Volumetric Analysis 1
To make a standard solution of sodium carbonate
Introduction
A standard solution is one whose concentration is known exactly. Standard solutions of liquids, for example acids, are easy to prepare and are usually supplied. Standard solutions of solids can be prepared by weighing a mass of solid, and dissolving it in a known volume of solution in a volumetric flask. Today, you are going to prepare a standard solution of sodium carbonate to use later in another practical.
Apparatus
Goggles
Bench mat
100 cm3 beaker
250 cm3 beaker
250 cm3 volumetric flask with stopper
Filter funnel
Glass rod
Teat pipette
Spatula
Label
De-ionised water
Anhydrous sodium carbonate, Na2CO3(s) Copyright © Nigel Saunders N-ch1-35
Methods
Read through the Methods. Make a suitable blank results table, complete with units in the headings.
1. Using the ± 0.1g balance, weigh approximately between 1.2g and 1.4g of sodium carbonate into the small beaker. Do not record the mass.
2. Using the ± 0.01g balance, weigh the small beaker and its contents accurately.
Record this mass.
3. Transfer the contents of the small beaker into the large beaker.
Weigh the small beaker again using the ± 0.01g balance.
Record this mass.
The difference between the two accurate masses is the mass of sodium carbonate in your beaker.
4. Add de-ionised water cautiously down the side of the large beaker.
Use about 150cm3 of water, and swirl the beaker to mix the contents.
5. Stir using a glass rod to dissolve the solid completely.
6. Transfer the solution into the volumetric flask using the funnel.
Remember: pour down the glass rod;
remove the last drop of solution from the glass rod onto the funnel.
Wash the beaker, rod and funnel several times using de-ionised water from the wash bottle, letting the washings go into the flask.
7. Make up to the mark on the volumetric flask with de-ionised water.
Stopper firmly, and shake the flask thoroughly to mix the contents.
8. Label the flask clearly with your name, the date, and the contents of the flask. Copyright © 2003 N. Saunders N-ch1-35
Volumetric Analysis 1
To make a standard solution of sodium carbonate:
Technician's Notes
Prior to practical
Sodium carbonate
Heat required amount of sodium carbonate (Na2CO3) to drive off water of crystallisation.
Either: heat in an evaporating dish over a Bunsen burner for 30 minutes approx., or
heat in a drying oven at about 110°C for 1 hour.
Agitate the solid periodically with a clean glass rod.
Transfer to a desiccator after heating, and label it "sodium carbonate".
Care: Use tongs and eye protection.
Beware of hot solid and apparatus.
Sodium carbonate forms caustic alkaline solutions with water; if spilt on skin wash with plenty of water.
Analytical balances
Please check cleanliness and correct functioning of analytical balances.
De-ionised water
Please check 6th Form wash bottles are clean and filled with de-ionised water.
Make sure that additional de-ionised water is available in the aspirator.
Requirements per class
Sodium carbonate solid (see above). Minimum of 3g per student approx.
Analytical balances (see above).
Top pan digital balances (minimum of two if possible).
De-ionised water (see above).
Requirements per student
1 x 250cm3 beaker (dry)
1 x 250cm3 volumetric flask with stopper to fit
1 x glass funnel (check that it will enter the neck of the volumetric flask easily)
1 x glass rod
1 x weighing bottle with lid
1 x 6th Form wash bottle containing de-ionised water
1 x teat pipette
1 x small spatula
1 x self-adhesive label
Tuesday, April 21, 2009
Estimating Errors in Chemistry
Copyright © 2003 Nigel Saunders N-ch1-24 Estimating Errors in Chemistry
Introduction
You must be able to estimate the size and importance of errors in practical work to gain full marks in your practical assessments. The table summarises what you need to do about errors in each Skill:
Skill
You must be able to:
Skill P
Planning
Organise the procedure to be followed, selecting appropriate techniques, reagents and apparatus, with due regard to precision of measurement and scale of working.
Skill I
Implementing
Make and record measurements to a degree of precision allowed by the apparatus used.
Skill A
Analysing
Identify sources of error, and recognise the limitations of experimental measurements.
Skill E
Evaluating
Assess the reliability of your data and the conclusions drawn from them, taking into account the errors in the data obtained.
Introduction
You must be able to estimate the size and importance of errors in practical work to gain full marks in your practical assessments. The table summarises what you need to do about errors in each Skill:
Skill
You must be able to:
Skill P
Planning
Organise the procedure to be followed, selecting appropriate techniques, reagents and apparatus, with due regard to precision of measurement and scale of working.
Skill I
Implementing
Make and record measurements to a degree of precision allowed by the apparatus used.
Skill A
Analysing
Identify sources of error, and recognise the limitations of experimental measurements.
Skill E
Evaluating
Assess the reliability of your data and the conclusions drawn from them, taking into account the errors in the data obtained.
Making Standard Solutions
Making Standard Solutions
What is a standard solution?
A standard solution is a solution whose concentration is known accurately. Its concentration is usually given in mol dm–3. When making up a standard solution it is important that the correct mass of substance is accurately measured. It is also important that all of this is successfully transferred to the volumetric flask used to make up the solution. The following procedure will make sure that this happens.
Background calculations
1. Work out the number of moles needed to make up a solution with the required volume and concentration.
Show your working in the space below.
2. Now work out the relative formula mass, Mr, of the chosen substance. Show your working in the space below.
3. Work out the mass of substance needed using your answers from steps 1 and 2. Show your working in the space below.
Making up the solution
• Take a watch glass and place it on the balance. Tare the balance (set it to zero).
Carefully weigh out the required mass of substance.
• Transfer this amount to a beaker. Add water from a wash bottle to dissolve it.
Use some of the water to rinse all the substance off the watch glass. Do this at least twice.
• Stir with a glass rod until all the solid is dissolved, then transfer the solution to the volumetric flask.
Use more water from the wash bottle to rinse out the beaker and the glass rod. Do this at least twice.
• Add water to just below the line on the volumetric flask.
Add the final drops with a teat pipette to ensure that the bottom of the meniscus is on the line.
• Put the lid on the flask and turn the flask over a couple of times to mix the solution.
• Label your solution with your name, the date, and the contents, e.g. 2.0M NaCl. Then tidy up!
moles = concentration x volume
concentration is in mol dm–3 … and …
volume is in dm3, so if the volume is given in cm3, divide it by 1000 to get dm3
To find the Mr of carbon dioxide, CO2:
CO2 has … 1 carbon atom … 1 x 12 = 12
2 oxygen atoms … 2 x 16 = 32
add together … = 44
It helps to remember that:
"mass is mister mole", or
mass = Mr x mole Copyright © 2003 Nigel Saunders & Sally Burch N-ch1-49
What is a standard solution?
A standard solution is a solution whose concentration is known accurately. Its concentration is usually given in mol dm–3. When making up a standard solution it is important that the correct mass of substance is accurately measured. It is also important that all of this is successfully transferred to the volumetric flask used to make up the solution. The following procedure will make sure that this happens.
Background calculations
1. Work out the number of moles needed to make up a solution with the required volume and concentration.
Show your working in the space below.
2. Now work out the relative formula mass, Mr, of the chosen substance. Show your working in the space below.
3. Work out the mass of substance needed using your answers from steps 1 and 2. Show your working in the space below.
Making up the solution
• Take a watch glass and place it on the balance. Tare the balance (set it to zero).
Carefully weigh out the required mass of substance.
• Transfer this amount to a beaker. Add water from a wash bottle to dissolve it.
Use some of the water to rinse all the substance off the watch glass. Do this at least twice.
• Stir with a glass rod until all the solid is dissolved, then transfer the solution to the volumetric flask.
Use more water from the wash bottle to rinse out the beaker and the glass rod. Do this at least twice.
• Add water to just below the line on the volumetric flask.
Add the final drops with a teat pipette to ensure that the bottom of the meniscus is on the line.
• Put the lid on the flask and turn the flask over a couple of times to mix the solution.
• Label your solution with your name, the date, and the contents, e.g. 2.0M NaCl. Then tidy up!
moles = concentration x volume
concentration is in mol dm–3 … and …
volume is in dm3, so if the volume is given in cm3, divide it by 1000 to get dm3
To find the Mr of carbon dioxide, CO2:
CO2 has … 1 carbon atom … 1 x 12 = 12
2 oxygen atoms … 2 x 16 = 32
add together … = 44
It helps to remember that:
"mass is mister mole", or
mass = Mr x mole Copyright © 2003 Nigel Saunders & Sally Burch N-ch1-49
Sunday, April 19, 2009
Anion analysis
Chloride Cl-
AgCl, PbCl2, Hg2Cl2 and CuCl are insoluble in water.
Concentrated sulphuric acid liberates steamy acidic fumes of HCl from solid chlorides:
NaCl(s) + H2SO4(l) " NaHSO4(s) + HCl(g)
Silver nitrate solution added to a solution of a chloride that has been acidified (test with blue litmus paper) with dilute nitric acid gives a white precipitate of silver chloride. The precipitate is readily soluble in dilute ammonia or in sodium thiosulphate solution:
Ag+(aq) + Cl-(aq) " AgCl(s)
AgCl(s) + 2NH3(aq) " [Ag(NH3)2]+(aq) + Cl-(aq)
AgCl(s) + 2S2O32-(aq) " [Ag(S2O3)2]3-(aq) + Cl-(aq)
Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates.
Concentrated solutions of sulphates can give a precipitate of silver sulphate in this test; its appearance is wholly different from AgCl. The latter is truly white; the sulphate is a pearly white, rather like pearlescent nail varnish.
Bromide Br-
AgBr, PbBr2, Hg2Br2 and CuBr are insoluble in water.
