School Science Lessons
(topic12E)
2024-09-17
Tests, Tests for all substances, gases and vapours, metals, Titration
Contents
12.11.1 Tests for all substances
Tests for acetates: 12.11.1
Tests for acetic acid in vinegar: 9.3.1
Tests for acetylene (ethyne): 16.4.6.2
Tests for acid radicals in solution: 12.11.2
Tests for air and dissolved oxygen in water: 18.7.8
Tests for air pollution from burning refuse: 18.6.5
Tests for albumin and gelatine: 9.3.2
Tests for aldehydes, Tollens' test: 9.3.3
Tests for aldehydes with Fehling's solution: 9.5.2
Tests for alcohol, breath tests: 15.2.11
Tests for aluminium: 12.3.12
Tests for aluminium compounds: 12.3.13
Tests for aluminium compounds in solution: 12.3.14
Tests for ammonia: 3.5.0
Tests for amylose and amylopectin: 9.3.4
Tests for anions: 12.11.0
Tests for antimonates, borates, oxalates: 12.11.2a
Tests for antimony: 12.11.3.12
Tests for arsenates: 12.11.3
Tests for arsenic, flame tests: 12.11.3.8 (See: 2. Arsenic)
Tests for ash content of plant dry matter: 9.3.5
Tests for aspirin: 5.5.6.2
Tests for barium: 12.11.3.13
Tests for benzidine: 12.11.3.10
Tests for bicarbonates: 12.11.4
Tests for bismuth: 12.11.3.14
Tests for bismuth: 12.11.3.8, (See: 3.)
Tests for blood: 9.9.1
Tests for borates: 12.11.5
Tests for borax / turmeric adulteration of food: 19.1.32
Tests for breakdown of starch to sugars: 9.3.6
Tests for bromides: 12.11.6
Tests for cadmium: 12.11.3.8, (See: Heat with charcoal 4.)
Tests for cadmium: 12.11.3.15
Tests for calcium: 16.5.5
Tests for calcium, flame test: 12.11.3.16
Tests for calcium carbonate (limestone): 35.6.14 (Geology)
Tests for carbohydrates, Molisch's test: 9.3.7
Tests for carbon dioxide: 3.5.0
Tests for carbon dioxide in the breath with limewater: 9.1.18
Tests for carbon monoxide with a gas detector: 18.6.6
Tests for carbonates: 12.11.7
Tests for cellulose: 9.3.8
Tests for cement brick strength (contents): 34.2.3
Tests for cement brick strength (water content): 34.2.4
Tests for cement change in weight when setting: 34.2.5
Tests for chewing gum quality by comparing bubbles: 3.4.2
Tests for chlorine: 12.11.7
Tests for chromates: 12.11.9
Tests for chromium: 12.11.3.17
Tests for cigarette smoke: 16.5.6
Tests for cobalt: 12.11.3.18
Tests for concrete alkalinity: 34.2.6
Tests for copper: 12.4.0
Tests for dextrins in toast: 19.2.14
Tests for diastase activity: 9.3.9
Tests for dinitrogen oxide, nitrous oxide, N2O: 13.3.23
Tests for dissolved oxygen, DO (Winkler method): 18.3.2
Tests for fabrics, Burning tests for fabrics: 4.0.0
Tests for Faraday's first law: 15.1.1.1
Tests for fats and oils: 9.3.11
Tests for food, food tests: 9.3.0
Tests for gases and vapours: 12.11.2
Tests for gases collected in a respirometer: 9.1.19
Tests for gases from burning hydrocarbons: 16.4.6.0
Tests for gases with hot concentrated sulfuric acid:12.11.3.6
Tests for gases and vapours: 12.11.2
Tests for glucose: 12.11.5
Tests for glycerol: 12.2.11
Tests for glycerine: 12.7.4
Tests for gypsum added to the soil: 9.15.5
Tests for haemoglobin, Hematrace test: 9.1.14
Tests for halides, Cl, Br, I: 12.11.10
Tests for hard water. water hardness: 12.4.0
Tests for hardness, lead, tin, and tin alloys: 3.62.0
Tests for heat: 12.11.6
Tests for hydrogen chloride: 13.2.24
Tests for hydrogen chloride with ammonia solution: 3.3.8
Tests for hydrogen gas: 13.3.4.4
Tests for hydrogen peroxide: 17.7.23
Tests for hydrogen sulfide solution: 13.3.26
Tests for hydrolysis of starch, dilute hydrochloric acid: 9.5.3
Tests for hydrolysis of starch, salivary amylase: 9.5.4
Tests for hydroxides: 12.11.11
Tests for hydroxyl ions, ammonia solution: 3.3.7
Tests for insoluble solids in rain water: 18.2.1
Tests for iodides: 12.11.12
Tests for ions in a water sample: 18.4.0
Tests for iron: 12.11.3.20
Tests for iron in cooking water: 19.2.15
Tests for ketones: 19.5.1
Tests for lactic acid solution: 12.7.11
Tests for lead: 12.11.3.21
Tests for lemon juice effect on apple browning: 19.2.3.4 (Cooking)
Tests for lignin: 9.3.12
Tests for limestone: 35.6.14 (Geology)
Tests for lipase: 9.3.13
Tests for magnesium
Tests for manganese: 12.11.3.23
Tests for melting point of lead, tin, and lead-tin alloys: 3.63
Tests for metals: 12.11.3
Tests for methane gas, burn methane: 16.5.1.2
Tests for milk: 16.0.0
Tests for moisture content of plant organs: 9.3.14
Tests for multiple reagent strips: 19.5.6, Tests with
Tests for natural fabrics, burning tests: 4.2.0
Tests for nickel, DMG: 12.11.3.24
Tests for nitrates / nitrites with dipsticks: 19.5.2
Tests for nitrates: 12.11.13, the brown ring test
Tests for nitrogen content in food, Kjeldahl method: 16.5.7
Tests for nitrogen content in food, soda lime test: 9.3.15
Tests for nitrous oxide, (dinitrogen oxide), N2O: 13.3.23
Tests for organic acids and alcohols: 9.3.16
Tests for organic functional groups
Tests for oxalates: 12.11.14
Tests for oxidase and peroxidase in plant tissues: 9.3.14a
Tests for oxidation of glucose, blue bottle experiment: 9.3.10a
Tests for oxidizing agents: 15.4.17
Tests for oxygen absorption during plant respiration: 9.1.20
Tests for oxygen content of water, (dissolved oxygen): 18.3.2
Tests for oxygen gas: 3.49.1
Tests for pectin in jelly and jam: 16.5.8
Tests for pH with acid-base indicators: 5.6.1
Tests for pH of soil samples:../topics/topic18.html#18.7.11H">18.7.11
Tests for phosphate ions in water: 18.4.1
Tests for phosphates: 12.11.15
Tests for plant tissues water content: 9.3.14
Tests for plant tissues oxidase and peroxidase: 9.3.14aH
Tests for plastics, Burning tests for fabrics: 4.0.0
Tests for polymers: 4.0.0
Tests for potassium, sodium perchlorate: 12.11.3.25
Tests for potassium, tetraphenylborate test: 12.11.3.1.1
Tests for proteins: 9.4.0
Tests for rain water (soluble solids): 18.2.3
Tests for reagents, multiple reagent strips: 19.5.6
Tests for reducing sugars, Benedict's test: 9.4.1
Tests for reducing sugars, Fehling's test: 9.5.0
Tests for respiration of soaked peas with limewater: 9.1.21
Tests for salt effect on buffer solutions: 12.12.11
Tests for salts with flame tests: 12.11.3.7
Tests for salts with flame test sprays: 12.11.3.9
Tests for saturated hydrocarbons, bromine water test: 9.3.33.6
Tests for silver: 12.11.3.8 (See 8.)
