School Science Lessons
(topic 07b)
2024-12-26
Chemistry terminology
Contents
7.1.6 Chemical changes and physical changes
7.9.0 Chemistry terminology
7.1.5 Prepare forms of sulfur
7.2.0 Pure substances and impure substances
7.2.1 Classify substances
7.1.6 Chemical changes and physical changes
7.1.6.1 Chemical changes and physical changes
Experiments
7.1.6.2 Breakdown starch to sugars
7.1.6.3 Breakdown ethanol to ethene (ethylene)
7.1.6.4 Burn magnesium and weigh the products
7.1.6.5 Burn steel wool and weigh the products
7.1.6.6 Empirical formula of magnesium oxide
7.1.6.7 Chemical changes, heat organic substances
7.1.6.8 Chemical changes, heat metals in chlorine
7.1.6.9 Physical changes, magnetize and demagnetize iron wire
7.1.5 Prepare forms of sulfur
7.1.5.0 Prepare forms of sulfur, allotropes of sulfur
7.1.5.1 Prepare monoclinic sulfur from powdered sulfur (flowers of sulfur)
7.1.5.2 Prepare monoclinic crystals from roll sulfur
7.1.5.3 Prepare plastic sulfur then rhombic sulfur
7.1.5.4 Prepare forms of sulfur
7.9.0 Chemistry terminology
7.9.1.0 Acid, acids
7.9.2 Aerobic
7.9.3 Aliquot
7.9.4.2 Allotropes, sulfur, carbon
7.9.6 Amphoteric
7.9.12 Anti-bumping granules, boiling chips
7.9.6.2 Aqueous solutions, states of matter
3.0.0 Atoms
7.9.7 Constant boiling mixture, azeotrope
7.9.8 Base
7.9.51.1a Battery sulfation
7.9.9 Bessemer process
7.9.10 Borosilicate glass, Pyrex
7.9.11 Buffers
7.0.0 Chemical compounds
7.9.14.1 Chemical potential energy, enthalpy
7.9.14.3 Clathrate
7.9.15 Continuous phase / outer phase
7.9.16 Detergent, SYNDET, Synthetic Detergent
8.0.0 Direct union of elements to form compounds, (Experiments)
2.0.0 Elements
8.2.17 Elements combine with oxygen gas when heated in air
7.9.19 Enzymes
7.9.20 Equilibrium
7.9.20.1 Etchants
7.9.22 Flammable
7.9.24 Flocculent
7.9.26 Flux
7.9.27 Froth flotation
7.9.28 Fuel cell
7.9.29 Galvanize
7.9.30 Group formula
7.9.31 Heavy metals
7.9.32 Hydronium ion (hydroxonium ion, oxonium ion)
7.9.33 Hydrophilic, hydrophobic
7.9.35 Inhibit, (Chemistry)
5.0.0 Ions
7.9.36 Labile, (Chemistry)
7.9.37 Martensite
1.0.0 Matter
1.1.0 Micronization
4.0.0 Molecules
7.9.39 Molecular mass
7.9.40 Napalm
7.9.41.1 Petroleum ether
7.9.41 Petroleum fraction
7.9.41.1 Petroleum spirit
7.9.42 Photolysis
7.9.46 Radical, (Chemistry)
7.9.47 Sequester, (Chemistry)
7.9.48 Solute
7.9.49 Solvent
7.9.50 Spontaneous process
7.9.51 Substrate
7.9.51.1 Sulfates, battery sulfation
7.9.51.2 Sulfides
7.9.51.3 Sulfites
7.9.52 Surfactants
7.9.53 Synergism
34.7.6 Tempering
7.9.54.2 Urethane
7.9.55.1 Waste chemical bottles
7.9.56.1 Xanthene dyes
7.9.57 Zymase
1.0.0 Matter
Matter is composed of atoms, which, in turn, contain protons and neutrons in a nucleus, and electrons outside the nucleus.
The number of positively charged protons is equal to the number of negatively charged electrons in a neutral atom, and determines all the chemical properties of an atom.
Materials may be elements, compounds or mixtures.
1.1.0 Micronization
Size reduction mills are used in the pharmaceutical industry to increase surface area, improve formulation dissolution and maintain a consistent average particle size distribution for tablets and capsules.
Impact size reduction is usually by mechanical impact and impact via fluid energy using hammer mills and liquidized bed jet mills.
Micronization that uses supercritical fluids to cause the small diameter of the solid particles by the supersaturation at the time of the particle formation, includes the following:
* RESS process (Rapid Expansion of Supercritical Solutions), where the solution is expanded through a nozzle,
* SAS method (Supercritical Anti-Solvent), where the material precipitates out of the solution as a solid with a very small particle diameter,
* PGSS method (Particles from Gas Saturated Solutions), where the solid is melted and the supercritical fluid is dissolved in it as in the SAS method, but the solution is forced to expand through a nozzle.
The PGSS method is used for pharmaceuticals and foodstuffs.
2.0.0 Elements
Elements cannot be broken down into simpler substances by a chemical change.
An element is a substance in which all atoms have the same number of protons.
Atoms of an element may contain different numbers of neutrons, and are known as isotopes.
Every element is assigned a unique chemical symbol.
At room temperature and atmospheric pressure, eleven elements are gases: H2, He, N2, O2, F2, Ne, Cl2, Ar, Kr, Xe, Rn.
Two elements are liquids at room temperature: Hg and Br.
Reactive elements have atoms weakly bound together and have electrons available for bonding, e.g. F.
Unreactive elements have atoms joined by strong bonds, e.g. diamond.
Elements that exist as separate small molecules have low boiling points and melting points, e.g. He, O2.
Experiment
Describe each example.
1. Note the state of matter at room temperature, solid, liquid or gas.
2. Note whether the solid has a shiny surface or has a lustre when the surface is clean.
3. Note whether the metal can be bent or twisted with pliers, or whether it fractures.
4. Note whether the element conducts electricity when held between two alligator clips as electrical contacts.
5. Put a piece of the element on a combustion spoon and set it alight with a burner flame.
Observe the burning element.
6. Shake the products of the combustion in a test-tube containing water.
Test the solution with moist litmus paper.
8.2.17 Elements combine with oxygen gas when heated in air
1. Put a small quantity of sulfur on a deflagrating spoon and set it alight with a Bunsen burner flame.
Note the appearance of the burning sulfur and then lower it into a test-tube containing oxygen gas.
Be careful!
The sulfur dioxide produced has a very irritating odour and may cause distress to people with asthma.
When you heat sulfur, it melts, turns brown, and burns with a blue flame.
It burns more vigorously in oxygen gas than in air.
During the burning it combines with oxygen gas to form the compound sulfur dioxide:
sulfur (s) + oxygen (g) --> sulfur dioxide (g)
2. Repeat the experiment using the following:
* Iron (steel wool),
* Magnesium.
Be careful!
Do not look at the burning magnesium!
The light may injure the eyes!
* Carbon.
In each case, note whether the substance burns more rapidly in oxygen gas than in air.