Concentrated sulphuric acid gives a mixture of hydrogen bromide, bromine and sulphur dioxide with solid bromides; the HBr produced is oxidised by sulphuric acid. The mixture evolves steamy brownish acidic fumes:
NaBr(s) + H2SO4(l) " NaHSO4(s) + HBr(g)
2HBr + H2SO4 " Br2 + SO2 + 2H2O
Silver nitrate solution added to a solution of a bromide that has been acidified (test with blue litmus paper) with dilute nitric acid gives a cream precipitate of silver bromide. The precipitate is readily soluble in concentrated ammonia:
Ag+(aq) + Br-(aq) " AgBr(s)
AgBr(s) + 2NH3(aq) " [Ag(NH3)2]+(aq) + Br-(aq)
Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates.
Oxidising agents oxidise bromide to bromine, which is yellow or orange in aqueous solution. Bromine can be extracted from the solution by shaking with an immiscible organic solvent, for example hexane, the organic layer then turning orange.
A suitable oxidising agent is sodium chlorate(I); this is added to the test solution, followed by a little dilute hydrochloric acid and a few cm3 of hexane:
OCl-(aq) + 2H+(aq) + 2Br-(aq) " Br2(aq) + Cl-(aq) + H2O(l)
Iodide I-
AgI, PbI2, Hg2I2 and CuI are insoluble in water.
Concentrated sulphuric acid gives a mixture of hydrogen iodide, iodine, hydrogen sulphide, sulphur and sulphur dioxide when added to solid iodides; the HI produced is oxidised by sulphuric acid. The mixture evolves purple acidic fumes, turns to a brown slurry, and is a mess:
NaI(s) + H2SO4(l) " NaHSO4(s) + HI(g)
2HI + H2SO4 " I2 + SO2 + 2H2O
6HI + H2SO4 " 3I2 + S + 4H2O
8HI + H2SO4 " 4I2 + H2S + 4H2O
There are no state symbols in these equations because the mixture is such a mess, and is more sulphuric acid than water.
Silver nitrate solution added to a solution of an iodide that has been acidified (test with blue litmus paper) with dilute nitric acid gives a yellow precipitate of silver iodide. The precipitate is insoluble even in concentrated ammonia:
Ag+(aq) + I-(aq) " AgI(s)
Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates.
Oxidising agents oxidise iodide to iodine, which is yellow or orange in aqueous solution. Iodine can be extracted from the solution by shaking with an immiscible organic solvent, for example hexane, the organic layer then turning purple.
A suitable oxidising agent is sodium chlorate(I); this is added to the test solution, followed by a little dilute hydrochloric acid and a few cm3 of hexane:
OCl-(aq) + 2H+(aq) + 2Ir-(aq) " I2(aq) + Cl-(aq) + H2O(l)
Lead ethanoate or lead nitrate solutions give a bright yellow precipitate of lead(II) iodide with iodides:
Pb2+(aq) + 2I-(aq) " PbI2(s)
The colour of lead(II) iodide comes from interactions in the lattice; dissolving the salt in boiling water gives a colourless solution which deposits glittering yellow plates on cooling.
Solutions of copper(II) salts give a brown mixture containing iodine and copper(I) iodide when added to solutions of iodides. Addition of sodium thiosulphate solution decolourises the iodine and leaves pinkish-cream copper(I) iodide as a precipitate.
2Cu2+(aq) + 4I-(aq) " 2CuI(s) + I2(aq)
2S2O32-(aq) + I2(aq) " 2I-(aq) + S4O62-(aq)
This reaction is the basis for the volumetric estimation of copper; the iodine liberated from a known amount of a copper(II) solution is titrated with standard sodium thiosulphate solution.
Sulphite SO32-
Sulphurous acid is considerably stronger than carbonic acid, so sulphites do not give the effervescence that is characteristic of carbonates when dilute acid is added.
Dilute hydrochloric acid on warming with a sulphite evolves sulphur dioxide; this turns acidified potassium dichromate(VI) solution (or paper) green:
SO32-(aq) + 2H+(aq) " H2O(l) + SO2(g)
Barium chloride solution gives a white precipitate of barium sulphite; addition of dilute hydrochloric acid causes the precipitate to dissolve without effervescence:
SO32-(aq) + Ba2+(aq) " BaSO3(s)
Sulphate SO42-
BaSO4, SrSO4 and PbSO4 are insoluble; CaSO4 is sparingly soluble.
Barium chloride solution added to the test solution acidified with dilute hydrochloric acid gives a white precipitate of barium sulphate:
Ba2+(aq) + SO42-(aq) " BaSO4(s)
HSO4- does the same thing with barium ions; however the original test solution would then be very acidic, so that should be tested for.
The addition of HCl destroys any carbonate or sulphite ions present so prevents the spurious positive result due to the precipitation of these barium salts. Barium nitrate solution can be used instead of barium chloride.
18.2 Lead ethanoate solution gives a precipitate of white lead sulphate:
Pb2+(aq) + SO42-(aq) " PbSO4(s)
Nitrate NO3-
Since all nitrates are water soluble, there is no precipitation reaction for this ion.
Solid nitrates decompose on heating; those of group 1 (except Li) give the nitrite and oxygen;
2NaNO3(s) " 2NaNO2(s) + O2(g)
All others give the metal oxide, nitrogen dioxide, and oxygen. A brown gas is emitted that re-lights a glowing splint:
2Pb(NO3)2(s) " 2PbO(s) + O2(g) + 2NO2(g)
Nitrate ions are reduced to ammonia by boiling with aluminium or with Devarda’s Alloy in sodium hydroxide solution. Devarda’s Alloy contains aluminium, zinc and copper. Since ammonium ions also give ammonia with NaOH, the test solution must be boiled with NaOH and the vapour tested for ammonia; if present heating must continue until all the ammonia has gone. The mixture is then cooled, Devarda’s Alloy (or a piece of aluminium foil) added, and the mixture re-heated. A gas that turns moist red litmus paper blue indicates nitrate in the original solution:
3NO3-(aq) + 8Al(s) + 18H2O(l) + 21 OH-(aq) " 8[Al(OH)6]3-(aq) + 3NH3(g)
Not an equation to be remembered!
Carbonate CO32-
Only the alkali metal and ammonium carbonates are water soluble. Some carbonates (e.g. zinc, copper(II)) are basic carbonates and contain a proportion of the hydroxide in their structure.
Heating decomposes all but the alkali and alkaline earth metal carbonates (at Bunsen temperatures) giving the oxide and carbon dioxide:
CuCO3(s) " CuO(s) + CO2(g)
Dilute hydrochloric acid gives vigorous effervescence with carbonates, evolving carbon dioxide:
CO32-(aq or s) + 2H+(aq) " H2O(l) + CO2(g)
Bicarbonates also give this effervescence. The reaction of carbonates with acid is exothermic; bicarbonates react endothermically.
Hydrogen carbonate (bicarbonate) HCO3-
Only the alkali metal and ammonium bicarbonates are obtainable as solids; group 2 bicarbonates exist only in solution.
Calcium chloride solution on addition to a bicarbonate solution gives no precipitate since calcium bicarbonate is soluble; this distinguishes it from carbonate, which does give a precipitate. On heating the calcium chloride/bicarbonate mixture a white precipitate appears since the bicarbonate decomposes to carbonate:
Ca2+(aq) + 2HCO3-(aq) " CaCO3(s) + CO2(g) + H2O(l)
Ethanoate CH3COO-
Ethanoates on heating with dilute hydrochloric acid give ethanoic acid, recognisable by its vinegary smell.
Neutral iron(III) chloride solution added to neutral solutions of ethanoate ion give a deep red colouration owing to formation of iron(III) ethanoate.
Ethanedioate C2O42-
Concentrated sulphuric acid added to a solid ethanedioate salt gives a mixture of carbon monoxide and carbon dioxide from dehydration of the ethanedioic acid formed:
HOOC-COOH " H2O + CO2 + CO
Potassium manganate(VII) solution acidified with dilute sulphuric acid added to a solution of an ethanedioate causes the purple colour to disappear:
2MnO4-(aq) + 5C2O42-(aq) + 16H+(aq) " 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
Calcium chloride solution added to a solution of an ethanedioate gives a white precipitate of calcium ethanedioate:
Ca2+(aq) + C2O42-(aq) " CaC2O4(s)
The precipitate dissolves readily in dilute hydrochloric acid.
Chromate(VI) CrO42- and dichromate(VI) Cr2O72-
These ions are related through the equilibrium
Cr2O72-(aq) + 2OH-(aq) D 2CrO42-(aq) + H2O(l)
In alkaline solution the yellow chromate(VI) dominates, in acidic solution orange dichromate(VI). All dichromate(VI) salts are soluble; addition of dichromate(VI) ions to solutions of ions of metals which have insoluble chromate(VI) salts leads to the precipitation of chromates. This means that the only dichromates that can exist are those of group 1 metals, ammonium, magnesium, calcium and strontium.
Barium chloride solution added to a chromate(VI) or dichromate(VI) solution precipitates bright yellow barium chromate(VI):
Ba2+(aq) + CrO42-(aq) " BaCrO4(s)
The addition of a heavy metal ion to potassium dichromate solution precipitates the chromate and therefore moves the equilibrium to the right hand side. If the chromate is sparingly soluble (e.g. strontium) the supernatant liquid will remain yellow. Very insoluble chromates, such as lead, remove all the colour from the supernatant liquid.