Tests for silver, potassium chromate: 12.11.3.26
Tests for soap: 12.2.12
Tests for sodium: 12.11.3.27
Tests for sodium bicarbonate in a stomach powder: 12.1.20
Tests for sodium chloride, flame test: 12.1.29
Tests for soils, Soil tests: 6.10.0
Tests for solubility: 12.11.3.3
Tests for soluble solids in rain water: 18.2.3
Tests for starch: 12.11.8
Tests for strength of mud, clay and sand bricks: 34.2.10
Tests for strength of plaster of Paris bricks: 34.2.11
Tests for strontium: 12.11.3.28
Tests for substances by loss on heating: 12.11.3.1
Tests for substances by sublimation, melting, decrepitation: 12.11.3.4
Tests for substances with dilute hydrochloric acid: 12.11.3.5
Tests for substances with heated charcoal and fusion mixture: 12.11.3.8
Tests for sugars: 9.3.17
Tests for sulfates: 12.11.16
Tests for sulfates in groundwater: 18.2.2.3
Tests for sulfides: 12.11.17
Tests for sulfites: 19.5.3
Tests for sulfur dioxide: 13.3.0
Tests for sulfur in proteins: 9.3.18
Tests for swimming pools: 18.5.0
Tests for synthetic fibres, burning tests: 4.3.0
Tests for tannic acid in tea: 9.3.19
Tests for tartaric acid: 19.5.4
Tests for thermometer calibration: 22.7.3
Tests for tin: 12.11.3.8 (See: Heat with charcoal 9.)
Tests for tin: 12.11.3.29
Tests for toxic effect of drugs on water flea: 16.5.9
Tests for turbidity: 18.7.15
Tests for unsaturation: 9.3.20
Tests for urine: 19.5.5
Tests for vitamin C, DCPIP: 9.3.21
Tests for water: 12.11.9
Tests for wheat starch for gluten: 16.6.2
Tests for wood: 9.3.22
Tests for zinc: 12.11.3.8, (See: Heat with charcoal 10.)
Tests for zinc: 12.11.3.30
Tests for zymase and catalase in yeast: 9.3.23
9.3.0 Tests for food
9.3.1 Tests for acetic acid in vinegar
9.3.2 Tests for albumin and gelatine
9.3.3 Tests for aldehydes, Tollens' test
9.5.2 Tests for aldehydes with Fehling's solution:
9.3.4 Tests for amylose and amylopectin
9.3.5 Tests for ash content of plant dry matter.
9.3.6 Tests for breakdown of starch to sugars
9.3.7 Tests for carbohydrates, Molisch's test
9.3.8 Tests for cellulose
9.3.9 Tests for diastase activity
12.11.5 Tests for glucose
9.3.11 Tests for fats and oils
9.3.12 Tests for lignin
9.3.13 Tests for lipase
9.3.14 Tests for plant tissues, moisture content
9.3.15 Tests for nitrogen content in food, soda lime test
9.3.16 Tests for organic acids and alcohols
9.4.0 Tests for proteins
9.5.0 Tests for starch
9.3.17 Tests for sugars, sugar test solution
9.3.18 Tests for sulfur in proteins
9.3.19 Tests for tannic acid in tea
9.3.20 Tests for unsaturation
9.3.21 Tests for vitamin C, DCPIP
9.3.22 Tests for wood
9.3.23 Tests for zymase and catalase in yeast.
12.11.2 Tests for gases and vapours
Tests for acetylene, (ethyne): 16.1.3.2
Tests for ammonia: 3.5.0
Tests for carbon dioxide: 3.34.1.1
Tests for carbon dioxide in the breath with limewater : 9.1.18
Tests for carbon monoxide with a gas detector: 18.6.6
Tests for chlorine: 3.40.1
Tests for chlorine levels in swimming pools, test kit: 18.7.21.0
Tests for gases collected in a respirometer: 9.1.19
Tests for gases from burning hydrocarbons: 16.4.6.0
Tests for gases with hot concentrated sulfuric acid: 12.11.3.6
Tests for gases: Lighted splint tests
Tests for hydrogen chloride: 3.42.1.0
Tests for hydrogen gas: 3.41.1.0
Tests for hydrogen sulfide solution: 3.43.1
Tests for methane gas, burn methane: 16.5.1.2
Tests for nitrous oxide, dinitrogen oxide, N2O: 3.45.1
Tests for oxygen absorption during plant respiration: 9.6.7
Tests for oxygen gas: 3.49.1
Tests for sulfur dioxide: 3.51.6
12.11.3 Tests for metals
Tests for metal ions in water, EDTA chelates: 12.13.11
Tests for metals with borax beads: 12.11.3.1a
Tests for metals with flame tests: 12.11.3.2
Tests for aluminium: 12.11.3.11
Tests for antimony: 12.11.3.12
Tests for barium: 12.11.3.13
Tests for bismuth: 12.11.3.14
Tests for cadmium: 12.11.3.15
Tests for calcium: 12.11.3.16
Tests for chromium: 12.11.3.17
Tests for cobalt: < a href="#12.11.3.18H">12.11.3.18
Tests for copper: 12.4.0
Tests for iron: < a href="#12.11.3.20H">12.11.3.20
Tests for lead: < a href="#12.11.3.21H">12.11.3.21
Tests for lead, heat charcoal with mixture: < a href="#12.11.3.8H">12.11.3.8
Tests for lead ions: < a href="../topics/topicIndexL.html#15.8.3H">15.8.3
Tests for magnesium: < a href="#12.11.3.22H">12.11.3.22
Tests for manganese: < a href="#12.11.3.23H">12.11.3.23
Tests for nickel: < a href="#12.11.3.24H">12.11.3.24
Tests for potassium: < a href="#12.11.3.25H">12.11.3.25
Tests for silver: < a href="#12.11.3.26H">12.11.3.26
Tests for silver: 12.12.11.3.8, (See 8.)
