuric acid acts as an oxidizing agent
See 3: 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
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)
1. 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)
2. 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. BE CAREFUL! 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 slowly ADD ACID TO WATER. Never add water to 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.
12.9.0 Phosphoric acid
Ionization reaction
H3PO4 + H2O <--> H3O+ + H2PO4-
H2PO4- + H2O <--> H3O+ + HPO42-
HPO42-+ H2O <--> H3O+ + PO43-
12.10.0 Boric acid, ionization reaction, prepare boric acid crystals
1. Orthoboric acid, trioxoboric acid (III) acid, boracic acid, sassolite, H3BO3 is a weak acid. White to colourless triclinic crystals, m.p. 169oC, 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-
H3BO3 <--> H+ + H2BO3-
H2BO3- <--> H+ + HBO32-
HBO32- <--> H+ + BO32-
2. Prepare boric acid crystals
Use 5 g of boric acid crystals. Pour some into 2 cm boiling water in a test-tube and leave to dissolve. Continue adding crystals and heat to boiling until all crystals dissolved. Leave to cool to see fine white crystals form.