Dichromate(VI) ion solution in sulphuric acid is an oxidising agent; oxidation is shown by the solution turning from orange to green (Cr(III)). The following can be oxidised:
(a) iron(II) to iron(III):
Cr2O72-(aq) + 14H+(aq) + 6Fe2+(aq) " 2Cr3+(aq) + 7H2O(l) + 6Fe3+(aq)
(b) iodide to iodine (the solution turns murky greenish-brown):
Cr2O72-(aq) + 14H+(aq) + 6I-(aq) " 2Cr3+(aq) + 7H2O(aq) + 3I2(aq)
(c) sulphite to sulphate:
Cr2O72-(aq) + 8H+(aq) + 3SO32-(aq) " 2Cr3+(aq) + 4H2O(l) + 3SO42-(aq)
(d) nitrite to nitrate:
Cr2O72-(aq) + 8H+(aq) + 3NO2-(aq) " 2Cr3+(aq) + 4H2O(l) + 3NO3-(aq)
(e) hydrogen peroxide reacts with acidified dichromate(VI) solutions to give a blue compound that rapidly turns green and evolves oxygen. The blue compound can be extracted into an organic solvent such as butan-1-ol. The blue compound is CrO5, which contains a peroxy structure. It is covalent, and is stable in organic solvents though not in water.
Cr2O72-(aq) + 8H+(aq) + 3H2O2 (aq) " 2Cr3+(aq) + 7H2O(l) + 3 O2(g)
Alcohols are oxidised by acidified potassium dichromate(VI) solution. Primary alcohols give aldehydes and then acids, secondary alcohols give ketones. Ethanol can be used to test for dichromate(VI) ions, therefore, the solution turning green and the apple smell of ethanal being evident.
AgCl, PbCl2, Hg2Cl2 and CuCl are insoluble in water.
Concentrated sulphuric acid liberates steamy acidic fumes of HCl from solid chlorides:
NaCl(s) + H2SO4(l) " NaHSO4(s) + HCl(g)
Silver nitrate solution added to a solution of a chloride that has been acidified (test with blue litmus paper) with dilute nitric acid gives a white precipitate of silver chloride. The precipitate is readily soluble in dilute ammonia or in sodium thiosulphate solution:
Ag+(aq) + Cl-(aq) " AgCl(s)
AgCl(s) + 2NH3(aq) " [Ag(NH3)2]+(aq) + Cl-(aq)
AgCl(s) + 2S2O32-(aq) " [Ag(S2O3)2]3-(aq) + Cl-(aq)
Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates.
Concentrated solutions of sulphates can give a precipitate of silver sulphate in this test; its appearance is wholly different from AgCl. The latter is truly white; the sulphate is a pearly white, rather like pearlescent nail varnish.
Bromide Br-
AgBr, PbBr2, Hg2Br2 and CuBr are insoluble in water.
Concentrated sulphuric acid gives a mixture of hydrogen bromide, bromine and sulphur dioxide with solid bromides; the HBr produced is oxidised by sulphuric acid. The mixture evolves steamy brownish acidic fumes:
NaBr(s) + H2SO4(l) " NaHSO4(s) + HBr(g)
2HBr + H2SO4 " Br2 + SO2 + 2H2O
Silver nitrate solution added to a solution of a bromide that has been acidified (test with blue litmus paper) with dilute nitric acid gives a cream precipitate of silver bromide. The precipitate is readily soluble in concentrated ammonia:
Ag+(aq) + Br-(aq) " AgBr(s)
AgBr(s) + 2NH3(aq) " [Ag(NH3)2]+(aq) + Br-(aq)
Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates.
Oxidising agents oxidise bromide to bromine, which is yellow or orange in aqueous solution. Bromine can be extracted from the solution by shaking with an immiscible organic solvent, for example hexane, the organic layer then turning orange.
A suitable oxidising agent is sodium chlorate(I); this is added to the test solution, followed by a little dilute hydrochloric acid and a few cm3 of hexane:
OCl-(aq) + 2H+(aq) + 2Br-(aq) " Br2(aq) + Cl-(aq) + H2O(l)
Iodide I-
AgI, PbI2, Hg2I2 and CuI are insoluble in water.
Concentrated sulphuric acid gives a mixture of hydrogen iodide, iodine, hydrogen sulphide, sulphur and sulphur dioxide when added to solid iodides; the HI produced is oxidised by sulphuric acid. The mixture evolves purple acidic fumes, turns to a brown slurry, and is a mess:
NaI(s) + H2SO4(l) " NaHSO4(s) + HI(g)
2HI + H2SO4 " I2 + SO2 + 2H2O
6HI + H2SO4 " 3I2 + S + 4H2O
8HI + H2SO4 " 4I2 + H2S + 4H2O
There are no state symbols in these equations because the mixture is such a mess, and is more sulphuric acid than water.
Silver nitrate solution added to a solution of an iodide that has been acidified (test with blue litmus paper) with dilute nitric acid gives a yellow precipitate of silver iodide. The precipitate is insoluble even in concentrated ammonia:
Ag+(aq) + I-(aq) " AgI(s)
Acidification with nitric acid is necessary to eliminate carbonate or sulphite, both of which interfere with the test by giving spurious precipitates.
Oxidising agents oxidise iodide to iodine, which is yellow or orange in aqueous solution. Iodine can be extracted from the solution by shaking with an immiscible organic solvent, for example hexane, the organic layer then turning purple.
A suitable oxidising agent is sodium chlorate(I); this is added to the test solution, followed by a little dilute hydrochloric acid and a few cm3 of hexane:
OCl-(aq) + 2H+(aq) + 2Ir-(aq) " I2(aq) + Cl-(aq) + H2O(l)
Lead ethanoate or lead nitrate solutions give a bright yellow precipitate of lead(II) iodide with iodides:
Pb2+(aq) + 2I-(aq) " PbI2(s)
The colour of lead(II) iodide comes from interactions in the lattice; dissolving the salt in boiling water gives a colourless solution which deposits glittering yellow plates on cooling.
Solutions of copper(II) salts give a brown mixture containing iodine and copper(I) iodide when added to solutions of iodides. Addition of sodium thiosulphate solution decolourises the iodine and leaves pinkish-cream copper(I) iodide as a precipitate.
2Cu2+(aq) + 4I-(aq) " 2CuI(s) + I2(aq)
2S2O32-(aq) + I2(aq) " 2I-(aq) + S4O62-(aq)
This reaction is the basis for the volumetric estimation of copper; the iodine liberated from a known amount of a copper(II) solution is titrated with standard sodium thiosulphate solution.
Sulphite SO32-
Sulphurous acid is considerably stronger than carbonic acid, so sulphites do not give the effervescence that is characteristic of carbonates when dilute acid is added.
Dilute hydrochloric acid on warming with a sulphite evolves sulphur dioxide; this turns acidified potassium dichromate(VI) solution (or paper) green:
SO32-(aq) + 2H+(aq) " H2O(l) + SO2(g)
Barium chloride solution gives a white precipitate of barium sulphite; addition of dilute hydrochloric acid causes the precipitate to dissolve without effervescence:
SO32-(aq) + Ba2+(aq) " BaSO3(s)
Sulphate SO42-
BaSO4, SrSO4 and PbSO4 are insoluble; CaSO4 is sparingly soluble.
Barium chloride solution added to the test solution acidified with dilute hydrochloric acid gives a white precipitate of barium sulphate:
Ba2+(aq) + SO42-(aq) " BaSO4(s)
HSO4- does the same thing with barium ions; however the original test solution would then be very acidic, so that should be tested for.
The addition of HCl destroys any carbonate or sulphite ions present so prevents the spurious positive result due to the precipitation of these barium salts. Barium nitrate solution can be used instead of barium chloride.
18.2 Lead ethanoate solution gives a precipitate of white lead sulphate:
Pb2+(aq) + SO42-(aq) " PbSO4(s)
Nitrate NO3-
Since all nitrates are water soluble, there is no precipitation reaction for this ion.
Solid nitrates decompose on heating; those of group 1 (except Li) give the nitrite and oxygen;
2NaNO3(s) " 2NaNO2(s) + O2(g)
All others give the metal oxide, nitrogen dioxide, and oxygen. A brown gas is emitted that re-lights a glowing splint:
2Pb(NO3)2(s) " 2PbO(s) + O2(g) + 2NO2(g)
Nitrate ions are reduced to ammonia by boiling with aluminium or with Devarda’s Alloy in sodium hydroxide solution. Devarda’s Alloy contains aluminium, zinc and copper. Since ammonium ions also give ammonia with NaOH, the test solution must be boiled with NaOH and the vapour tested for ammonia; if present heating must continue until all the ammonia has gone. The mixture is then cooled, Devarda’s Alloy (or a piece of aluminium foil) added, and the mixture re-heated. A gas that turns moist red litmus paper blue indicates nitrate in the original solution:
3NO3-(aq) + 8Al(s) + 18H2O(l) + 21 OH-(aq) " 8[Al(OH)6]3-(aq) + 3NH3(g)
Not an equation to be remembered!
Carbonate CO32-
Only the alkali metal and ammonium carbonates are water soluble. Some carbonates (e.g. zinc, copper(II)) are basic carbonates and contain a proportion of the hydroxide in their structure.
Heating decomposes all but the alkali and alkaline earth metal carbonates (at Bunsen temperatures) giving the oxide and carbon dioxide:
CuCO3(s) " CuO(s) + CO2(g)
Dilute hydrochloric acid gives vigorous effervescence with carbonates, evolving carbon dioxide:
CO32-(aq or s) + 2H+(aq) " H2O(l) + CO2(g)
Bicarbonates also give this effervescence. The reaction of carbonates with acid is exothermic; bicarbonates react endothermically.
Hydrogen carbonate (bicarbonate) HCO3-
Only the alkali metal and ammonium bicarbonates are obtainable as solids; group 2 bicarbonates exist only in solution.