Tests for sodium: < a href="#12.11.3.27H">12.11.3.27
Tests for strontium: < a href="#12.11.3.28H">12.11.3.28
Tests for tin: < a href="#12.11.3.29H">12.11.3.29
Tests for zinc: < a href="#12.11.3.30H">12.11.3.30c
12.11.5 Tests for glucose
Tests for glucose: 19.5.8
Tests for glucose and fructose with Fehling's reagent: 9.5.4
Tests for glucose and starch, "Testape": 9.5.2
Tests for glucose, blood glucose, glucose tolerance: 19.1.26
Tests for glucose, Clinitest tablet: 19.1.24
Tests for glucose, Clinistix strip: 19.1.25
Tests for glucose concentration, ferricyanide test: 19.1.27
Tests for glucose concentration, ferricyanide test: 19.1.26
Tests for glucose, blood glucose: 19.1.26
Tests for glucose, glucose oxidase test: 19.1.29
Tests for glucose, glycosylated haemoglobin test: 19.1.30
Tests for glucose, Nelson-Somogyi test: 19.1.28
Tests for glucose: 19.5.8
Tests for glucose oxidation, blue bottle experiment: 9.3.10
Tests for glucose, Tollens' test: 9.3.10
Tests for glucose, urine test: 19.1.31
Tests for glucose, urine test, Diastix strip: 19.1.33
12.11.6 Tests for heat
Heat of reaction, chromium (VI) oxide with ethanol: 14.1.9
Heat of reaction, potassium permanganate with ethanol: 14.1.8
Heat of reaction, potassium permanganate with glycerol: 14.1.7
Heat of reaction, potassium with diethyl ether: 14.1.10
12.11.7 Tests for chlorine
Tests for chlorides: 12.11.8
Tests for chlorides in groundwater: 18.7.2
Tests for chlorine 1: 15.5.12a
Tests for chlorine 2: 3.40.1
Tests for chlorine 3: 18.7.21.0, (Test kit for swimming pools)
Tests for chlorine 4: 18.7.4.0, (Tests for free swimming pools)
12.11.8 Tests for starch
9.3.6 Tests for breakdown of starch to sugars
9.5.3 Tests for hydrolysis of starch, dilute hydrochloric acid
9.5.4 Tests for hydrolysis of starch, salivary amylase
9.5.5 Tests for starch in potato tuber cells
9.5.6 Tests for starch, iodine tests
9.5.2 Tests for glucose and starch, "Testape"
1.6 Tests for starch, iodine solution
9.5.3 Tests for starch with Fehling's solution
12.19.6.10 Tests for starch, Iodine with starch
16.1.15 Starches, amylum, glycogen, iodine test for starch
12.7.14 Tests for starch in adhesive paste
12.1.42 Starch with water, iodine test
12.11.9 Tests for water
Tests for air and dissolved oxygen in water: 18.7.8
Tests for anions in sewage and tap water: 18.5.2
Tests for colour of water: 18.7.9
Tests for colour of water: 18.6.6.1
z
Tests for contamination of groundwater from refuse deposits: 18.2.4
Tests for dissolved oxygen: 18.3.3, titration
Tests for electrical conductivity with a conductivity meter: 18.2.6.1
Tests for environment of river, lake or ocean: 18.5.1
Tests for hydrogen ion concentration of water: 18.7.10
Tests for insoluble solids in rain water: 18.2.1
Tests for ions in a water sample: 18.4.0
Tests for moisture content of plant organs: 9.3.14
Tests for oxygen content of water: 18.3.2, dissolved oxygen, DO, (Winkler method)
Tests for pH of soil samples.18.7.11
Tests for pH of rain water: 18.1.3
Tests for pH of standing water: 18.1.4
Tests for pH of water in the laboratory: 18.1.2
Tests for pH with universal indicator: 18.1.1
Tests for phosphate ions in water18.4.1
Tests for smell of water, hydrogen sulfide: 18.7.12
Tests for soluble solids in rain water18.2.3
Tests for sulfates in groundwater18.2.2.3
Tests for temperature of water: 18.7.13
Tests for total dissolved solids and suspended solids in water: 18.2.0
Tests for total solids: 18.7.14 Tests for total solids
Tests for turbidity: 18.7.15
Tests for water hardness: 12.4.0
Tests for water pollution: 18.7.0
Tests for water, pH tests: 18.1.0
Tests for water samples: 18.5.3
Tests for water with anhydrous copper(II) sulfate: 8.2.1
Tests for water with cobalt (II) chloride paper: 8.2.2
Using rivers for water testing: 18.7.17
12.11.4 Titration, acid-base neutralization
12.11.4.1 Acid-base neutralization, acid with base forms a salt and water
12.11.4.2 Acidity of vinegar and wine
12.11.4.3 Ammonia with sulfuric acid
12.11.4.4 Carbon dioxide affects acid-base titration
12.11.4.5 Group analysis
12.11.4.6 Heat of neutralization titration
12.11.4.7 Microscale titration, sodium hydroxide with dilute acids
12.11.4.8 Prepare monoprotic acid solution from unknown molarity acid
12.11.4.9 Simple titration of acids and bases
12.11.4.10 Sodium hydroxide with hydrochloric acid
12.11.4.11 Titrate dilute hydrochloric acid with sodium hydroxide solution, with a burette
12.11.4.12 Titrate dilute sulfuric acid with sodium hydroxide solution and isolate sodium sulfate crystals.
12.7.14 Tests for starch in adhesive paste
Shake a drop of the paste with water in a test-tube and add a drop of iodine solution.
If the paste contains starch the contents of the test-tube turns blue-black.
If a red colour forms the paste contains a chemical called dextrin, that is made from starch.
12.1.42 Starch with water, iodine test
Shake a small pinch of powdered starch with half a test-tube of water.
The starch does not dissolve.
Tests the liquid by adding one drop of iodine solution.