3.0.0 Atoms
Atom is the smallest division of an element that can chemically exist alone and have the characteristics of the element.
Atomic mass (atomic weight), of an atom is arbitrarily defined relative to the mass of the isotope carbon.
The relative atomic mass of an element is the ratio of average mass of atoms of the element to 1 / 12 of the mass of one atom of the isotope C-12.
Atomic mass unit (a.m.u.), is a unit of mass used to express the relative atomic mass.
So 1 atomic mass unit = 1 / 12 of the mass of the carbon-12 isotope, i.e. 1.66 × 10-27 kg.
The mass of 1 atom of oxygen = 16 a.m.u.
The atomic theory dates from the 5th century B.C. when Greek philosophers, e.g. Democritus, said that matter consists of indivisible indestructible particles.
The modern atomic theory started with the hypothetical thinking of John Dalton, England (1766 - 1844).
However, he did not envisage the structure of atoms, i.e. nucleus, electrons and other particles.
4.0.0 Molecules
Molecules have two or more atoms joined chemically to form the simplest stable structure of that element or compound.
So a hydrogen molecule must contain 2 atoms, H2, because the hydrogen atom, H, cannot exist by itself.
A molecule of water, H2O, contains 2 atoms of hydrogen, H, and 1 atom of oxygen, O.
Molecular formula represents the number of atoms of each element in the molecule.
The molecular formula of water is H2O, so water contains two H atoms and one O atom.
The empirical formula of a compound shows the ratio of elements present in the compound.
Glucose has a molecular formula of C6H12O6, so it contains 2 moles of hydrogen for every mole of carbon and oxygen.
The empirical formula for glucose is C6H12O6/6 = CH2O.
5.0.0 Ions
Ions form when an atom or group of atoms covalently bound together may gain or lose one or more electrons.
Ionic bonding occurs when positive and negative ions are held together in a crystal lattice by electrostatic forces.
The structure of metals involves positive ions embedded in a sea of electrons.
7.0.0 Chemical compounds
1. Compounds are composed of two or more elements that are chemically united in fixed proportion.
Compounds can be broken down to simple substances.
Chemical compounds form when chemical bonds, whether ionic or covalent, are formed between different elements.
A chemical compound can be represented by a chemical formula.
Forces weaker than covalent bonding exist between molecules.
In compounds containing carbon-hydrogen bonds, organic compounds, the carbon atoms bind to one another through single, double or triple covalent bonds to form chains or rings.
Observe samples of iron, carbon, copper, magnesium, mercury in a thermometer and solid sticks of sulfur.
2. Classify each compound according to the following characteristics:
* Is it a hard solid, a liquid or a gas at room temperature?
* If the element is a solid, does it shine and have a lustre?
If necessary, scratch its surface.
* Can the element be bent or twisted or does it fracture easily? You may require pliers to do this.
* Will the element conduct electricity?
Place the element between two electrical contacts in a circuit.
7.1.6.1 Chemical changes and physical changes
1. In a chemical change, one or more substances changes into one or more new substances, e.g. hydrogen gas combines with oxygen gas to form water.
In this document the expression used is as follows: "hydrogen gas with oxygen gas forms water".
The new substances, the products, have properties different from the original substances, the reactants.
For example burning wood forms black ash containing carbon, white ash containing mineral salts, carbon dioxide gas and water vapour.
New substances form and the change cannot be reversed.
The arrow symbol " -->" represents this type of change.
2H2 (g) + O2 (g) --> 2H2O (l).
2. In a physical change, the properties: of a substance change, but the substance is still the same.
The change can be reversed.
The physical properties: of water change when water freezes to ice.
However, the ice is composed of water molecules and the change can be reversed.
When electricity passes through a tungsten filament in a light bulb, the filament becomes hot and emits light.
It is still tungsten.
When you turn off the light, the tungsten filament is the same as before.
Physical or chemical changes may be fast or slow, e.g. hammering iron, wearing away rock by wind erosion, an explosion when hydrogen gas burns, rusting of iron.
3. Classify common examples of changes into chemical change or physical change.
Do the following activities and in each case say whether it is a chemical or physical change:
1. Light a match (chemical change).
2. Turn on the light (physical change).
3. Let a nail rust (chemical change takes place in air causing a gain in weight).
4. Chew food (physical change for biting and masticating, and chemical change for the reactions of starch with salivary amylase).
7.1.6.4 Burn magnesium and weigh the products
Put a piece of magnesium ribbon in a crucible.
Weigh the crucible and lid and magnesium ribbon.
Burn the magnesium in the nearly closed crucible so that the ash is not lost.
Weigh the crucible and lid and ash.
The weight of the ash is greater than the weight of the magnesium.
The reaction forms a new substance, magnesium oxide.
Mg, density 1.74 g / cm3 oxidizes to MgO, density 3.58 g / cm3, so when oxygen attaches to magnesium the volume decreases, because of the strength of the (Mg2+-- O2-) bond.
2Mg (s) + O2 (g) --> 2MgO (s).
7.1.6.5 Burn steel wool and weigh the products
Cover the end of a wooden ruler with aluminium foil then note where the ruler balances as a first order lever, see-saw, over a fulcrum.
Put the ruler balance on a sink bench.
Put 5 g loose steel wool on the aluminium foil end of the ruler then put a weight on the other end of the ruler so that this end is just down.
Heat the steel wool with a Bunsen burner flame.
The steel wool glows and its side of the balance moves down as the iron becomes iron oxide.
7.1.6.6 Empirical formula of magnesium oxide
Method 1.
To obtain useful data when burning magnesium, problems may occur with igniting the Mg, keeping the Mg burning and recording very little or no mass gain.
Using only 3 cm is not enough so coil it up and use about 25cm in a crucible over a Bunsen burner to obtain reasonable results.
Use sandpaper on the magnesium before experimenting tp remove the oxide layer, then place the magnesium in the crucible, leaving the lid open until it was just about to react.
Make sure that as much of as possible of the coiled longer piece is in contact with the crucible.
Leave the lid off until it ignites, then, while still heating, keep raising the crucible lid with tongs to let in more oxygen, without releasing too much ash. until the reaction appears complete.
Method 2.
Weigh magnesium then heat it in a crucible so that it reacts with oxygen to produce magnesium oxide.
There has been an increase in mass.
Use the results can be used to find the formula of magnesium oxide.
Procedure
If the magnesium ribbon looks tarnished or black, clean it with sand paper.
1. Cut a piece of magnesium ribbon about 10cm long and twist it into a loose coil.
2. Weigh the crucible with the lid and then weigh the magnesium inside the crucible with the lid.
3. Set up the Bunsen burner on the heat resistant mat with the tripod.
Place a pipe clay triangle over the tripod.
Place the crucible containing the magnesium in the pipe clay triangle and put the lid on at an angle so air can still get into the crucible.
4. Light the Bunsen burner and begin to heat the crucible.
Start with a gentle blue flame, but you will need to use a roaring flame, with the air hole fully open, to get the reaction.
5. When the crucible is hot, gently lift the lid with the tongs a little to allow some oxygen to get in.
The magnesium may begin to flare up.