Calcium chloride solution on addition to a bicarbonate solution gives no precipitate since calcium bicarbonate is soluble; this distinguishes it from carbonate, which does give a precipitate. On heating the calcium chloride/bicarbonate mixture a white precipitate appears since the bicarbonate decomposes to carbonate:
Ca2+(aq) + 2HCO3-(aq) " CaCO3(s) + CO2(g) + H2O(l)
Ethanoate CH3COO-
Ethanoates on heating with dilute hydrochloric acid give ethanoic acid, recognisable by its vinegary smell.
Neutral iron(III) chloride solution added to neutral solutions of ethanoate ion give a deep red colouration owing to formation of iron(III) ethanoate.
Ethanedioate C2O42-
Concentrated sulphuric acid added to a solid ethanedioate salt gives a mixture of carbon monoxide and carbon dioxide from dehydration of the ethanedioic acid formed:
HOOC-COOH " H2O + CO2 + CO
Potassium manganate(VII) solution acidified with dilute sulphuric acid added to a solution of an ethanedioate causes the purple colour to disappear:
2MnO4-(aq) + 5C2O42-(aq) + 16H+(aq) " 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
Calcium chloride solution added to a solution of an ethanedioate gives a white precipitate of calcium ethanedioate:
Ca2+(aq) + C2O42-(aq) " CaC2O4(s)
The precipitate dissolves readily in dilute hydrochloric acid.
Chromate(VI) CrO42- and dichromate(VI) Cr2O72-
These ions are related through the equilibrium
Cr2O72-(aq) + 2OH-(aq) D 2CrO42-(aq) + H2O(l)
In alkaline solution the yellow chromate(VI) dominates, in acidic solution orange dichromate(VI). All dichromate(VI) salts are soluble; addition of dichromate(VI) ions to solutions of ions of metals which have insoluble chromate(VI) salts leads to the precipitation of chromates. This means that the only dichromates that can exist are those of group 1 metals, ammonium, magnesium, calcium and strontium.
Barium chloride solution added to a chromate(VI) or dichromate(VI) solution precipitates bright yellow barium chromate(VI):
Ba2+(aq) + CrO42-(aq) " BaCrO4(s)
The addition of a heavy metal ion to potassium dichromate solution precipitates the chromate and therefore moves the equilibrium to the right hand side. If the chromate is sparingly soluble (e.g. strontium) the supernatant liquid will remain yellow. Very insoluble chromates, such as lead, remove all the colour from the supernatant liquid.
Dichromate(VI) ion solution in sulphuric acid is an oxidising agent; oxidation is shown by the solution turning from orange to green (Cr(III)). The following can be oxidised:
(a) iron(II) to iron(III):
Cr2O72-(aq) + 14H+(aq) + 6Fe2+(aq) " 2Cr3+(aq) + 7H2O(l) + 6Fe3+(aq)
(b) iodide to iodine (the solution turns murky greenish-brown):
Cr2O72-(aq) + 14H+(aq) + 6I-(aq) " 2Cr3+(aq) + 7H2O(aq) + 3I2(aq)
(c) sulphite to sulphate:
Cr2O72-(aq) + 8H+(aq) + 3SO32-(aq) " 2Cr3+(aq) + 4H2O(l) + 3SO42-(aq)
(d) nitrite to nitrate:
Cr2O72-(aq) + 8H+(aq) + 3NO2-(aq) " 2Cr3+(aq) + 4H2O(l) + 3NO3-(aq)
(e) hydrogen peroxide reacts with acidified dichromate(VI) solutions to give a blue compound that rapidly turns green and evolves oxygen. The blue compound can be extracted into an organic solvent such as butan-1-ol. The blue compound is CrO5, which contains a peroxy structure. It is covalent, and is stable in organic solvents though not in water.
Cr2O72-(aq) + 8H+(aq) + 3H2O2 (aq) " 2Cr3+(aq) + 7H2O(l) + 3 O2(g)
Alcohols are oxidised by acidified potassium dichromate(VI) solution. Primary alcohols give aldehydes and then acids, secondary alcohols give ketones. Ethanol can be used to test for dichromate(VI) ions, therefore, the solution turning green and the apple smell of ethanal being evident.
Cation Analysis
Sodium Na+
Flame test: covered earlier. There are no simple reagents that will precipitate sodium compounds.
Potassium K+
Flame test: covered earlier. There are no simple reagents that will precipitate potassium compounds.
Magnesium Mg2+
Ions
Sodium hydroxide solution precipitates white magnesium hydroxide, insoluble in excess NaOH:
Mg2+(aq) + 2OH-(aq) " Mg(OH)2(s)
Ammonia solution only partially precipitates magnesium hydroxide, and not at all in the presence of ammonium ions. This is because magnesium hydroxide is fairly soluble, and the small concentration of OH- ions in ammonia becomes even smaller in the presence of ammonium ions and is not sufficient to produce the precipitate.
Sodium or potassium carbonate solution gives a white gelatinous precipitate of the basic carbonate Mg(OH)2.4MgCO3.5H2O.
Calcium Ca2+
Dilute sulphuric acid gives a white precipitate of calcium sulphate if the original solution is fairly concentrated:
Ca2+(aq) + SO42-(aq) " CaSO4(s)
Calcium sulphate solution is permanently hard water; its solubility is about 0.05 mol dm-3 at 25oC. Sodium or potassium carbonate solution precipitates white calcium carbonate:
Ca2+(aq) + CO32-(aq) " CaCO3(s)
Ammonium ethanedioate solution precipitates white calcium ethanedioate from neutral or alkaline solutions. The precipitate dissolves readily in dilute acid:
Ca2+(aq) + C2O42-(aq) " CaC2O4(s)
Calcium ethanedioate (calcium oxalate) is found in rhubarb leaves - it is what makes them poisonous.
Chromium(III) [Cr(H2O)6]3+
Sodium hydroxide solution precipitates grey-green chromium(III) hydroxide, which reacts with excess NaOH to give a deep green solution of the chromite ion:
[Cr(H2O)6]3+(aq) + 3OH-(aq) " Cr(OH)3(s) + 6H2O(l)
Cr(OH)3(s) + 3OH-(aq) " [Cr(OH)6]3-(aq)
Ammonia solution precipitates grey-green chromium(III) hydroxide, which with excess ammonia very slowly (see below) forms pinkish solutions of ammines. Ammines are transition metal complexes with ammonia, NH3. Do not confuse them with amines RNH2.The initial equation is as the first one above; then
Cr(OH)3(s) + 4NH3(aq) + 2H2O(l) " [Cr(NH3)4(H2O)2]3+(aq) + 3OH-(aq)
The exact composition of the ammine will depend on the amount of ammonia used. However, the reaction is so slow that several days are needed to see anything at all.
Oxidising agents convert chromium(III) to yellow chromate(VI) in the presence of alkali. Hydrogen peroxide is favoured:
2[Cr(H2O)6]3+(aq) + 3H2O2(aq) + 10OH-(aq) " 2CrO42-(aq) + 14H2O(l)
The test solution is made alkaline with NaOH solution, then an equal volume of ‘20 volume’ hydrogen peroxide solution is added and the mixture boiled. 20-volume hydrogen peroxide solution gives 20cm3 of oxygen per cm3 of solution that is decomposed to oxygen and water.
Manganese(II) [Mn(H2O)6]2+
Sodium hydroxide solution precipitates beige manganese(II) hydroxide. This is insoluble in excess reagent, and rapidly darkens owing to oxidation to hydrated manganese(IV) oxide:
[Mn(H2O)6]2+(aq) + 2OH-(aq) " Mn(OH)2(s) + 6H2O(l)
4Mn(OH)2(s) + 2H2O(l) + O2(aq) " 4MnO2.H2O(s)
Sodium or potassium carbonate solution gives a white precipitate of manganese(II) carbonate:
[Mn(H2O)6]2+(aq) + CO32-(aq) " MnCO3(s) + 6H2O(l)
Lead(IV) oxide in the presence of nitric acid converts manganese(II) into purple manganate(VII); an alternative oxidising agent is sodium bismuthate(V), NaBiO3:
5PbO2(s) + 2Mn2+(aq) + 6H+(aq) " 2MnO4-(aq) + 5Pb2+(aq) + 2H2O(l)
5BiO3-(aq) + 14H+(aq) + 2Mn2+(aq) " 5Bi3+(aq) + 2MnO4-(aq) + 7H2O(l)
A small amount of lead(IV) oxide or of sodium bismuthate(V) is added to the test solution, 6 or so drops of concentrated nitric acid added, and the mixture boiled. Filtration will give a purple filtrate if Mn2+ was present.
Iron(III) [Fe(H2O)6]3+
Aqueous iron(III) ions are not [Fe(H2O)6]3+; this ion is an amethyst (pale purple) colour, and is found only in solids such as aluminium iron(III) sulphate. In solution the hexaaqua ion is readily deprotonated by solvent water; the solution is acidic, and the yellow ion is [Fe(H2O)5OH]2+:
[Fe(H2O)6]3+ + H2O " [Fe(H2O)5OH]2+ + H3O+.
Sodium hydroxide solution precipitates foxy-red iron(III) hydroxide, insoluble in excess alkali:
[Fe(H2O)6]3+(aq) + 3OH-(aq) " Fe(OH)3(s) + 6H2O(l)
Ammonia solution reacts in the same way as sodium hydroxide – iron does not form complexes with ammonia.
Potassium hexacyanoferrate(II) precipitates dark blue potassium iron(III) hexacyanoferrate(II), Prussian Blue:
K+(aq) + Fe3+(aq) + [Fe(CN)6]4-(aq) " KFeIII[FeII(CN)6](s)
Prussian Blue is the pigment that is used to print 'blueprints'. It was also used for Prussian army uniforms, and for the locomotives and rolling stock of the Somerset & Dorset Joint Railway, pre-1923.