No reaction is given.
Boil a small pinch of starch with half a test-tube of water for seconds.
The starch dissolves.
Cool the test-tube under the tap and add one drop of iodine solution.
A deep blue-black liquid forms.
12.11.3.1 Tests for substances by loss on heating
1. Loss of water vapour indicates the presence of water of crystallization and some basic hydroxides, basic carbonates, and acid salts.
Also, it may indicate that the substance has absorbed moisture from the atmosphere.
2. Loss of oxygen indicates oxides of silver, peroxides, sodium or potassium nitrate, permanganates and chlorates.
3. Loss of carbon dioxide indicates carbonate or bicarbonate.
4. Loss of ammonia indicates an ammonium compound.
5. Loss of nitrogen dioxide (dinitrogen tetroxide), (N2O4), indicates nitrates of heavy metals, e.g. copper, lead, zinc.
6. Loss of sulfur trioxide indicates some sulfates.
7. Loss of halogens indicates oxidation of a halide.
12.11.3.1.1 Tests for potassium, tetraphenylborate test
To 15 drops of the solution, add 5 drops of 1 M NaOH, boil the solution, add 2 drops of 6 M HCl and 15 drops of 1 M sodium acetate.
Add 3 drops of 3% sodium tetraphenylborate, NaB(C6H5)4.
If a white precipitate forms the test is positive.
12.11.3.2 Tests for metals with flame tests
Cations in an unknown solution can be identified by using flame tests.
Add drops of concentrated hydrochloric acid to the solution.
Dip a clean piece of platinum wire into it then hold it in a Bunsen burner flame.
Dip platinum wire into concentrated hydrochloric acid (12 M) then into powdered solid and heat in a non-luminous edge of a Bunsen burner flame.
When a salt is heated in the flame, it dissociates into neutral atoms and electrons are excited into a higher energy level then return to the ground state and emits light of characteristic colour for that atom.
Remember that each person observes colours differently.
The non-metal atoms in anions do emit light, but at wavelengths shorter than ultraviolet light, so we cannot see the colours.
Experiments
Check the flame test colours by doing the test for all the cations against a dark background.
Compare with the following list of colours:
Ammonium compounds: green (faint colour).
Antimony: blue-green to light blue (faint colour)
Arsenic: light blue (moistened with hydrochloric acid)
Barium: pale green to yellow-green:
Bismuth: blue
Calcium: red
Calcium compounds: brick-red to yellow (masked by barium)
Copper: blue-green
Copper compounds: green (not halides) (CuBr2 blue-green)
Lead: light blue to blue
Lithium compounds: crimson (masked by barium or sodium)
Molybdenum: yellow-green:
Phosphates: blue-green (if when moistened with sulfuric acid).
Potassium: lilac, but crimson through blue glass, violet through cobalt glass
Potassium compounds: pink-lilac to violet (not borates, phosphates, and silicates.) (masked by sodium or lithium).
Selenium: blue
Sodium: strong golden yellow, but no colour viewed through blue glass
Sodium compounds: yellow (if the yellow flame persists and is not intensified by adding 1% NaCl to the dry compound.)
Strontium: crimson
Strontium compounds: scarlet (masked by barium)
Zinc: green-white.
12.11.3.3 Tests for solubility, prepare a solution for group analysis
Dissolve 1 g of the substance in the first reagent below that can dissolve the substance.
1. Water: Try to dissolve the salt in deionized water.
If the salt does not dissolve, heat it in a test-tube to observe if the salt dissolves in hot water.
2. Dilute hydrochloric acid: If the salt does not dissolve in hot water, add dilute hydrochloric acid to observe if it dissolves.
3. Concentrated hydrochloric acid, 2.0 to 5.0 mL: When all substance is dissolved, dilute solution to five times its bulk, then leave to cool.
If the dilution produces a precipitate, because of hydrolysis of chlorides of bismuth, antimony or tin, add drops of concentrated hydrochloric acid.
4. Dilute nitric acid: It may dissolve compounds of lead, silver and mercury, but avoid using nitric acid, because it may oxidize hydrogen sulfide.
5. Aqua regia, Be careful!
Heat with concentrated hydrochloric acid, then add a few drops of concentrated nitric acid.
Dilute as in 3. above.
6. Concentrated nitric acid: Warm with 2.0 to 5.0 mL of the acid.
The solution must be evaporated to dryness and the residue dissolved in water or hydrochloric acid.
7. Insoluble residue: Filter it off, wash it, and fuse it in a crucible with four times its bulk of fusion mixture.
Cool, boil with water, filter.
Test filtrate for acid radicals.
Dissolve precipitate of metal carbonates in hydrochloric acid, and analyse separately.
If the solution so obtained gives no precipitate with that obtained by 1. to 6. above, analyse them together.
12.11.3.4 Tests for substances by sublimation, melting, decrepitation
1. Sublimation indicates the presence of ammonium halides, other halides, and some oxides.
2. Melting indicates the presence of sodium, potassium or ammonium nitrate, potassium chlorate, and other less common substances.
3. Decrepitation (crackling noise of some heated crystals), indicates the presence of sodium chloride, lead nitrate and potassium chlorate.
12.11.3.5 Tests for substances with dilute hydrochloric acid
1. Carbon dioxide produced indicates a carbonate or bicarbonate.
2. Hydrogen gas produced indicates some free metals.
3. Sulfur dioxide produced indicates sulfite or bisulfite.
4. Sulfur dioxide and sulfur produced indicates thiosulfate.
5. Hydrogen sulfide produced indicates sulfide.
6. Nitrogen dioxide (N2O4) produced indicates nitrite.
7. Chlorine produced indicates hypochlorite or oxidizing agent.
12.11.3.6 Tests for gases with hot concentrated sulfuric acid
If organic acid present, substance should be ignited, extracted with dilute hydrochloric acid and filtered before proceeding with main group separation.
Gas evolved indication
1. Hydrogen chloride produced indicates chloride.
2. Nitric acid produced indicates nitrate.
3. Oxygen produced indicates peroxide, permanganate, chromate, dichromate.
4. Chlorine peroxide (yellow green gas, violent action) produced indicates chlorate.
5. Sulfur dioxide produced indicates sulfite, thiosulfate or reducing agent.
6. Hydrogen bromide, bromine and sulfur produced indicates bromide.
7. Hydrogen iodide, iodine and hydrogen sulfide produced indicates iodide.
8. Carbon monoxide and carbon dioxide produced indicates oxalate.
9. Carbon monoxide only produced indicates formate.
10. Acetic acid produced indicates acetate.
12.11.3.7 Tests for salts with flame tests
Soak paper in the following salts, leave to dry then ignite
Calcium chloride: orange, Copper (II) chloride: blue, Copper (II) sulfate: green, Lithium chloride: red, Potassium chloride: purple, Sodium borate, borax: green.