Do not lift the lid is off for too long, because the magnesium oxide product will begin to escape from the crucible.
. Repeat this procedure until you see no further reaction.
6. Turn off the Bunsen burner, allow the apparatus to cool, then re-weigh the crucible with lid containing the product.
7. Heat the crucible again for a couple of minutes, allow to cool, and repeat this step until the mass readings are consistent.
This procedure is called ""heating to constant mass".
2Mg + O2 --> 2MgO
Magnesium + oxygen --> magnesium oxide
To find the formula of magnesium oxide, divide mass by the atomic mass for each element, to give the number of moles of each.
The relative atomic mass of magnesium is 24 and for oxygen is 16.
Find the ratio between the two by dividing them both by the smallest number.
The ratio should be close to 1:1, because the formula of magnesium oxide is MgO.
Example calculation
Mass magnesium = 2.39 g, Mass magnesium oxide = 3.78 g, So mass oxygen = 1.39 g
Number moles Mg = 2.39/24 = 0.0995
Number moles O = 1.39/16 = 0.0868
Divide by the smallest to give the ratio approx. 1 Mg: 1O.
This result suggests a formula of MgO, which is the correct formula.
One teacher got students to all do different lengths, and them plot moles of Mg on X axis and moles of O on Y axis (by subtraction and processing) and tried to get a gradient of 1.
Possible errors include the following:
* the magnesium oxide product may escape as students lift the lid,
* not all the magnesium may have reacted, because the product may still look a bit grey rather than white,
* students may have prodded the product with their splint, so not all of it got weighed,
* the balanced was not tared correctly for one of the weighings,
* the magnesium was coiled too tightly so that not all of it reacted,
* the tripods not at the correct height so that the hottest part of the flame not exactly where the crucible is sitting in the clay triangle,
* some Bunsen burners produce much hotter flames than others, so find one that produce a good flame or use two Bunsen burners if the gas turrets, have two taps, to put two burners on with it,
* if start seeing puffs of whitish “smoke coming out, MgO is being lost,
* students stopped heating before all of the magnesium was powdered ash.
* students failed to record the mass of the empty crucible,
* students could only obtain an around 70% yield.
7.1.6.7 Chemical changes, heat organic substances
1. Often one product is black, because carbon forms.
Heat the following substances in a hard glass or Pyrex test-tube.
Observe any decomposition of the substances.
When smoke appears, leave the test-tube to cool.
No chemical change occurs when you heat charcoal.
For the following substances, a chemical change occurs and the residue in the test-tube is carbon:
Experiments
1. Heat wood or saw dust, and note the brown substances that stick to the sides of the test-tube and the black substance that remain in the bottom.
2. Heat the following substances, sugar, bread, fruit, fat or oil, starch, potato,rice, paper, wool, hair, nail clippings, meat.
Heat the tube gently then more strongly.
The wood decomposes into a black solid and brown liquids and vapours.
Stop heating when you see no more dense smoke.
Let the test-tube and observe the substances on the sides of the test-tube.
Put any remaining solid substance in another clean test-tube.
Heat it gently and then more strongly.
You produce no further liquids or gases from this substance.
Compare the black substance with a sample of the original wood for colour and can be bent.
Place the end of each in a Bunsen flame for a few seconds.
Remove and observe them.
The black substance does not decompose or burn with a flame but, glows red hot.
The black substance is charcoal, i.e. carbon.
3. Heat small quantities of sugar, paper, wool, meat and compare the residue in the test-tubes with the original substance.
7.1.6.8 Chemical changes, heat metals in chlorine
Be Careful! Chlorine is poisonous.
These experiment may not be allowed on your school system.
Do the experiments in a fume cupboard.
Experiments
1. Use tongs to heat steel wool in a Bunsen burner flame, then place it in a test-tube containing chlorine gas.
Observe whether any heat is given out during the reaction.
Iron and chlorine react together to form a brown yellow iron chloride.
2. Using a wire gauze, scrape some powdered antimony from a lump of the element so that it falls into a test-tube of chlorine.
The sparks and burning show that a chemical change is taking place.
The new white substance is antimony chloride.
7.1.6.9 Physical changes, magnetize and demagnetize iron wire
When physical changes occur, no new substance forms.
Pull a thick iron wire (fencing wire), through iron filings.
The iron does not attract the iron filings.
Magnetize the iron by stroking it with a bar magnet.
Hold the iron wire in the iron filings.
The iron wire now attracts the iron filings, so a physical change has occurred.
Hammer the iron wire, the iron wire cannot attract the iron filings so strongly, so the physical change has been reversed.
7.1.5.0 Physical changes, prepare forms of sulfur, allotropes of sulfur
See diagram 12.18.1: Sulfur crystals.
1. Sulfur is a non-metallic element that occurs in several allotropic forms.
Allotropes are variations of the same element with bonding and crystal structure.
2. Sulfur occurs in three forms:
Form 1: Rhombic sulfur is a light yellow powder.
At 100oC, it changes to:
Form 2: Monoclinic sulfur with a deeper colour.
The monoclinic and rhombic forms differ in the arrangement of the S8 molecules.
At 160oC, sulfur melts to form a sticky dark brown liquid that can be cooled quickly to form red brown plastic sulfur.
Form 3: Plastic sulfur contains long chains of S atoms.
Experiments
Put sulfur powder in a test-tube.
Heat the sulfur extremely gently until it slowly melts to a golden yellow liquid.
Continue to heat more strongly until a red gas appears above the liquid.
Leave the test-tube to cool.
Sulfur forms deposits on the sides of the tube and in the bottom of the tube.
7.1.5.1 Prepare monoclinic sulfur from powdered sulfur (flowers of sulfur)
Monoclinic sulfur has a deep yellow colour, m.p. = 119oC, and density = 1.96.
Heat powdered sulfur extremely gently in an evaporating dish.
The sulfur changes to liquid.
Add more powdered sulfur.
The colour should stay pale yellow.
If the sulfur turns dark, you have overheated it.
Repeat the experiment with gentler heating.
Leave the sulfur to cool, without moving the evaporating dish.
Monoclinic crystals form between 96oC and 114oC.
After a thin crust forms, punch two holes through the crust with a nail.
Pour out the hot sulfur through one hole.
Remove the crust and note the monoclinic sulfur crystals on the underside.
7.1.5.2 Prepare monoclinic crystals from roll sulfur
Put small pieces of roll sulfur in a test-tube.
Heat the test-tube extremely gently until the sulfur melts, m.p. = 114.5oC.
The liquid is lemon yellow.
Pour the liquid sulfur into a folded filter paper.
A crust forms on the surface.
When the crust forms, open the filter paper.
Needle-shaped crystals remain on the filter paper.
7.1.5.3 Prepare plastic sulfur then rhombic sulfur
When heated to the melting point, sulfur usually ignites and forms sulfur dioxide gas that may distress people suffering from asthma.
Rhombic sulfur is a yellow powder, m.p. = 119oC, and density = 1.96.
Experiment
Put roll sulfur in a test-tube and heat the test-tube slowly.