Iron(II) [Fe(H2O)6]2+
Sodium hydroxide solution precipitates dirty green iron(II) hydroxide; on standing the surface of the precipitate turns foxy-red owing to air oxidation to iron(III) hydroxide:
[Fe(H2O)6]2+(aq) + 2OH-(aq) " Fe(OH)2(s) + 6H2O(l)
Ammonia solution behaves similarly to sodium hydroxide – again there are no ammine complexes.
Potassium hexacyanoferrate(III) precipitates dark blue potassium iron(II) hexacyanoferrate(III), Turnbull’s Blue:
K+(aq) + Fe2+(aq) + [Fe(CN)6]3-(aq) " KFeII[FeIII(CN)6](s)
Although Turnbull's Blue and Prussian Blue don't have quite the same colour, they are in fact the same substance.
Potassium manganate(VII) oxidises iron(II) to iron(III) in acidic solution. The purple manganate colour is lost and the resulting solution of iron(III) is yellow:
MnO4-(aq) + 5Fe2+(aq) + 8H+(aq) " Mn2+(aq) + 5Fe3+(aq) + 4H2O(aq)
Potassium dichromate(VI) oxidises iron(II) to iron(III) in acidic solution. The orange dichromate colour is lost and the resulting solution is green owing to presence of chromium(III) which masks the iron(III) colour:
Cr2O72-(aq) + 6Fe2+(aq) + 14H+(aq) " 2Cr3+(aq) + 6Fe3+(aq) + 7H2O(l)
Copper(II) [Cu(H2O)6]2+
Sodium hydroxide solution gives a pale blue precipitate usually described as copper(II) hydroxide:
[Cu(H2O)6]2+(aq) + 2OH-(aq) " Cu(OH)2(s) + 6H2O(l)
In fact this precipitate is the basic salt – thus from copper(II) sulphate solution the product is basic copper sulphate:
2[Cu(H2O)6]2+(aq) + 2OH-(aq) + SO42-(aq) " Cu(OH)2.CuSO4(s) + 12H2O(l)
Basic copper sulphate or basic copper carbonate is responsible for the green patina - verdigris - found on weathered copper roofs. The sulphate is present in industrial surroundings. To get true copper(II) hydroxide, the solution of the copper(II) salt must be added to excess sodium hydroxide; the precipitate is then blue rather than turquoise. In both cases CAUTIOUS warming of the mixture causes the precipitate to lose water and turn black as CuO(s) is formed. The loss of water by heating in an aqueous medium is unusual.
Ammonia solution initially gives a blue precipitate as for sodium hydroxide. Further addition of ammonia gives deep blue soluble cuprammine complexes whose composition depends on the amount of ammonia present. The reaction is usually represented:
Cu(OH)2(s) + 4NH3(aq) + 2H2O(l) " [Cu(NH3)4(H2O)2]2+(aq) +2OH-(aq)
The cuprammine solution has the bizarre property of being able to dissolve cellulose - bits of filter paper, for e example. The resulting viscous substance can be extruded into a bath of dilute sulphuric acid to give viscose fibre, which is a form of artificial silk.
Concentrated HCl or a saturated solution of sodium chloride gives a bright green solution containing CuCl42- ions; dilution with water causes the solution to turn pale blue as the [Cu(H2O)6]2+ ion re-forms.
[Cu(H2O)6]2+(aq) + 4Cl-(aq) " CuCl42-(aq) + 6H2O(l)
Sodium carbonate solution precipitates greenish-blue basic carbonates of indefinite composition.
Potassium iodide solution precipitates iodine and copper(I) iodide; this reaction is used volumetrically for the estimation of copper, the liberated iodine being titrated with sodium thiosulphate solution:
2[Cu(H2O)6]2+(aq) + 4I-(aq) " 2CuI(s) + I2(aq) + 6H2O(l)
The mixture turns a sludgy brown colour; addition of sodium thiosulphate solution leaves a creamy-pink precipitate of copper(I) iodide.
Zinc(II) [Zn(H2O)6]2+
Sodium hydroxide solution precipitates white zinc(II) hydroxide, easily soluble in excess sodium hydroxide to give sodium zincate(II), because the hydroxide is amphoteric:
[Zn(H2O)6]2+(aq) + 2OH-(aq) " Zn(OH)2(s) + 6H2O(l)
Zn(OH)2(s) + 2OH-(aq) " Zn(OH)42-(aq)
Ammonia solution initially precipitates zinc(II) hydroxide, as with sodium hydroxide. Excess ammonia causes the precipitate to disappear owing to the formation of ammine complexes – a different reason from that in 10.1:
Zn(OH)2(s) + 4NH3(aq) " [Zn(NH3)4]2+(aq) + 2OH-(aq)
Sodium carbonate solution precipitates a white basic carbonate:
5Zn2+(aq) + 2CO32-(aq) + 6OH-(aq) " 2ZnCO3.3Zn(OH)2(s)
This latter substance is what you get when you buy zinc carbonate – which is why on heating it gives off water vapour as well as carbon dioxide.
Aluminium Al[H2O] 3+
Sodium hydroxide solution precipitates white gelatinous aluminium hydroxide. This reacts with excess NaOH to give a colourless solution of sodium aluminate:
[Al(H2O)6] 3+(aq) + 3OH-(aq) " Al(OH)3(s) + 6H2O(l)
Al(OH)3(s) + 3OH-(aq) " [Al(OH)6]3-(aq)
Ammonia solution precipitates white aluminium hydroxide, as with NaOH; however it does not react further with excess ammonia.
Lead(II) Pb2+
Dilute hydrochloric acid or other soluble chlorides precipitate white lead(II) chloride from moderately concentrated solutions of lead(II) salts.
Pb2+(aq) + 2Cl-(aq) " PbCl2(s)
The precipitate dissolves in an excess of concentrated hydrochloric acid to give yellow solutions containing various species including PbCl3- and PbCl42-.
Sodium hydroxide solution gives a white precipitate of lead(II) hydroxide which is soluble in excess NaOH to give colourless sodium plumbate(II)
Pb2+(aq) + 2OH-(aq) " Pb(OH)2(s)
Pb(OH)2(s) + 2OH-(aq) " [Pb(OH)4]2-(aq)
Ammonia solution precipitates white lead(II) hydroxide; this does not dissolve in excess ammonia, because no complexes are formed and ammonia is not a strong enough base to bring out the amphoteric properties of lead(II) hydroxide.
Dilute sulphuric acid or other soluble sulphates give a white precipitate of lead(II) sulphate:
Pb2+(aq) + SO42-(aq) " PbSO4(s)
Potassium chromate(VI) solution gives a bright yellow precipitate of lead(II) chromate(VI):
Pb2+(aq) + CrO42-(aq) " PbCrO4(s)
Potassium iodide solution gives a bright yellow precipitate of lead(II) iodide. This will dissolve in boiling water to give a colourless solution – the colour comes from interactions in the crystal lattice rather than from coloured ions:
Pb2+(aq) + 2I-(aq) " PbI2(s)
Ammonium NH4+
Heat: all ammonium salts decompose on heating; ammonium nitrate may explode. Ammonium chloride and sulphate give products that recombine on cooling, so that the salts apparently sublime.
NH4Cl(s) " NH3(g) + HCl(g)
NH4NO3(s) " N2O(g) + 2H2O(g)
2(NH4)2SO4(s) D 2NH3(g) + SO3(g) + H2O(g)
Ammonium dichromate(VI) decomposes spectacularly on ignition in a reaction that is oxidation of the cation by the anion; the initially orange solid gives a fluffy green product of much larger volume:
(NH4)2Cr2O7(s) " N2(g) + Cr2O3(s) + 4H2O(g)
Alkalis (sodium hydroxide, calcium hydroxide) liberate ammonia from ammonium salts on warming with the solution, or even from a mixture of the solids. This is because OH- is a stronger base than ammonia, so removes a hydrogen ion from the ammonium ion:
NH4+(aq) + OH-(aq) " NH3(g) + H2O(l)
The test solution is warmed with sodium hydroxide solution and the vapours tested with moist red litmus paper. It is important to test the vapours immediately heating begins, since the ammonia is lost very quickly and by the time the solution boils it may well have all gone.
Cobalt(II) [Co(H2O)6]2+
The (+2) state is the most stable for simple cobalt salts; they are coloured pink or blue. The pink colour of the hydrated ion [Co(H2O)6]2+ changes to blue on heating, on dehydration, or in the presence of concentrated acids.
Sodium hydroxide solution precipitates a blue basic salt which on warming with excess sodium hydroxide forms solid pink cobalt(II) hydroxide. With a solution of cobalt(II) nitrate the equations are:
[Co(H2O)6]2+ (aq) + NO3 – (aq) + OH – (aq) "
Co(OH)NO3 (s) + 6H2O(l)
blue
Co(OH)NO3 (s) + OH – (aq) "
Co(OH)2 (s) + NO3 – (aq)
pink
Basic salts, usually hydroxides or carbonates, contain not only the expected hydroxide or carbonate ion but also the anion from the original salt. The precipitate from cobalt(II) nitrate solution is therefore different from that which would be given from cobalt(II) chloride. The latter would be Co(OH)Cl.