Sodium carbonate: yellow, Sodium chloride: yellow, Strontium chloride: red.
12.1.29 Sodium chloride flame tests
12.11.3.8 Test substances with heated charcoal and fusion mixture
Test substances with charcoal, heat charcoal with fusion mixture, note heated metal appearance.
1. Aluminium produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a blue mass, but this is also caused by fusible phosphates, arsenates, borates and silicates.
2. Arsenic produces fumes smelling of garlic and forms a white crust if seen at some distance from the flame.
3. Bismuth forms pink globules, becomes brittle and forms a yellow crust.
4. Cadmium forms a brown crust.
5. Copper forms red scales.
6. Lead forms grey-white soft globules and forms a red crust when hot and a yellow crust when cold.
7. Magnesium produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a green mass.
8. Silver has shining metal particles.
9. Tin forms hard white beads.
10. Zinc forms a yellow crust when hot and a white crust when cold.
Zinc also produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a green mass.
12.11.3.9 Tests for salts with flame test sprays
Be careful! Wear eye protection.
Use spray bottles, e.g. window cleaners or garden sprays, to spray saturated solutions of metal salts in ethanol onto roaring Bunsen burner flames in a darkened room.
The salts can include sodium chloride, potassium chloride, lithium chloride, and copper sulfate.
The spray bottles should have a trigger mechanism and not a scent bottle spray pump, which may allow flash back.
12.11.3.10 Benzidine
Organic reagent, C12H12N2, 0.05% solution in 10% acetic acid
To one drop of solution on filter paper add one drop 0.05% NaOH then one drop of reagent.
Use in Group IV when in solution in dilute acid.
Dissolve Group IV precipitate in very dilute acid and use the solution, rejecting any undissolved solid.
12.11.3.11 Tests for aluminium
1. Heat charcoal with fusion mixture, note heated metal appearance.
Aluminium produces a white residue.
Add drops of cobalt nitrate solution and heat again to form a blue mass, but this is also caused by fusible phosphates, arsenates, borates and silicates.
2. Test aluminium
Test aluminium after dipping it in concentrated hydrochloric acid then press it on absorbent paper to remove the layer of aluminium oxide.
The voltage reading will start at a low value then increase as remaining aluminium oxide dissolves.
Record the maximum value.
12.11.3.12 Tests for antimony
1. Dilute with its own volume of water.
Pass H2S.
An orange-red precipitate of antimony sulfide, Sb2S3, indicates the presence of antimony.
2Sb3+ + 3S2- --> Sb2S3 (s).
2. Organic reagent: Gallocyanine (Fast violet), C15H13ClN2O5, 0.05% in M HCl.
To one drop of antimony solution on filter paper, add one drop of reagent.
A colour change from wine red to blue indicates the presence of antimony.
Use Group IIb precipitate dissolved in concentrated HCl and diluted.
12.11.3.13 Tests for barium
Confirm by flame test: Light green
Tests for Barium and strontium
Organic reagent: Rhodizonic acid, C6H2O6, C6H2O6.2H2O, 0.1% aqueous solution
Put one drop of test liquid on filter paper then add one drop of reagent.
A red-brown spot indicates the presence of Sr and Ba.
When one drop of dilute HCl is added, a barium spot is intensified and a Sr spot disappears.
Prepare fresh solution of reagent if it has decolorized.
12.11.3.14 Tests for bismuth
Organic reagent: Thiourea, H2N.CS.NH2, 10% aqueous solution (10 mL Bi solution + 10 mL dilute HNO3 + 1 mL reagent)
A yellow colour indicates the presence of bismuth.
12.11.3.15 Tests for cadmium
1. Ammonium hydroxide gives white precipitate easily soluble in excess.
2. Organic reagent: Diphenyl carbazide, C13H14N4O, CO(NH.NH.C6H5)2, in saturated alcoholic solution
Add few drops of reagent to Cd to give violet coloration.
Use solution in dilute HNO3 in Group separation.
If Cu present as a blue solution, first saturate reagent with KCNS and add crystal KI, then Cu is reduced and does not interfere.
12.11.3.16 Tests for calcium
Confirm by flame test: Brick-red (green through blue glass).
12.11.3.17 Tests for chromium
1. Fuse with sodium carbonate and a little potassium nitrate in a porcelain crucible.
Dissolve in water, add acetic acid and lead acetate solution.
A yellow precipitate forms.
A filtrate may contain chromium and aluminium as sodium chromate and sodium aluminate.
A yellow precipitate indicates the presence of chromium.
Pb2+ + CrO42- --> PbCrO4 (s), [yellow lead chromate].
2. Chromium (as chromate}
Organic reagent: Diphenyl carbazide, C13H14N4O, CO(NH.NH.C6H5)2, 0.2% solution in one
part glacial acetic acid and nine parts methylated spirit.
Make the chromate solution acidic with acetic acid or sulfuric acid.
Add reagent.
A deep violet-red colour indicates the presence of chromate.
Use in Group III when in form of chromate.
12.11.3.18 Tests for cobalt
Organic reagent: Nitroso-beta-naphthol, 1-nitroso-2-naphthol, C10H7NO2, 1 g in 50 mL acetic acid
Dilute to 100 mL Add reagent to neutral or slightly acid solution.
A brown colour indicates the presence of cobalt.
Use in Group IV when in solution after treatment with KClO3 and acid, or use the solution after Group III.
Cu, Fe, Sn, Ag, Cr, Bi, all interfere with the test.
12.11.3.20 Tests for iron
1. Dissolve ammonium thiocyanate in water and heat the solution.
The solution turns a characteristic red colour with iron (III) compounds, ferric compounds.
2. Organic reagent: Cupferron
NH4[C6H5N(O)NO] (ammonium salt of N-nitroso-N-phenylhydroxylamine), 5% aqueous solution
Filter if reagent is turbid.
Add reagent to strongly acidic HCl solution.
A red-brown compound indicates the presence of iron.
The reagents is unstable over long periods, but decomposition may be delayed by a piece of solid ammonium carbonate added to the reagent.
12.11.3.21 Tests for lead
1. Add potassium iodide solution to solutions of lead salts to form a yellow precipitate that is soluble in boiling water.