Note the changes at the melting point from a light yellow colour to a red liquid.
The red brown sulfur becomes viscous and it does not flow out if you hold the test-tube upside down.
Continue heating until the reaction forms a brown black liquid.
Continue to heat until the sulfur boils at 445oC.
BE CAREFUL!
Pour the melted sulfur into a beaker of cold water.
The reaction forms strands of amorphous sulfur.
When the strands are cool, twist them to show that they are elastic.
Later the sulfur hardens, because it returns to the rhombic form of sulfur as a ring of eight sulfur atoms in tiny crystals.
When heated, the ring breaks open to form long chains.
7.1.5.4 Prepare forms of sulfur
See diagram 12.18.1: Sulfur crystals
Almost fill a dry test-tube with sulfur powder and heat slowly to boiling, using a safety holder.
Pour the boiling sulfur into a beaker of water.
Immerse any floating sulfur with a stirring rod.
Remove and examine the plastic sulfur.
Note the gradual loss of elasticity as the plastic sulfur changes to rhombic sulfur.
7.1.5.5 Prepare sulfur dioxide by heating sulfur
Use a test-tube with a stopper to heat sulfur in a fume cupboard or well-ventilated area.
When heated to the melting point, sulfur usually ignites and forms sulfur dioxide gas that may distress people suffering from asthma.
1. When you heat sulfur, it melts, turns brown, then burns with a blue flame.
Sulfur combines with oxygen gas to form sulfur dioxide.
2. Heat sulfur in air
Be careful!
The gases that form have an irritating odour and may cause distress to people who suffer from asthma.
Put sulfur in a combustion spoon and set it alight with a flame.
Observe the burning sulfur and then lower it into a test-tube containing oxygen gas.
The gases turn moist blue litmus red.
S(s) + O2 (g) ---> SO2 (g).
3. Heat sulfur gently in a crucible in a fume cupboard.
The solid melts to form a clear yellow liquid.
Melting sulfur without also igniting it is difficult.
Ignition is indicated by the formation of a blue flame on the surface, and the formation of acrid sulfur dioxide gas.
4. Put sulfur powder in a large test-tube.
Clamp the test-tube in a horizontal position and insert a piece of coiled copper wire 3 cm from the sulfur.
Heat the sulfur and copper alternately for five minutes with a strong Bunsen burner flame, with most of the heat on the copper.
The sulfur vapours blacken the copper and changes its electrical properties.
5. Put a small quantity of sulfur on a deflagrating spoon and set it alight with a Bunsen burner flame.
Note the appearance of the burning sulfur and then lower it into a test-tube containing oxygen gas.
Be careful! The sulfur dioxide produced has a very irritating odour and may cause distress to people who suffer from asthma.
When you heat sulfur, it melts, turns brown and burns with a blue flame.
It burns more vigorously in oxygen gas than in air.
During the burning it combines with oxygen gas to form the compound sulfur dioxide:
sulfur (s) + oxygen (g) --> sulfur dioxide (g).
7.2.0 Pure substances and impure substances, elements and compounds
1. Pure substances contains only one kind of atom or molecule, e.g. iron, sulfur, water, and oxygen.
It can consist of elements or compounds.
Elements may be metals or non-metals.
Impure substances contain more than one kind of substance.
They may be mixtures, e.g. sulfur and iron filings, or air, or solutions, e.g. sea water.
Mixtures in the home include flour, milk, ink, face powder, and tooth paste.
Solutions in the home include fruit drinks, (although often an emulsion), lemonade, mineral water.
You can make toffee or candy with a supersaturated sugar solution.
French polish is shellac with methylated spirit solution.
2. An element cannot be broken down into simpler substances by a chemical reaction and all the atoms in it have the same atomic number, the number of protons and electrons is the same, but the number of neutrons may vary.
Ninety-two (92), naturally occurring elements exist.
3. A compound is composed of two or more elements combined in fixed proportions as a result of a chemical reaction and which cannot be separated into simpler substances by any physical process, e.g. shaking.
Most compounds are ionic, e.g. common salt (NaCl), or covalent molecular, e.g. CO2 or H2O, or covalent network, e.g. SiO2.
Each compound has a name or formula.
4. Inorganic chemistry is the chemistry of the elements and their compounds, including CO2, CO and carbonates.
Organic chemistry is the chemistry of carbon and its compounds.
7.2.1 Classify substances, pure substances, mixtures, solutions
Describe materials at home and in the classroom from their observable physical properties:
1. colour (shiny or dull),
2. opaque (or transparent or translucent),
3. shape (or shape of crystals),
4. odour,
5. state (solid, liquid, gas) (change in state when heated or cooled),
6. mass (heavy, light), more dense or less dense than water,
7. taste (sweet, sour, bitter, other),
8. can be magnetized,
9. can conduct heat,
10. can conduct electricity,
11. can absorb liquids,
12. flexibility (can be bent, fragile).
Describe the materials used to make items in the classroom and in the home.
7.2.2 Describe materials as pure substance, or solution or mixture
If a solution is a homogeneous mixture of a liquid with a gas or a solid, then you would not classify air as a solution, nor brass as an alloy, unless they were in liquid form.
However, if a solution is a homogeneous mixture of two or more components in a single phase, and usually refers to a solution in water (aqueous solution), then perhaps you can classify air and solid alloys as solutions.
Table 7.2.1
Substances Pure substance Solution Mixture
1. Ice floats in water + -
 -
2. Tincture of iodine -
 + -
3. Washing powder -
 -
 +
4. Tap water -
 + -
5. Air -
 + or mixture +
6. Brass alloys -
 + or mixture +

7.9.1.0 Acid, acids
(Latin: acidus, sour tasting)
An acid is a good electrolyte, reacts with active metals, turns blue litmus red and has a sour taste.
Acids contain hydrogen replaceable by metals.
Acids neutralize alkalis and alkalis neutralize acids.
Acids are sour corrosive, mainly liquids, that can dissolve metals.
Acids with water produce hydrogen ions, H+.
An acid is a proton donor (H+) (Bronsted-Lowry definition).
Acids an donate protons or accept pairs of electrons.
The "acid test" was originally the nitric acid test for gold, because only gold would not dissolve in it.
To acidify is to add acid to usually a solution.
Acidic substance have pH < 7.
An acid salt is formed by an acid where incomplete exchange of replaceable hydrogen occurs.
An acid dye is a dye that is metallic salt of an acid and c an be applied in an acid medium.
7.9.2 Aerobic
(Greek aēr air, bios live)
A bacterial processes that occurs only in the presence of oxygen is an aerobic process, the opposite is an anaerobic process.
7.9.3 Aliquot
An aliquot is a portion, a known fraction of the whole sample.
7.9.4.2 Allotropes, sulfur, carbon
See diagram 12.18.1: Sulfur crystals.
If the same element may have different bonding structure and crystal structure and so has different properties within, each variation of structure and properties is called an allotrope.
Sulfur has 3 allotropes, rhombic or α sulfur, monoclinic or β sulfur, and plastic sulfur.
Carbon has different bonding and crystal structure in the allotropes graphite and diamond.