If the pink precipitate is boiled for some time (or is warmed with hydrogen peroxide solution) it is converted to brownish-black cobalt(III) hydroxide:
4Co(OH)2 (s) + 2H2O (l) + O2 (aq) " 4Co(OH)3 (s)
Ammonia solution gives a blue basic precipitate as with sodium hydroxide; excess ammonia converts this to a brown solution of a cobaltammine which turns red on exposure to air giving an ammine of cobalt(III). The overall reaction is:
Co(OH)NO3 (s) + 28NH3(aq) + 6H2O(l) + O2(aq) " 4[Co(NH3)6](OH)3(aq) + 4NH4NO3(aq)
Nickel(II) [Ni(H2O)6]2+
Sodium hydroxide solution gives a green gelatinous precipitate of nickel(II) hydroxide, insoluble in excess of the reagent:
[Ni(H2O)6]2+ (aq) + 2OH – (aq) " Ni(OH)2 (s) + 6H2O(l)
Ammonia solution gives a green precipitate of a basic salt which dissolves readily in excess ammonia solution to form a blue-violet solution of complex nickel ammines. With nickel(II) chloride the reactions are:
[Ni(H2O)6]2+(aq) + Cl – (aq) + NH3(aq) " Ni(OH)Cl(s) + NH4+(aq) + 5H2O(l)
Ni(OH)Cl(s) + 5NH3(aq) + H2O(l) " [Ni(NH3)4](OH)2(aq) + NH4+(aq) + Cl – (aq)
Ni(OH)Cl(s) + 7NH3(aq) + H2O(l) " [Ni(NH3)6](OH)2(aq) + NH4+(aq) + Cl – (aq)
Flame test: covered earlier. There are no simple reagents that will precipitate sodium compounds.
Potassium K+
Flame test: covered earlier. There are no simple reagents that will precipitate potassium compounds.
Magnesium Mg2+
Ions
Sodium hydroxide solution precipitates white magnesium hydroxide, insoluble in excess NaOH:
Mg2+(aq) + 2OH-(aq) " Mg(OH)2(s)
Ammonia solution only partially precipitates magnesium hydroxide, and not at all in the presence of ammonium ions. This is because magnesium hydroxide is fairly soluble, and the small concentration of OH- ions in ammonia becomes even smaller in the presence of ammonium ions and is not sufficient to produce the precipitate.
Sodium or potassium carbonate solution gives a white gelatinous precipitate of the basic carbonate Mg(OH)2.4MgCO3.5H2O.
Calcium Ca2+
Dilute sulphuric acid gives a white precipitate of calcium sulphate if the original solution is fairly concentrated:
Ca2+(aq) + SO42-(aq) " CaSO4(s)
Calcium sulphate solution is permanently hard water; its solubility is about 0.05 mol dm-3 at 25oC. Sodium or potassium carbonate solution precipitates white calcium carbonate:
Ca2+(aq) + CO32-(aq) " CaCO3(s)
Ammonium ethanedioate solution precipitates white calcium ethanedioate from neutral or alkaline solutions. The precipitate dissolves readily in dilute acid:
Ca2+(aq) + C2O42-(aq) " CaC2O4(s)
Calcium ethanedioate (calcium oxalate) is found in rhubarb leaves - it is what makes them poisonous.
Chromium(III) [Cr(H2O)6]3+
Sodium hydroxide solution precipitates grey-green chromium(III) hydroxide, which reacts with excess NaOH to give a deep green solution of the chromite ion:
[Cr(H2O)6]3+(aq) + 3OH-(aq) " Cr(OH)3(s) + 6H2O(l)
Cr(OH)3(s) + 3OH-(aq) " [Cr(OH)6]3-(aq)
Ammonia solution precipitates grey-green chromium(III) hydroxide, which with excess ammonia very slowly (see below) forms pinkish solutions of ammines. Ammines are transition metal complexes with ammonia, NH3. Do not confuse them with amines RNH2.The initial equation is as the first one above; then
Cr(OH)3(s) + 4NH3(aq) + 2H2O(l) " [Cr(NH3)4(H2O)2]3+(aq) + 3OH-(aq)
The exact composition of the ammine will depend on the amount of ammonia used. However, the reaction is so slow that several days are needed to see anything at all.
Oxidising agents convert chromium(III) to yellow chromate(VI) in the presence of alkali. Hydrogen peroxide is favoured:
2[Cr(H2O)6]3+(aq) + 3H2O2(aq) + 10OH-(aq) " 2CrO42-(aq) + 14H2O(l)
The test solution is made alkaline with NaOH solution, then an equal volume of ‘20 volume’ hydrogen peroxide solution is added and the mixture boiled. 20-volume hydrogen peroxide solution gives 20cm3 of oxygen per cm3 of solution that is decomposed to oxygen and water.
Manganese(II) [Mn(H2O)6]2+
Sodium hydroxide solution precipitates beige manganese(II) hydroxide. This is insoluble in excess reagent, and rapidly darkens owing to oxidation to hydrated manganese(IV) oxide:
[Mn(H2O)6]2+(aq) + 2OH-(aq) " Mn(OH)2(s) + 6H2O(l)
4Mn(OH)2(s) + 2H2O(l) + O2(aq) " 4MnO2.H2O(s)
Sodium or potassium carbonate solution gives a white precipitate of manganese(II) carbonate:
[Mn(H2O)6]2+(aq) + CO32-(aq) " MnCO3(s) + 6H2O(l)
Lead(IV) oxide in the presence of nitric acid converts manganese(II) into purple manganate(VII); an alternative oxidising agent is sodium bismuthate(V), NaBiO3:
5PbO2(s) + 2Mn2+(aq) + 6H+(aq) " 2MnO4-(aq) + 5Pb2+(aq) + 2H2O(l)
5BiO3-(aq) + 14H+(aq) + 2Mn2+(aq) " 5Bi3+(aq) + 2MnO4-(aq) + 7H2O(l)
A small amount of lead(IV) oxide or of sodium bismuthate(V) is added to the test solution, 6 or so drops of concentrated nitric acid added, and the mixture boiled. Filtration will give a purple filtrate if Mn2+ was present.
Iron(III) [Fe(H2O)6]3+
Aqueous iron(III) ions are not [Fe(H2O)6]3+; this ion is an amethyst (pale purple) colour, and is found only in solids such as aluminium iron(III) sulphate. In solution the hexaaqua ion is readily deprotonated by solvent water; the solution is acidic, and the yellow ion is [Fe(H2O)5OH]2+:
[Fe(H2O)6]3+ + H2O " [Fe(H2O)5OH]2+ + H3O+.
Sodium hydroxide solution precipitates foxy-red iron(III) hydroxide, insoluble in excess alkali:
[Fe(H2O)6]3+(aq) + 3OH-(aq) " Fe(OH)3(s) + 6H2O(l)
Ammonia solution reacts in the same way as sodium hydroxide – iron does not form complexes with ammonia.
Potassium hexacyanoferrate(II) precipitates dark blue potassium iron(III) hexacyanoferrate(II), Prussian Blue:
K+(aq) + Fe3+(aq) + [Fe(CN)6]4-(aq) " KFeIII[FeII(CN)6](s)
Prussian Blue is the pigment that is used to print 'blueprints'. It was also used for Prussian army uniforms, and for the locomotives and rolling stock of the Somerset & Dorset Joint Railway, pre-1923.
Iron(II) [Fe(H2O)6]2+
Sodium hydroxide solution precipitates dirty green iron(II) hydroxide; on standing the surface of the precipitate turns foxy-red owing to air oxidation to iron(III) hydroxide:
[Fe(H2O)6]2+(aq) + 2OH-(aq) " Fe(OH)2(s) + 6H2O(l)
Ammonia solution behaves similarly to sodium hydroxide – again there are no ammine complexes.
Potassium hexacyanoferrate(III) precipitates dark blue potassium iron(II) hexacyanoferrate(III), Turnbull’s Blue:
K+(aq) + Fe2+(aq) + [Fe(CN)6]3-(aq) " KFeII[FeIII(CN)6](s)
Although Turnbull's Blue and Prussian Blue don't have quite the same colour, they are in fact the same substance.
Potassium manganate(VII) oxidises iron(II) to iron(III) in acidic solution. The purple manganate colour is lost and the resulting solution of iron(III) is yellow:
MnO4-(aq) + 5Fe2+(aq) + 8H+(aq) " Mn2+(aq) + 5Fe3+(aq) + 4H2O(aq)
Potassium dichromate(VI) oxidises iron(II) to iron(III) in acidic solution. The orange dichromate colour is lost and the resulting solution is green owing to presence of chromium(III) which masks the iron(III) colour:
Cr2O72-(aq) + 6Fe2+(aq) + 14H+(aq) " 2Cr3+(aq) + 6Fe3+(aq) + 7H2O(l)
Copper(II) [Cu(H2O)6]2+
Sodium hydroxide solution gives a pale blue precipitate usually described as copper(II) hydroxide:
[Cu(H2O)6]2+(aq) + 2OH-(aq) " Cu(OH)2(s) + 6H2O(l)
In fact this precipitate is the basic salt – thus from copper(II) sulphate solution the product is basic copper sulphate:
2[Cu(H2O)6]2+(aq) + 2OH-(aq) + SO42-(aq) " Cu(OH)2.CuSO4(s) + 12H2O(l)
Basic copper sulphate or basic copper carbonate is responsible for the green patina - verdigris - found on weathered copper roofs. The sulphate is present in industrial surroundings. To get true copper(II) hydroxide, the solution of the copper(II) salt must be added to excess sodium hydroxide; the precipitate is then blue rather than turquoise. In both cases CAUTIOUS warming of the mixture causes the precipitate to lose water and turn black as CuO(s) is formed. The loss of water by heating in an aqueous medium is unusual.
Ammonia solution initially gives a blue precipitate as for sodium hydroxide. Further addition of ammonia gives deep blue soluble cuprammine complexes whose composition depends on the amount of ammonia present. The reaction is usually represented:
Cu(OH)2(s) + 4NH3(aq) + 2H2O(l) " [Cu(NH3)4(H2O)2]2+(aq) +2OH-(aq)
The cuprammine solution has the bizarre property of being able to dissolve cellulose - bits of filter paper, for e example. The resulting viscous substance can be extruded into a bath of dilute sulphuric acid to give viscose fibre, which is a form of artificial silk.