Organic reagent: Rhodizonic acid, C6H2O6, (CO)4(COH)2, dihydrate: C6H2O6.2H2O,
sodium salt (CO-CO.C.ONa)2, 0.1% aqueous solution
Add two drops of reagent to a sample of Group I precipitate still wet with acid.
A violet colour indicates the presence of lead.
Make a fresh solution of the reagent, if it has decolorized.
2. If test substances with heated charcoal and fusion mixture, lead forms grey-white soft globules and a red crust when hot, and a yellow crust when cold.
12.11.3.22 Tests for magnesium
1. Heat on charcoal with sodium carbonate.
Add a few drops of cobalt nitrate solution and heat again to produce a pink residue.
2. Organic reagent: The complex dye Titan yellow
Na2C28H19S4O6, as 0.1% aqueous solution.
Add 2 mL 1% KOH to 2 drops of test solution.
Boil to remove NH4+ and add 2 drops of Titan yellow.
A red colour or red precipitate indicates the presence of magnesium.
Tests for Mg in a Group VI solution.
Ammonium ions interfere with the test and must be removed.
12.11.3.23 Tests for manganese
1. Fuse manganese with sodium carbonate and some potassium nitrate in a crucible to form a blue-green mass.
2Mn(OH)2 + 5O --> 2MnO4- + 2H+ + H2O, [5O from oxidizing agents], [MnO4- = purple permanganate ion].
12.11.3.24 Tests for nickel
Organic reagent: Dimethylglyoxime, DMG (CH3C(NOH)C(NOH)CH3), 1% solution in methylated spirit, Toxic if ingested.
Warm a slightly acid test solution, add reagent then ammonium hydroxide until solution is alkaline.
A bright red precipitate indicates the presence of nickel.
Bismuth interferes with the test.
12.11.3.25 Tests for potassium
Reagent: Sodium perchlorate, 20% solution in equal parts of water and alcohol
Add reagent to equal volume of test solution.
A white precipitate of KClO4 indicates the presence of potassium.
Test in Group VI solution concentrated by evaporation and let cool.
Do the test on a glass plate above a black background.
12.11.3.26 Tests for silver
1. Add potassium chromate solution to neutral solution of a silver salt to form a brick-red precipitate.
2. Organic reagent: 5-(4-dimethylaminobenzylidene) rhodanine
C12H12N2OS2, 0.03% in acetone
The reagent detects AgCl in solution in water.
A red colour indicates the presence of silver.
Use a Group I precipitate.
12.11.3.27 Tests for sodium
Organic reagent: Uranyl magnesium acetate.
UO2(CH3COO)2.Mg(CH3COO)2, as a saturated aqueous solution
Add reagent to cold solution.
A yellow precipitate indicates the presence of sodium.
Test in Group VI solution concentrated by evaporation and let cool.
Do the test on a glass plate above a black background.
12.11.3.28 Tests for strontium
1. Confirm by flame test: Crimson
Sr2+ + SO42- --> SrSO4 (s).
2. Strontium and barium
Organic reagent: Rhodizonic acid (CO-CO.CONa)2, 0.1% aqueous solution
Put one drop of test liquid on filter paper then add one drop of reagent.
A red-brown spot indicates the presence of Sr and Ba.
When one drop of dilute HCl is added, a barium spot is intensified and a Sr spot disappears.
Use Group V precipitate after solution in dilute acetic acid.
Prepare fresh solution of reagent if it has decolorized.
12.11.3.29 Tests for tin
1. Use a borax bead containing some copper (II) sulfate.
Add a sample of the original solid and heat again to produce a red bead.
2. Organic reagent: Cacotheline, C21H21O7N3, saturated solution in water
Tin must be as Sn (II) in M HCl.
Add drops of reagent.
A violet colour indicates the presence of tin.
Stability of reagent is about 14 days.
Cu, Ni, Co, Cr and Fe interfere with the reaction.
Do the test with Group IIb solution when as Sn (II).
12.11.3.30 Tests for zinc
Filter and dissolve the precipitate in concentrated nitric acid.
Add a little cobalt nitrate solution, evaporate to concentrate, and soak a filter paper in the mixture.
Ignite the filter paper.
A green ash (Rimann's green) indicates the presence of zinc.
The green ash is a compound of zinc and cobalt oxides.
2ZnO22- + 8H+ + 2S2- --> 2ZnS (s) + 4H2O.
12.11.4.1 Acid-base neutralization, acid with base forms a salt and water
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.
Experiments
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.
Test 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, magnesium hydroxide and ammonia solution into 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.
HCl (aq) + NaOH (s) --> NaCl (aq) + H2O (l)
HCl (aq) + KOH (s) --> KCl (aq) + H2O (l)
2HCl (aq) + Ca(OH)2 (s) --> CaCl2 (aq) + 2H2O (l)
2HCl (aq) + Mg(OH)2 (s) --> MgCl2 (aq) + 2H2O (l)
HCl (aq) + NH3 (aq) --> NH4Cl (aq) + H2O (l)
[ H2O (l) + NH3 (aq) --> NH4+ (aq) + OH- (aq)].
4. Add 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.
H2SO4 (aq) + Mg(OH)2 (s) --> MgSO4 (aq) + 2H2O (l).
12.11.4.2 Acidity of vinegar and wine
Vinegar must have a minimum of 5% acidity if it is sold.
The percent acidity, % acidity (percent acetic acid) = (grams of acetic acid, CH3COOH / grams of vinegar) x 100
CH3COOH (aq) + NaOH (aq)--> CH3COONa (aq) + H2O (l)
Titrate 5.00 mL of household vinegar with sodium hydroxide solution.
The molarity of the sodium hydroxide solution is stated on the bottle label.
The molar mass of acetic acid (ethanoic acid) = 60.05.
The density of vinegar = 1.00 g / mL.
One mole of acetic acid reacts with one mole of sodium hydroxide
So the number of moles of acetic acid = molarity of the sodium hydroxide solution, in moles / litre X volume of sodium hydroxide solution, in litres.
Pour 5 mL of vinegar clean 250 mL Erlenmeyer flask and add 20 mL of distilled water and four drops of phenolphthalein indicator solution.
Add sodium hydroxide solution from the burette while swirling the flask, until the solution is a faint permanent pink and record the volume.
Discard the solution if you overshoot the end point and the solution is red.