Graphite is slippery, because weak van der Waals' forces between the flat layers allow them to slide over each other.
So graphite is used as a dry lubricant in machinery and in "lead" pencils.
Diamond has very high melting point and is extremely hard, because of strong chemical bonds holding the carbon atoms into a rigid three-dimensional structure, a network solid with a tetrahedral arrangement, making it chemically unreactive.
However, diamond is a poor conductor of electricity, because the electrons are held in relatively fixed positions around the carbon atoms.
7.9.6 Amphoteric
Can act as an acid or a base, e.g. water, bicarbonate ion.
Amphoteric oxides react with both acids and bases, e.g. Al2O3, PbO, SnO ZnO.
Their hydroxides are also amphoteric.
7.9.6.2 Aqueous solutions, states of matter
The states of matter are solid (s), liquid (l), gas (g), aqueous solution (dissolved in water) (aq).
An aqueous solution is a solution in water.
In this document, "solution" is always an aqueous solution, unless otherwise specified.
So a sugar solution contains sugar dissolved in water, but a solution of liquid sucrose contains no water.
7.9.7 Constant boiling mixture, azeotrope
An azeotrope is a mixture of liquids that has a constant boiling point, because the vapour has the same composition as the liquid mixture, so the components of the solution cannot be separated by distillation.
The boiling point of an azeotropic mixture may be higher or lower than any component.
An example of a constant boiling mixture is 4.37% water and 95.63% ethanol.
The mixture boils at 78.2oC, but pure ethanol boils at 78.4oC and water boils at 100oC.
7.9.8 Base
A base or an alkali is a good electrolyte, that turns red litmus blue and has a slippery feel.
Bases with water form hydroxide ions, OH-.
Bases react with hydrogen ions, H+.
A base is a proton acceptor (H+) (Bronsted-Lowry definition).
7.9.9 Bessemer process
The Bessemer process, invented by Sir Henry Bessemer, 1813-1898, converts pig iron from a blast furnace into steel by blowing.
air or pure oxygen into the molten impure metal to convert impurities into a separating slag.
Oxygen is blown through the molten iron to react with carbon in the iron and remove it as carbon dioxide gas.
The reaction which gives off a lot of heat to raise and maintain the temperature of the molten steel.
The process was changed to first removing all the carbon from the melt, then introducing 1 per cent of carbon to make the steel.
7.9.10 Borosilicate glass, Pyrex
Addition of borate allows the formation of a glass that melts at a lower temperate than silica, and expands less on heating than soda glass, as well as more plastic over a wider temperature range, e.g. Pyrex and glass wool.
So borosilicate glass has a very low coefficient of thermal expansion and a softening temperature above 800oC.
The composition may be 70% silica, 10% boron oxide and some sodium, potassium and calcium oxides.
The chemical composition of the Pyrex used in laboratory glassware may be different from the Pyrex used in kitchenware.
Mixing bowls, Pyrex, glass 1 litre.
Casserole dish, freezer and dishwasher safe, 2 litre.
7.9.11 Buffers
A mixture of substances that tend to hinder large changes in acid or basic properties of a solution.
The term "buffer" is used in a more general sense outside chemistry.
The pH of a buffer solution is not greatly changed by the addition of an acid or an alkali.
Most buffer solutions are a mixture of a weak acid or base with one of its salts.
In body fluids, the buffers include H2CO3 with HCO3-.
Acidic buffer, e.g. sodium hydrogen carbonate with carbonic acid solutions, the salt of the weak acid is completely dissociated into ions.
However, the weak acid is only partly dissociated.
Basic buffer, e.g. ammonium chloride in ammonia solutions.
7.9.12 Anti-bumping granules, boiling chips
Anti-bumping granules, boiling chips (- ceramic, silicon carbide, fused alumina)
Boiling chips, usually fused alumina, also flower pot bits, prevents large bubbles of gas forming that could cause explosive emissions from a beaker containing a heated solution.
Sudden formation of a large amount of vapour from the bottom of a heated vessel of liquid, rather than the usual controlled boiling.
So boiling chips (anti-bumping granules), are added to chemical reactions to keep the bubbles small and aid steady boiling.
Hydrocarbons with longer chains will have higher boiling points than similar hydrocarbons with branched chains, because they have more van der Waals intermolecular bonds between one molecule and another molecule.
7.9.13 Catalyst
A catalyst is an agent that speeds up a chemical reaction without itself being used up in the process, e.g. the transition metals Co, Ni, Pt.
Enzymes are catalysts for biological reactions.
7.9.14.1 Chemical potential energy, enthalpy
1. The chemical potential energy stored in a substance is called the heat content or enthalpy.
In a chemical reaction, chemical bonds are broken in the reactants (energy is absorbed), and formed in the products, (energy is released).
The energy is measured in joules, J or more commonly in kilojoules, kJ, where 1000 J = 1 kJ.
2. In a chemical reaction energy is neither created nor destroyed, the law of conservation of energy and the First Law of Thermodynamics.
In an endothermic reaction, the amount of energy absorbed when chemical bonds are broken is greater than the amount of energy released when chemical bonds are formed.
In an exothermic reaction, the amount of energy absorbed when chemical bonds are broken is less than the amount of energy released when chemical bonds are formed.
3. The heat of reaction, δH, is the heat change for the reaction and is measured in a calorimeter.
If δH is negative, the reaction is endothermic.
The reaction releases heat energy and the container feels hotter.
If δH is positive the reaction is endothermic.
The reaction takes in heat energy and the container feels cooler.
Energy is usually measured at 100 kPa and 298 K.
4.1.0 Enthalpy of reaction, heat of reaction
Energy from chemical reactions
The heat of reaction, DH (δH) is the heat change for a reaction in kJ per mol of reactant or product.
This is also called the enthalpy of reaction.
For endothermic reactions, DH (δH) is positive.
For exothermic reactions, DH (δH) is negative.
A --> B + xkJ
i.e. A --> B, DH is -xkJ / mol A
Be careful! The reactions may be vigorous!.
7.9.14.3 Clathrate
Inclusion compound where the "guest molecule" is in a lattice cage formed by the host molecule, e.g. Dianin's compound, 4-p-hydroxyphenyl-2, 2, 4-trimethylchroman.
7.9.15 Continuous phase /outer phase
The continuous phase is the continuous "outside" liquid that surrounds a second liquid, its droplets being discontinuous, in an emulsion.
7.9.16 Detergent, SYNDET, Synthetic Detergent
A detergent is a synthetic surfactant, not a soap, the sodium salts of natural fats.
A detergent has the cleansing properties of a soap, but it does not combine with any salts present as soap does in hard water.
So detergents hold dirt in suspension.
7.9.19 Enzymes
Enzymes are biological catalysts that can speed up, and control, chemical reactions that would otherwise virtually never occur at normal body temperature, 37oC.
Thousands of chemical reactions are occurring in the human body every moment of life, and each of these reactions is controlled by a particular enzyme, e.g. catalase breaks down potentially poisonous hydrogen peroxide into water and oxygen.