Concentrated HCl or a saturated solution of sodium chloride gives a bright green solution containing CuCl42- ions; dilution with water causes the solution to turn pale blue as the [Cu(H2O)6]2+ ion re-forms.
[Cu(H2O)6]2+(aq) + 4Cl-(aq) " CuCl42-(aq) + 6H2O(l)
Sodium carbonate solution precipitates greenish-blue basic carbonates of indefinite composition.
Potassium iodide solution precipitates iodine and copper(I) iodide; this reaction is used volumetrically for the estimation of copper, the liberated iodine being titrated with sodium thiosulphate solution:
2[Cu(H2O)6]2+(aq) + 4I-(aq) " 2CuI(s) + I2(aq) + 6H2O(l)
The mixture turns a sludgy brown colour; addition of sodium thiosulphate solution leaves a creamy-pink precipitate of copper(I) iodide.
Zinc(II) [Zn(H2O)6]2+
Sodium hydroxide solution precipitates white zinc(II) hydroxide, easily soluble in excess sodium hydroxide to give sodium zincate(II), because the hydroxide is amphoteric:
[Zn(H2O)6]2+(aq) + 2OH-(aq) " Zn(OH)2(s) + 6H2O(l)
Zn(OH)2(s) + 2OH-(aq) " Zn(OH)42-(aq)
Ammonia solution initially precipitates zinc(II) hydroxide, as with sodium hydroxide. Excess ammonia causes the precipitate to disappear owing to the formation of ammine complexes – a different reason from that in 10.1:
Zn(OH)2(s) + 4NH3(aq) " [Zn(NH3)4]2+(aq) + 2OH-(aq)
Sodium carbonate solution precipitates a white basic carbonate:
5Zn2+(aq) + 2CO32-(aq) + 6OH-(aq) " 2ZnCO3.3Zn(OH)2(s)
This latter substance is what you get when you buy zinc carbonate – which is why on heating it gives off water vapour as well as carbon dioxide.
Aluminium Al[H2O] 3+
Sodium hydroxide solution precipitates white gelatinous aluminium hydroxide. This reacts with excess NaOH to give a colourless solution of sodium aluminate:
[Al(H2O)6] 3+(aq) + 3OH-(aq) " Al(OH)3(s) + 6H2O(l)
Al(OH)3(s) + 3OH-(aq) " [Al(OH)6]3-(aq)
Ammonia solution precipitates white aluminium hydroxide, as with NaOH; however it does not react further with excess ammonia.
Lead(II) Pb2+
Dilute hydrochloric acid or other soluble chlorides precipitate white lead(II) chloride from moderately concentrated solutions of lead(II) salts.
Pb2+(aq) + 2Cl-(aq) " PbCl2(s)
The precipitate dissolves in an excess of concentrated hydrochloric acid to give yellow solutions containing various species including PbCl3- and PbCl42-.
Sodium hydroxide solution gives a white precipitate of lead(II) hydroxide which is soluble in excess NaOH to give colourless sodium plumbate(II)
Pb2+(aq) + 2OH-(aq) " Pb(OH)2(s)
Pb(OH)2(s) + 2OH-(aq) " [Pb(OH)4]2-(aq)
Ammonia solution precipitates white lead(II) hydroxide; this does not dissolve in excess ammonia, because no complexes are formed and ammonia is not a strong enough base to bring out the amphoteric properties of lead(II) hydroxide.
Dilute sulphuric acid or other soluble sulphates give a white precipitate of lead(II) sulphate:
Pb2+(aq) + SO42-(aq) " PbSO4(s)
Potassium chromate(VI) solution gives a bright yellow precipitate of lead(II) chromate(VI):
Pb2+(aq) + CrO42-(aq) " PbCrO4(s)
Potassium iodide solution gives a bright yellow precipitate of lead(II) iodide. This will dissolve in boiling water to give a colourless solution – the colour comes from interactions in the crystal lattice rather than from coloured ions:
Pb2+(aq) + 2I-(aq) " PbI2(s)
Ammonium NH4+
Heat: all ammonium salts decompose on heating; ammonium nitrate may explode. Ammonium chloride and sulphate give products that recombine on cooling, so that the salts apparently sublime.
NH4Cl(s) " NH3(g) + HCl(g)
NH4NO3(s) " N2O(g) + 2H2O(g)
2(NH4)2SO4(s) D 2NH3(g) + SO3(g) + H2O(g)
Ammonium dichromate(VI) decomposes spectacularly on ignition in a reaction that is oxidation of the cation by the anion; the initially orange solid gives a fluffy green product of much larger volume:
(NH4)2Cr2O7(s) " N2(g) + Cr2O3(s) + 4H2O(g)
Alkalis (sodium hydroxide, calcium hydroxide) liberate ammonia from ammonium salts on warming with the solution, or even from a mixture of the solids. This is because OH- is a stronger base than ammonia, so removes a hydrogen ion from the ammonium ion:
NH4+(aq) + OH-(aq) " NH3(g) + H2O(l)
The test solution is warmed with sodium hydroxide solution and the vapours tested with moist red litmus paper. It is important to test the vapours immediately heating begins, since the ammonia is lost very quickly and by the time the solution boils it may well have all gone.
Cobalt(II) [Co(H2O)6]2+
The (+2) state is the most stable for simple cobalt salts; they are coloured pink or blue. The pink colour of the hydrated ion [Co(H2O)6]2+ changes to blue on heating, on dehydration, or in the presence of concentrated acids.
Sodium hydroxide solution precipitates a blue basic salt which on warming with excess sodium hydroxide forms solid pink cobalt(II) hydroxide. With a solution of cobalt(II) nitrate the equations are:
[Co(H2O)6]2+ (aq) + NO3 – (aq) + OH – (aq) "
Co(OH)NO3 (s) + 6H2O(l)
blue
Co(OH)NO3 (s) + OH – (aq) "
Co(OH)2 (s) + NO3 – (aq)
pink
Basic salts, usually hydroxides or carbonates, contain not only the expected hydroxide or carbonate ion but also the anion from the original salt. The precipitate from cobalt(II) nitrate solution is therefore different from that which would be given from cobalt(II) chloride. The latter would be Co(OH)Cl.
If the pink precipitate is boiled for some time (or is warmed with hydrogen peroxide solution) it is converted to brownish-black cobalt(III) hydroxide:
4Co(OH)2 (s) + 2H2O (l) + O2 (aq) " 4Co(OH)3 (s)
Ammonia solution gives a blue basic precipitate as with sodium hydroxide; excess ammonia converts this to a brown solution of a cobaltammine which turns red on exposure to air giving an ammine of cobalt(III). The overall reaction is:
Co(OH)NO3 (s) + 28NH3(aq) + 6H2O(l) + O2(aq) " 4[Co(NH3)6](OH)3(aq) + 4NH4NO3(aq)
Nickel(II) [Ni(H2O)6]2+
Sodium hydroxide solution gives a green gelatinous precipitate of nickel(II) hydroxide, insoluble in excess of the reagent:
[Ni(H2O)6]2+ (aq) + 2OH – (aq) " Ni(OH)2 (s) + 6H2O(l)
Ammonia solution gives a green precipitate of a basic salt which dissolves readily in excess ammonia solution to form a blue-violet solution of complex nickel ammines. With nickel(II) chloride the reactions are:
[Ni(H2O)6]2+(aq) + Cl – (aq) + NH3(aq) " Ni(OH)Cl(s) + NH4+(aq) + 5H2O(l)
Ni(OH)Cl(s) + 5NH3(aq) + H2O(l) " [Ni(NH3)4](OH)2(aq) + NH4+(aq) + Cl – (aq)
Ni(OH)Cl(s) + 7NH3(aq) + H2O(l) " [Ni(NH3)6](OH)2(aq) + NH4+(aq) + Cl – (aq)
Monday, December 15, 2008
Atomic structure and periodic table
1.5a Atomic structure and periodic table
1.5a Relative massesThe unit of mass is 1/12 of the mass of an atom of the isotope Carbon-12 (12C=12 exactly)
The Relative Atomic Mass (Ar)of an element is the ratio of the mass of an average atom of that element to 1/12 of the mass of an atom of the nuclide Carbon-12.
The Relative Isotopic Mass of a nuclide is the ratio of the mass of one atom of that nuclide to 1/12 of the mass of a Carbon-12 atom.
The Relative Molecular Mass (Mr)of a substance is the ratio of the mass of an average molecule of that substance to 1/12 of the mass of an atom of the nuclide Carbon-12.
Task1.5a Explain what is meant by
The relative atomic mass of helium is 4.002602.
The relative isopotic mass of deuterium is 2.0000.
The relative molecular mass of water is 18.
1.5b mass spectrometer
Uses of mass spectrometers
The mass spectrometer can detect the abundance of isotopes such as U-238 and Pb-206.
In any one place the Earth concentrates some minerals like uranium but not others like lead. So some rocks form containing uranium but not lead.
U-238 is radioactive and decays to form Pb-206 over time.
This process is controlled by the half life of U-238 which is 4.5*109 years. The relative abundance of U-238 to Pb-206 can be used to calculate the age of the rock.
Rock is found to have ages up to about 4 thousand million years, but not older. This is the age of the Earth when the first solid rock formed from the magma on the surface of the cooling planet.
Radiocarbon dating using C-14 can be used to similarly date organic matter.
The mass spectrum of an element shows the relative isotopic mass (m/e) and relative abundance of each isotope of the element being tested.
In a mass spectrum the height of each peak = the relative abundance.
The Relative Atomic Mass of an element can be found by finding the sum of the products of the relative abundance of each isotope and its relative isotopic mass and the dividing by the total relative abundance.
work sheets mass spetrometer ( AS level)
Write a sentence for each stage in the formation of a mass spectrum.