Moles of acetic acid = (molarity of the sodium hydroxide solution, in moles / litre) x (volume of sodium hydroxide solution, in litres)
Grams of acetic acid = (moles of acetic acid) x (molar mass)
Grams of vinegar = (volume of vinegar) x (density of vinegar)
% acetic acid = (grams of acetic acid / grams of vinegar) x 100.
Measure acidity using a titration kit
To measure T.A. in wine, use an inexpensive titration or acid test kit.
Test kits can be purchased cheaply and can be used over and over again.
The amount of acid in wine is measured by slowly adding a small amount of the reagent base NaOH until a change in colour occurs from indicator phenolphthalein.
Put a 15 cc sample (one cc equals one ml) of wine into a test-tube.
Most test tubes that come with the acid test kits are marked with a line indicating this volume, or
use a small plastic syringe in the test kit to measure the desired amount into the test tube.
Then rinse the syringe.
Put 3 drops of the phenolphthalein solution into the test-tube.
Swirl or shake the test tube to mix the indicator with the wine.
Use the syringe to draw out 10 cc of the sodium hydroxide reagent, with no bubbles in the liquid.
Be careful! Avoid contact of the sodium hydroxide solution with your skin or eyes.
Add the sodium hydroxide solution to the test-tube, by 0.5 cc at a time.
After each addition of 0.5 cc, swirl or shake the test-tube to mix the contents.
The colour of the liquid will momentarily change upon the addition of the reagent.
If testing white wines, the colour change will be pink.
If testing red wines, the colour change will be grey.
Keep swirling the test-tube until the colour subsides.
If the colour of the wine returns to the original colour, repeat adding 0.5 cc until the colour change is permanent.
So when the colour, pink or grey, does not go away, stop and record the amount of reagent used.
To determine the acidity of the wine, for each cc of reagent used, equals 0.1 % TA.
For example, if you used 6 cc of sodium hydroxide to react with the wine, the titrarable acidity is 0.6 %.
Discard the sample, because it is now toxic.
Do not add the sample back into your original wine!
Wash and dry the test equipment before storing it.
12.11.4.3 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) + 2H2O (aq)
H2SO4 (aq) + 2NH4OH (aq) --> (NH4)2SO4 (aq) + 2H2O (l).
12.11.4.5 Group analysis
A saturated solution can remain in equilibrium with undissolved molecules of the solute.
Two equilibria exist: 1. an equilibrium between the undissolved solute and dissolved molecules, and
2. an equilibrium between dissolved molecules and ions formed by dissociation
(XY) <--> XY <--> X+ + Y-
Equilibrium 1.: (XY) <--> XY
Equilibrium 2.: XY <--> X+ + Y-
(XY) = undissolved molecules
XY = dissolved, but unionized molecules
X+ + Y- = ions
The tendency of the solid to pass into solution depends on its active mass, solution pressure.
If the temperature remains constant, the active mass remains constant.
This happens because, by the law of mass action, concentration dissolved molecules / concentration undissolved molecules = the constant, K.
The concentration of dissolved molecules is also a constant.
By the law of mass action, if [concentration X+] [concentration Y-] / [concentration dissolved molecules] = a constant, in a saturated solution
Then the product of the concentrations of the ions is a constant.
If large concentrations of different ions are brought together into the same solution,
then ions of X+ and Y- will precipitate out of solution as solid molecules until the concentrations of the remaining ions in solution have the product of their concentrations equals the specific constant, the solubility product.
Group I
Lead, silver and mercury (I) are precipitated as chlorides by chloride ions from hydrochloric acid.
The concentrations of silver and chloride ions that can remain in solution are small.
When a solution of a silver salt containing silver ions is mixed with hydrochloric acid, most of the silver and chloride ions form molecular silver chloride and leave the solution as a solid phase until
the remaining ions attain equilibrium.
(concentration silver ions) × (concentration chloride ions) = 1 × 10-10, the solubility product.
The solubility product has a constant value, so adding excess chloride ions reduces the concentration of the silver ions to a negligible quantity.
Only silver chloride, lead chloride and mercury (I) chloride have low solubility products.
So the ions of other metals remain in solution in the presence of high concentrations of chloride ions.
Group II
Assume hydrogen sulfide is ionized
H2S <--> 2H+ + S2-
By the law of mass action (concentration H+)2 × (concentration S2-) / (concentration unionized H2S) = the constant, 1.1 × 10-22.
In a neutral solution, the concentration of sulfide ions is low, because hydrogen sulfide is a weak electrolyte.
The concentration of hydrogen ions is also low.
In the acid solution used for Group II, the concentration of hydrogen ions is increased by the presence of the strong acid.
So to maintain the value of the solubility product constant, the concentration of the sulfide ion is reduced below its already small value in neutral solution.
However, the amount of sulfide ions is enough to allow the solubility products of the sulfides of mercury (II), lead, copper and bismuth to be exceeded.
Also, cadmium sulfide may precipitate if the acid is not too concentrated.
So in Group II, all the sulfides of mercury (II), lead, copper, bismuth and cadmium precipitate.
Solubility products:
Lead sulfide 4 × 10-28, Copper sulfide 8 × 10-45, Mercury (II) sulfide 4 × 10-54,
Cadmium sulfide 3.6 × 10-29, Manganese sulfide 1.4 × 10-15, Zinc sulfide 1.2 × 10-24
The concentration of sulfide ion in acid solution is not enough to allow the higher solubility products of the sulfides of manganese, zinc, cobalt or nickel to be reached with any possible concentration of the metal
ion, so these sulfides do not precipitate.
They precipitate later in Group IV, where the precipitating agent is the highly ionized salt, ammonium sulfide, and the concentration of the sulfide ion from it is high.
Group III
The precipitating agent is ammonia solution.
NH4OH <--> NH4+ + OH-
By the law of mass action: (concentration NH4+) × (concentration OH-) / (concentration unionized NH4OH) = a constant.
Ammonia solution is a weak base, so does not ionize much and the value of the constant is only about 2 × 10-5.
Most of the ammonia solution will be dissolved, but not ionized.
The small hydroxyl ion concentration in a solution of ammonia solution that is also fairly concentrated with respect to ammonium chloride is still large enough to cause a precipitation of the hydroxides
of ferric iron, chromium and aluminium, but not great enough to precipitate the hydroxides of zinc, manganese, cobalt and nickel.
Manganese hydroxide may precipitate slightly if the concentration of ammonium chloride is not sufficiently great.
Group IV
In Group II the presence of hydrogen ions from the added acid reduces the concentration of sulfide ions, but this reduced value was enough to allow the solubility products of the metallic sulfides in the group to be exceeded.