Enzymes are proteins that are specific in their action, but are not altered by the reaction, so they can be used repeatedly.
An enzymes is neither a reactants nor a product of a chemical equation it influences.
As with all proteins, enzymes are destroyed by heat or by extreme values pH or salt concentration, denaturation.
So enzymes are sensitive to pH.
7.9.20 Equilibrium
An equilibrium occurs in reactions in which the forward and reverse rates are matched so that the composition of the mixture appears unchanging in time.
The symbol for equilibrium in this document is <--> or <=>.
For example, if nitrogen dioxide is placed in a closed flask, some of it changes to dinitrgen tetroxide.
2NO2 (g) <=> N2O4 (g) nitrogen dioxide <=> dinitrgen tetroxide.
7.9.20.1 Etchants
Etching may be done in schools to prepare printed circuit boards, metal specimens for examination, etchings and lithographic plates.
However, many etchants are hazardous.
1. Ammonium persulfate (NH4)2S2O8, a strong oxidizing agent, is used as an etchant for copper plates as a 20% solution (w / v), (20 g in 100 mL of water), and 5% solution (w / v) (5g in 100 mL of water), prepared before the lesson, and used at 80oC, to prepare copper and alloys for microscopic examination, etch copper in construction of printed circuit boards and prepare tin coating on steel for microscopic examination.
2. Copper ammonium chloride solution, [Cu(NH3)3Cl2], 10% solution (w /v) (10 g per 100 mL of water), to prepare steels for macroscopic examination.
3. Iron (III) chloride (ferric chloride), 20% solution (w / v) (20 g in 100 mL of water), prepared before the lesson, is an etchant for most metals and alloys, and is used etch copper in PCB construction, and etch aluminium, zinc and copper plates, but it may leave persistent stains.
4. Nitric acid, 25% solution (approximately 4M), is an etchant for copper plates.
5. Sodium hydroxide solution, 1% solution (w / v) (1 g per 100 mL of water), to prepare aluminium and alloys for microscopic examination.
6. Sulfuric acid with concentration < 2M, may be used with care.
7. The following etchant chemicals or mixtures containing them are not permitted in schools, because of high corrosive risk and, in some cases, potential fire hazard:
7.1. Ammonium hydroxide + hydrogen peroxide,
7.2. Chromic acid,
7.3. Hydrofluoric acid,
7.4. Nitric acid + potassium dichromate + water, which can produce nitric acid,
7.5 Nitric acid + methanol,
7.6 (Nital), Nitric acid + glycerol + acetic acid,
7.7. Picric acid.
7.9.22 Flammable
1. The word "flammable" means "easily set on fire".
Also, you can use "non-flammable", but in chemistry do not use "inflammable".
Flammability, explosion, limits: outer limits for the ratio of fuel to air within which the mixture will burn.
The mishandling of flammable solvents has probably caused fires and personal injuries in chemical laboratories, especially the burning of loose long hair.
Staff and students must have securely fixed and contained hair by tying back the hair or using caps or hair nets.
Flammable solvents become more difficult to ignite as their boiling points rise, so use the highest boiling point solvent possible.
Do not use water baths to heat volatile flammable solvents.
Solvents should only be used by staff with students after assessment of the risks, which include not only flammability, but their toxicity, including possible allergic reactions.
2. Carbon disulfide has a greater flammability than ether and forms more dense vapours, with a low ignition temperature < 100oC.
Carbon disulfide is not permitted in schools.
3. Diethyl ether evaporates readily to form a heavy vapour in air, which can travel along the bench or floor in an air current.
Diethyl ether is not permitted in schools.
4. Hydrogen forms violently explosive mixtures with air in almost any proportions and spontaneously combusts at concentrations greater than 4% in air.
Use this gas for demonstration purposes only in extremely small quantities or use soap bubble techniques.
5. Natural gas forms explosive mixtures with air so turn off heaters, Bunsen burners, and other equipment using natural gas and other flammable gases, e.g. acetylene.
6. Methylated spirit, ethanol and hydrocarbon solvents, e.g. petroleum spirit, hexane, pose the greatest risk in schools.
Mixtures of air with any of these materials are highly flammable, and ignition of vapour is usually followed by a fire in or around the solvent container.
7.9.24 Flocculent
A flocculent is usually a precipitate in cloud-like tufts, flocs.
In bacteriology, flocculation refers to the formation of floccules (agglutinated bacteria), in a precipitin test, especially for antigens of Salmonella.
In the mining industry, flocculation refers to coagulation of ore particles to form flocs and remove excess water.
7.9.26 Flux
A flux is a substance added to lower the melting temperature in metallurgy and soldering.
The fusion of metals to form alloys is often done under a flux that may promote liquefaction, prevent volatilization and unnecessary exposure to the air.
7.9.27 Froth flotation
The mining industry uses froth flotation to adsorb chemicals on solid particles along with a foam to preferentially float off certain minerals and leave others behind.
7.9.28 Fuel cell
1. A fuel cell is a device with a cathode and anode, which converts a fuel directly into electricity without burning.
The simplest case is hydrogen gas bubbled over a porous sintered nickel anode in alkali solution, while oxygen is bubbled over a similar cathode separated by a porous membrane.
2H2 (g) + O2 (g) --> 2H2O (l)
An electric current is produced in an external circuit.
A fuel cell is like a battery, except that fuels, e.g. methanol, rather than metals are consumed, and the reaction is not reversible.
Fuel cells need a continuous source of fuel and oxygen from the air.
However, in a battery, the energy comes from chemicals already stored in the battery.
2. The two types of fuel cell have the same reactants and products, and hydrogen is ionised at the anode.
The differences are the different movement of ions across the electrolyte, (indicative of the relevant H+ and OH- concentrations).
In acidic cells, the water product exits on the same side as the oxygen, (indicative of H+ moving across the electolyte).
In alkaline cells, the water product exits on the same side as the hydrogen, indicative of the Oxygen, (as hydroxide). moving across the electrolyte.
7.9.29 Galvanize
(Luigi Galvani 1737-1798)
To galvanize to cover metal by electrodeposition of zinc.
A common roofing material is galvanized iron.
7.9.30 Group formula
A group formula places atoms together in groups that correspond to the grouping in the actual molecule, e.g. aspirin, CH3CO.O.C6H4COOH.
7.9.31 Heavy metals
1. Heavy metals, metals of high density, specific gravity > 5, high relative atomic weight (atomic mass), especially if poisonous.
The term "heavy metals" is used in legislation related to chemical hazards and safe use of chemicals.
Heavy metals defined as elements commonly used in industry and toxic to animals and to aerobic and anaerobic processes.
Heavy metals may include As, Cd, Cr, Cu, Pb, Hg, Ni, Se, Zn.
The following are called "heavy metals", if they cause pollution: Copper, Lead, Mercury, Zinc.
The term "heavy metal" is not exact, for example, Aluminium and Beryllium are toxic, but they are not called "heavy metals".
Recycle heavy metals safety
From New South Wales Department of Education publication "Chemical safety in schools":
Not all heavy metal salts are toxic, but most are, i.e. arsenic, lead, cadmium and mercury compounds are extremely hazardous.