In Group IV, hydrogen sulfide is added to a solution made alkaline with ammonia solution and so contains excess of hydroxyl ions.
By the law of mass action: (concentration H+)2 × (concentration S2-) / (concentration unionized H2S) = a constant.
The ionic product of water [H+] × [OH-1] = (10-14).
The hydroxyl ions from the ammonia solution lower the concentration of hydrogen ions causing an increased concentration of S2- ion.
In Group IV, the metal sulfides not already precipitated in Group II, because of their high solubility products, precipitate.
So the ionic concentration of the sulfide ion is controlled by variation of the concentration of the ions it is associated with, i.e. hydrogen ions.
Group V
The metals still remaining in solution include barium, strontium, calcium, magnesium, sodium and potassium.
Barium, strontium and calcium are precipitated as carbonates by the addition of CO32- ions in alkaline solution, because of the low values of the solubility products,
[X2+] × [CO32-] = K< between 10-8 and 10-9.
The solubility product of magnesium carbonate is low, 10-5, and in neutral solution would be precipitated, but its precipitation is prevented in this group by the ammonium ions, mainly from
the ammonium
chloride added before Group III.
With the large NH4+ ion concentration from this source, the concentration of CO32- ions is reduced below the concentration to precipitate magnesium as a carbonate.
[NH4+]2 [CO32-] / [NH4)2CO3] = K.
12.11.4.10 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).
12.11.4.9 Simple titration of acids and bases
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.
12.11.4.11 Titrate dilute hydrochloric acid with sodium hydroxide solution, with a burette
See diagram 12.8.4: Titration
1. Pour a hydrochloric acid solution of known concentration (such as 0.11mol / 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.
NaOH (aq) + HCl (aq) --> NaCl (aq) + H2O (l).
2. 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.
NaOH (aq) + HCl (aq) --> NaCl (aq) + H2O (l).
12.11.4.12 Titrate dilute sulfuric acid with sodium hydroxide solution and isolate sodium sulfate crystals
Safety
* Wear protective gloves when handling corrosive chemicals.
* Handle acids and alkalis with care.
* If any alkali gets into your eyes or onto your skin, report to your teacher immediately, and wash the affected area under running water for at least three minutes.
* If any acid gets into your eyes, report to your teacher immediately, and flush your eyes with running water for at least three minutes.
* If any acid gets onto your skin, wash the affected area with plenty of water.
* Eye protection must be worn.
* Sodium hydroxide 2.0 M is corrosive.
* Sulfuric acid 1.0 M is an irritant.
1. Wash a burette as follows:
1.1 Use a filter funnel to fill a burette with 20 mL of water.
1.2 Hold the burette horizontally and rotate it slowly to wash the inner wall.
Open the stopcock to run out all the water into the sink.
Close the stopcock.
Repeat the above procedure with 1.0 M sulfuric acid.
2. Fill the burette as follows:
2.1 Fill the burette too below the zero mark with 1.0 M sulfuric acid.
2.2 Clamp the burette vertically in a stand.
2.3 Open the stopcock for a few seconds to fill the burette jet completely with the acid.
Do not leave any bubbles in the burette jet.
2.4 Take the initial burette reading (V1) = mL (to two decimal places).
3. Wash a pipette as follows:
3.1 Using a pipette filler, suck sufficient water into a pipette to fill part of the bulb.
(Never use your mouth to suck a pipette.)
3.2 Hold the pipette horizontally.
Rotate it slowly so that the water washes the inner wall, up to the graduation mark.
3.3 Allow the water to run out into the sink.
Repeat 3.1 to 3.3 using 2.0 M sodium hydroxide solution.
4. Titration:
4.1 Using the pipette filler and the washed pipette, transfer 25.0 cm of the sodium hydroxide solution into a clean conical flask.
4.2. Add 2 drops of methyl orange indicator to the conical flask.
Note the colour of the solution in the conical flask.
4.3 Run the 1.0 M sulfuric acid into the flask while swirling the flask continuously.
4.4 Continue adding the acid, until the solution in the flask just turns to an orange colour.
4.5 Take the final burette reading (V2) = mL (to two decimal places).
4.6 The volume of sulfuric acid needed to neutralize 25.0 mL of sodium hydroxide solution is (V2 - V1) mL.
5. Isolation of sodium sulfate crystals from solution:
5.1 Empty the conical flask and rinse it with water.
5.2 Use a pipette filler to pipette 25.0 mL of 2.0 M sodium hydroxide solution into the flask.
Do not add any indicator to allow preparation of pure sodium sulfate crystals.
5.3 From the burette, add the same volume of sulfuric acid as calculated from 4.6, i.e. (V2- V1) mL.
This results in complete neutralization to form a hot concentrated sodium sulfate solution and water.
5.4 Pour the resultant solution into a small beaker.
5.5 Boil the solution gently to concentrate it.
Heat with a small flame, until 1/3 of the solution is left.
Do not evaporate the solution until dry, because spitting will occur.
5.6 Every 10 seconds, dip a glass rod into the boiling solution and then take it out.
If the immersed end becomes "cloudy" within 5 or 6 seconds, stop heating.
The solution has now become concentrated enough to deposit crystals on cooling.
5.7 Leave the solution to cool overnight.
5.8 The next day, sodium sulfate crystals should have formed.
5.9 Filter the crystals from the remaining solution.
5.10 Wash the crystals with drops of distilled water from a plastic wash bottle.
5.11 Use a spatula to transfer the crystals onto a piece of filter paper or absorbent paper.
5.12 Dry the crystals by gently pressing them between filter paper of absorbent paper.
5.13 Store the crystals in a labelled enclosed container.
12.11.4.4 Carbon dioxide affects 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.
12.11.4.6 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 sodium hydroxide = volume of HCl X 1 / 1000 X 2.
12.11.4.7 Microscale titration, sodium hydroxide with dilute acids
See diagram 12.8.7: Microscale titration apparatus
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 spillage and breakage 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.
12.11.4.8 Prepare monoprotic acid solution from unknown molarity acid
(A monoprotic acid donates one hydrogen ion (H+) per molecule when it dissociates in water.
A monoprotic acid, (HZ),dissociates according to the following equation:
HZ (aq) + H2O (l) ⇌ H3O+ (aq) + Z− (aq)
Monoprotic acids include: | Acetic acid | Formic acid | Hydrochloric acid | Nitric acid |
Diprotic acids include: | Sulfuric acid | Carbonic acid | Oxalic acid | Citric acid |)
1. 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. Pour 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.
Pour 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 0.1M 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.