Heavy metals are toxic by ingestion, by inhalation of dusts or fumes, and by skin contact especially with chromates, which can cause dermatitis.
It is important to use safe working practices to prevent student contact with heavy metal salts.
Some heavy metal salts are powerful oxidants and may also present a reactive hazard.
Avoid or control contact of heavy metals with reducing agents.
Collect solution wastes of heavy metals, and reduce volume by allowing solvent (usually water), to evaporate.
Solid wastes of heavy metals should be recycled if practicable.
Insoluble heavy metal salts, e.g. barium sulfate, may be placed in garbage.
Very often chemical disposal can be incorporated into student activity to demonstrate a chemical process.
For example, recrystallization of copper sulfate or displacement of copper from solution using steel wool.
Heavy metals recycling
Solutions of the following heavy metals may be treated for recycling:
Antimony, barium, beryllium, chromium, cobalt, lead, manganese, molybdenum, nickel, selenium, strontium, tellurium, tin, vanadium, zinc.
Experiment
1. Precipitate insoluble metal salts with sodium carbonate, sodium hydroxide or sodium sulfide.
Decant the clear solution above the precipitate and wash it down the sink.
Store the dried precipitate in a waste disposal bottle or use a displacement reaction to recover the elemental metal.
2. Copper may be displaced from solution by adding steel wool to precipitate the copper.
Decant the clear solution above the precipitate and wash it down the sink.
Store the dried precipitate.
7.9.32 Hydronium ion, (hydroxonium ion, oxonium ion)
An hydronium ion H3O+ is formed when acids dissociate in water.
Water molecules are very polar.
Electrons tend to “build up” on the oxygen side (oxygen atoms are very electronegative), leaving the H-atom-side of the molecule relatively positive.
(H atoms are poorly electronegative).
If an H+ ion comes close enough to a water molecule, it will be attracted to the oxygen side of a water molecule, yielding an hydronium ion.
The rate of hydronium and hydroxyl ions forming water molecules is large, so few hydronium ions exist in water, 10-7 moles/litre.
7.9.33 Hydrophilic, hydrophobic
"Magic Sand" (Hydrophobic Sand), invisible in water, absorbent hydrophilic polymer (toy product).
The hydrophilic substances are "water-loving", polar molecule materials that mix with water.
Hydrophilic substancesare attracted to water and may dissolve in water to form hydrogen bonds, e.g. glucose, sugars.
They have an affinity for water and are readily absorbed or wetted by water.
Hydrophilic colloids readily form hydrosols or remain as hydrosols.
The hydrophobic substances are oily "water-hating", non-polar molecule materials that do not mix with water or repel water, e.g. oils, proteins, greases, clays.
Hydrophobic colloids do not form or remain as hydrosols.
Hydrophobia is the aversion to water of a person suffering from the disease rabies.
Soap molecules have one end polar and the other end non-polar so they can attach to oils yet dissolve in water.
7.9.35 Inhibit
To inhibit is to slow down a chemical reaction by blocking a part of the mechanism.
7.9.36 Labile
A labile substance is unstable and liable to change to another form or to move away.
7.9.37 Martensite
A martensite forms when a solid solution of carbon in iron forms on rapid cooling.
It is responsible for the hardness of quenched steel.
7.9.39 Molecular mass
The molecular mass, formerly molecular weight, is the mass of one mole of that material.
7.9.40 Napalm
Petrol gelled with the aluminium salts of naphthalene acids and palmitic acids, used for warfare in flame throwers and napalm bombs.
7.9.41 Petroleum fraction
A fraction of oil selected in a refinery distillation process on the basis of boiling point.
7.9.41.1 Petroleum spirit
Petroleum spirit, petroleum ether ACS reagent, ligroin, benzine (not "benzene"), petroleum fraction 60-150oC.
It is mixture of mostly pentanes and hexanes, petroleum benzine, petroleum solvent, 40oC-60oC, light petroleum, petroleum solvent, Highly flammable, Toxic by all routes, so do not inhale vapour.
Petroleum spirit, use higher boiling point fractions to avoid benzene, thinner, solvent in non-polar chromatography.
Highly flammable liquid and vapour, may be fatal if swallowed and enters airways, and may cause genetic defects and cancer.
Low cost, but substitute toluene or xylene from hardware store.
("benzine", in some countries = petrol).
7.9.42 Photolysis
A photolysis is chemical reaction brought about by light including ultra-violet light.
A radiolysis is the equivalent when radioactive emissions are involved.
7.9.46 Radical
A radical is a group of atoms that behaves like a single atom in a chemical reaction, e.g. the ammonium radical, NH4+.
A free radical has an unpaired single electron, e.g. the methyl radical, CH3-.
7.9.47 Sequester
To sequester is to take out of circulation, to tie up metal ions so that they do not interfere, to form a stable chelate complex or biochemical complex with an ion to remove it from solution or otherwise make it unreactive, e.g. by precipitating soaps.
Sequestering agents are used to link undesirable metal ions together to form a stable structure that does not readily decompose and limits the metal ions’ ability to react with other ions, clays or polymers.
Sequestering agents are commonly used for water treatment purposes to reduce water hardness.
They combine with calcium, magnesium and other heavy metal ions in hard water to form molecules in which the ions are held so securely (sequestered) that they can no longer react.
A sequestering agent surrounds another molecule or atom and holds it "in seclusion." so that it " hides" the molecule or atom and prevents it from entering into chemical reactions.
7.9.48 Solute
The solute is the dissolved material in a solution or the material to be dissolved.
The solute is the minor component in a solution.
7.9.49 Solvent
The solvent, usually a liquid, is the dissolving material in a solution or the liquid to dissolve the solute.
The solvent is the main component of a solution.
7.9.50 Spontaneous process
A spontaneous process that has the potential to occur on its own without further input of energy.
However, it may occur so slowly it is not measurable.
Diamond will eventually decompose to graphite.
7.9.51 Substrate
A substrate basis on which something else is placed, a starting material.
It is usually the underlying layer on which another substance reacts with and / or is deposited.
Also, an enzyme act on a substrate to facilitate a chemical change in it.
7.9.51.1 Sulfates, battery sulfation
Salt or ester of sulfuric acid, sulfuric (IV) acid, contain ion SO42-, normal and acid salts, organic sulfates R2SO4.
Sulfate ion is not toxic.
Toxicity depends on the cation present, especially heavy metals, e.g. leads in lead sulfate.
When sulfates are heated to decomposition, toxic sulfur dioxide gas forms.
Metal sulfates are used to study the following:
1. the reactivity series of metals,
2. the properties of the sulfate ion,
3. the heating to decomposition and
4. redox reactions, including the formation of precipitates.
Sulfate solutions, if not containing heavy metals, may be directly discharged to the sink if solution pH = 8-10.
7.9.51.1a Battery sulfation
Sulfation on plates of a lead-acid battery, secondary cell, causes reduced efficiency, because PbSO4 forms.
Sulfation is the effect of overcharging or other environmental conditions in lead acid batteries that leads to permanent capacity loss.
Flakes of lead sulfate break away from the plates and fall to the bottom of the cell, where they can no longer react and produce energy.
7.9.51.2 Sulfides
1. Sulfur with more electropositive element to form an inorganic compound, sulfur with metal by direct combination, sulfur with nonmetals to form covalent compounds, e.g. hydrogen sulfide H2S, salts of hydrogen sulfide are ionic sulfides containing S2- ion, organic thioethers containing two hydrocarbon groups, RSR (R is not H), e.g. diallyl sulfide with garlic smell, and dimethyl sulfide CH3SCH3.
The sulfide zone of a sulfide mineral lode contains unaltered sulfide mineral and leached sulfides from above.
Metal sulfides may react violently with oxidizing agents.
When heated to decomposition they produce toxic fumes of sulfur compounds.
They react with acid or water to produce toxic hydrogen sulfide gas.
Alkaline sulfides, i.e. calcium, sodium, ammonium and potassium sulfides, behave like alkalis, may cause softening and irritation of the skin.
Sulfides of heavy metals are usually insoluble and have low toxicity.
However, they may react with acids to release hydrogen sulfide gas.
2. Metal sulfides are used to study the following:
1. reactivity series of metals,
2. properties of the sulfide ion,
3. reactions with acids,
4. reactions with water,
5. the heating to decomposition,
6. formation of precipitates in redox reactions.
Sulfide solutions, if not containing heavy metals, may be discharged down the sink after treatment to be made acidic in a fume cupboard to expel hydrogen sulfide gas.
After completion of the reaction the solution should be made basic with sodium hydroxide to pH: 8-10, then discharged down the sink.
7.9.51.3 Sulfites
Salt or ester from sulfurous acid, sulfuric (IV) acid, with reducing properties, normal sulfites and acid sulfites, bisulfites, salts contain SO32-.
Bisulfites are used to digest wood pulp.
All sulfites are toxic if ingested.
Most metal sulfites irritate the stomach by production of sulfurous acid.
When sulfites are heated to decomposition, toxic sulfur dioxide forms.
Sulfite solutions, if not containing heavy metals, may be discharged down the sink after treatment.
Solutions should be made acidic in a fume cupboard to expel sulfur dioxide gas.
At completion of the reaction the solution should be made basic with sodium hydroxide to pH: 8-10, then discharged down the sink.
7.9.52 Surfactants
A surfactant is molecule attracted to the surface of water and capable of changing the properties of the surface, generally by lowering the surface tension to make a solution more wettable.
7.9.53 Synergism
A synergism occurs when two or more substances together produce an effect that is greater than the sum of the individual separate effects.
In pharmacology, the combined action a two drugs administered together may be greater than action of the combined action of the two drugs administered separately at different times.
7.9.54.2 Urethane
Urethane, ethyl carbamate ester, carbamic acid ethyl ester, C3H7NO2, NH2COOC2H5, CO(NH2)OC2H5, toxic, suspected carcinogen, veterinary anaesthetic, colourless, odourless, emits toxic fumes of nitrogen oxides when heated, used to make pesticides and pharmaceuticals, prepare polyurethane plastic, contaminant of fermented alcoholic beverages.
7.9.55.1 Waste chemicals bottles
Recommended waste chemicals bottles include:
Waste copper residues
Waste halogenated organic chemicals
Waste heavy metal mixtures
Waste non-halogenated organic liquids
Waste lead residues
Waste mercury residues
Waste silver residues
Waste zinc residues
In some countries, waste chemicals bottles are collected periodically by government or contractors.
7.9.56.1 Xanthene dyes
See diagram: Fluorescein.
From condensation of phthalic anhydride with resorcinol, have xanthene nucleus, e.g. fluorescein.
7.9.57 Zymase
Zymase is the mixture of enzymes, enzyme complex, in yeasts that catalyses the breakdown of sugar in alcoholic fermentation into alcohol and carbon dioxide.
C6H12O6 + yeast --> 2C2H5OH + 2 CO2
glucose + yeast --> ethanol + carbon dioxide>
The glucose combines with the yeast in an anerobic reaction to produce ethanol and carbon dioxide.
Zymase is an enzyme complex that catalyses glycolysis, the fermentation of sugar into ethanol and carbon dioxide.
8.0.0 Direct union of elements to form compounds
1. To form acids
H2 + Cl2 --> 2HCl
H2 + S --> H2S.
2. To form salts
2Na + Cl2 --> 2NaCl
Fe + S --> FeS
8Fe + S8 --> 8FeS (ferrimagnetic iron (II) sulfide)
7.1.6.2 Breakdown starch to sugars
Salivary amylase enzyme breaks down starch into the reducing sugars (+) glucose and maltose.
Reducing sugars do not react with iodine solution and starch does not react with Fehling's solution.
The sugars reduce copper (II) in Fehling's solution to brick-red copper (I) oxide.
1. Prepare a clear solution of laundry starch by adding a mixture of 1g starch in 10 ml of water to 500 mL of boiling water
Leave the solution to cool to room temperature.
2. Put 10 mL of dilute starch solution into a test-tube.
Add to this 1 mL of saliva and stir this into the starch solution.
Record the time of adding the saliva.
After 2 minutes use a dropper to put 2 drops of the solution on a white tile.
At 5 minute intervals, remove three drops with a dropper and put them on a clean white tile, taking care to keep them from running into each other.
The dropper must be washed between each test.
3. To test for starch, add iodine solution and note the intensity of the blue black colour.
The decreasing intensity of the blue colour shows the decreasing amount of starch.
4. To test for increasing amounts of sugar, put three drops of the reaction mixture into a small test-tube.
Add Fehling's No. 1 and No. 2 solutions and heat this mixture almost to boiling point.
Note the intensity of the brick-red colour increasing with time.
Repeat the experiment every 2 minutes with clean droppers.
Note the decreasing intensity of the blue colour that shows that starch is being used up.
Keep doing the test until it shows that there is more sugar after boiling.
7.1.6.3 Breakdown ethanol to ethene (ethylene)
See diagram 3.2.96: Breakdown of ethanol.
Push cotton wool soaked in methylated spirit to the bottom of a hard glass test-tube.
Pack small pieces of porous pot, unglazed porcelain, in the middle of the test-tube.
Fit a delivery tube to collect ethene gas over water in a receiving test-tube.
With the hard glass test-tube in a horizontal position, heat the porous pot strongly, then gently heat the cotton wool to produce ethanol vapour.
The ethanol vapour breaks down over the hot porous pot to produce ethene gas and water vapour.
Ethene is insoluble in water and collects in the receiving test-tube.
Collect three receiving test-tubes full of ethene then immediately disconnect the delivery tube when you stop heating to avoid a suck back of water on the hot porous pot.
In test-tube 1, burn ethene with a lighted taper.
Shake test-tube 2 with drops of dilute potassium manganate (VII) solution and sodium carbonate solution.
The colour disappears.
Shake test-tube 3 with bromine water.
The colour disappears.
C2H5OH (l) --> C2H4 (g) + H2O