Isotopes are atoms of the same element with different numbers of neutrons in the nucleus
The Relative Atomic Mass of an element is the weighted (to take account of relative abundance) average of the the Relative Isotopic Masses of all of the isotopes of that element.E.g. Chlorine has two isotopes with mass numbers (and relative isotopic masses
35 and37
35 37
75% is Cl and 25% is Cl
17 17
Let there be 100 atomsTotal mass of 100 atoms = (75 * 35) + (25 * 37) = 3550
Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms=3550/100so relative atomic mass of chlorine = 35.5
PRACTICE QUASTIONS
1. Bromine exists as two isotopes with mass numbers 81 and 79. If there is 50% of each isotope in a sample what is the relative atomic mass of bromine?
2. Calculate the relative atomic mass of the following:(a) gallium 60% Ga 69 and 40% Ga 71
(b) neon 90% Ne 20 and 10% Ne 22
(c) silver 50% Ag 107 and 50% Ag 109
(d) boron 20% B 10 and 80% B 11
(e) neon 90.92% Ne 20, 8.82% Ne 22 and 0.26% Ne 21.
(f) lead 52.3% Pb 208, 22.6% Pb 207, 23.6% Pb 206 and 1.5% Pb 204.
ionization energies
The first ionisation energy is the enthalpy change when 1 mole of gaseous atoms of an element each lose an electron to form gaseous ions each with a single positive charge.
It is the enthalpy change for the reaction: M(g) ---> M+(g) + e-
The second ionisation energy is the enthalpy change when 1 mole of gaseous ions of an element each with a single positive charge each lose an electron to form gaseous ions each with a double positive charge.
It is the enthalpy change for the reaction: M+(g)---> M2+(g) + e-
All of these ionizations involve the removal of a negative electron from a positive nucleus. The attraction between these must be over come by supplying energy. The process is always endothermic. Example 1st IE of Na = +500kJmol-1
Task1.5d.1 Define the third ionisation enthalpy for the element M and give an equation.
Task1.5d.2 Define the first ionisation enthalpy for sodium, magnesium, chlorine and neon.
Task1.5d.3 Define the second ionisation enthalpy for lithium, aluminium and oxygen.
1.5e Electronic structureThe graph of successive ionisation energies against ionisation number shows electrons grouped into three energy levels or quantum shells with similar energies.
Electron 1 (ionisation number = 1, removed first) from the third shell, electrons 2-9 in second shell, electrons 10 and 11(removed last) from first shell closest to the nucleus.
Electron 1 is on its own in the outer shell indicating that this element is in group 1. The graph of first ionisation energy against atomic number shows the grouping of electrons into s, p (and d) subshells (orbitals) within the (quantum shells) energy levels. The general trend is for ionisation energy to increase with increasing atomic number across a period but 1st IE B < 1st IE Be and 1st IE O < 1st IE N.
PRACTICE QUSTIONS
Label the diagram with symbols for the elements.
Label electrons removed as s or p.
Write two sentences using the frame: The ionisation energy (increases/decreases) (from left to right across a period/down a group) because (distance from outer electrons to nucleus is increasing/the nuclear charge in increasing)
1.5a Relative massesThe unit of mass is 1/12 of the mass of an atom of the isotope Carbon-12 (12C=12 exactly)
The Relative Atomic Mass (Ar)of an element is the ratio of the mass of an average atom of that element to 1/12 of the mass of an atom of the nuclide Carbon-12.
The Relative Isotopic Mass of a nuclide is the ratio of the mass of one atom of that nuclide to 1/12 of the mass of a Carbon-12 atom.
The Relative Molecular Mass (Mr)of a substance is the ratio of the mass of an average molecule of that substance to 1/12 of the mass of an atom of the nuclide Carbon-12.
Task1.5a Explain what is meant by
The relative atomic mass of helium is 4.002602.
The relative isopotic mass of deuterium is 2.0000.
The relative molecular mass of water is 18.
1.5b mass spectrometer
Uses of mass spectrometers
The mass spectrometer can detect the abundance of isotopes such as U-238 and Pb-206.
In any one place the Earth concentrates some minerals like uranium but not others like lead. So some rocks form containing uranium but not lead.
U-238 is radioactive and decays to form Pb-206 over time.
This process is controlled by the half life of U-238 which is 4.5*109 years. The relative abundance of U-238 to Pb-206 can be used to calculate the age of the rock.
Rock is found to have ages up to about 4 thousand million years, but not older. This is the age of the Earth when the first solid rock formed from the magma on the surface of the cooling planet.
Radiocarbon dating using C-14 can be used to similarly date organic matter.
The mass spectrum of an element shows the relative isotopic mass (m/e) and relative abundance of each isotope of the element being tested.
In a mass spectrum the height of each peak = the relative abundance.
The Relative Atomic Mass of an element can be found by finding the sum of the products of the relative abundance of each isotope and its relative isotopic mass and the dividing by the total relative abundance.
work sheets mass spetrometer ( AS level)
Write a sentence for each stage in the formation of a mass spectrum.
Isotopes are atoms of the same element with different numbers of neutrons in the nucleus
The Relative Atomic Mass of an element is the weighted (to take account of relative abundance) average of the the Relative Isotopic Masses of all of the isotopes of that element.E.g. Chlorine has two isotopes with mass numbers (and relative isotopic masses
35 and37
35 37
75% is Cl and 25% is Cl
17 17
Let there be 100 atomsTotal mass of 100 atoms = (75 * 35) + (25 * 37) = 3550
Average mass of an atom (relative atomic mass of chlorine) = Total mass /Number of atoms=3550/100so relative atomic mass of chlorine = 35.5
PRACTICE QUASTIONS
1. Bromine exists as two isotopes with mass numbers 81 and 79. If there is 50% of each isotope in a sample what is the relative atomic mass of bromine?
2. Calculate the relative atomic mass of the following:(a) gallium 60% Ga 69 and 40% Ga 71
(b) neon 90% Ne 20 and 10% Ne 22
(c) silver 50% Ag 107 and 50% Ag 109
(d) boron 20% B 10 and 80% B 11
(e) neon 90.92% Ne 20, 8.82% Ne 22 and 0.26% Ne 21.
(f) lead 52.3% Pb 208, 22.6% Pb 207, 23.6% Pb 206 and 1.5% Pb 204.
ionization energies
The first ionisation energy is the enthalpy change when 1 mole of gaseous atoms of an element each lose an electron to form gaseous ions each with a single positive charge.
It is the enthalpy change for the reaction: M(g) ---> M+(g) + e-
The second ionisation energy is the enthalpy change when 1 mole of gaseous ions of an element each with a single positive charge each lose an electron to form gaseous ions each with a double positive charge.
It is the enthalpy change for the reaction: M+(g)---> M2+(g) + e-
All of these ionizations involve the removal of a negative electron from a positive nucleus. The attraction between these must be over come by supplying energy. The process is always endothermic. Example 1st IE of Na = +500kJmol-1
Task1.5d.1 Define the third ionisation enthalpy for the element M and give an equation.
Task1.5d.2 Define the first ionisation enthalpy for sodium, magnesium, chlorine and neon.
Task1.5d.3 Define the second ionisation enthalpy for lithium, aluminium and oxygen.
1.5e Electronic structureThe graph of successive ionisation energies against ionisation number shows electrons grouped into three energy levels or quantum shells with similar energies.
Electron 1 (ionisation number = 1, removed first) from the third shell, electrons 2-9 in second shell, electrons 10 and 11(removed last) from first shell closest to the nucleus.
Electron 1 is on its own in the outer shell indicating that this element is in group 1. The graph of first ionisation energy against atomic number shows the grouping of electrons into s, p (and d) subshells (orbitals) within the (quantum shells) energy levels. The general trend is for ionisation energy to increase with increasing atomic number across a period but 1st IE B < 1st IE Be and 1st IE O < 1st IE N.
PRACTICE QUSTIONS
Label the diagram with symbols for the elements.
Label electrons removed as s or p.
Write two sentences using the frame: The ionisation energy (increases/decreases) (from left to right across a period/down a group) because (distance from outer electrons to nucleus is increasing/the nuclear charge in increasing)
Saturday, December 13, 2008
Revision sheet
Chemistry of elements
1. In the periodic table the elements are arranged according to the ………………
……………. Number
2. In the periodic table the elements are arranged according to the ………………
……………. No of particle in the nucleus
3. Give any two property of metal
………………………………………………………………………………….
4. Give any two properties of non metal
……………………………………...
5. NAME THE FIRST GROUP ELEMENTS
……………………………………………..
6. Name the second group elements
…………………………………………………….
7. Name the seventh group elements
………………………………….
8. Group 1 element are always has a …………………. Charge
9. Group 2 element are always has a …………………. Charge
9. Group 7 element are always has a …………………. Charge
10. Group 6 element are always has a …………………. Charge
11. Element A put in water to produce a gas B
And an alkali C It will give lilac color to the flame
Name the element A ……………………
Name the gas B ………………………………………….
NAME THE ALKALI C ………………………………………
Write three things that you see in this experiment
………………………………………………………………………………………………
……………………………………………………………………
What is color change shown by the universal indicator in the solution?
………………………………………….
12. Write down the word equation for the above reaction
…………………………………………………..
13. Write down the balanced equation for the above reaction
…………………………………………………..
14. What is the effect of blue litmus red litmus in the beaker
…………………………………………
15. How will you test the gas evolved………………………………
16. Name a non metal for the same period as lithium ………………………..
17. Name another metal in the same period as lithium……………………………..
18. What is similar in the arrangement of electron in alkali metals?
…………………………………………….
19. Why diamond and graphite have similar properties
………………………………………….
20. Why diamond and graphite have different properties
………………………………………….
21why graphite conducts electricity but diamond not. Why?
………………………………………………………………………………………..
22. Give the equation for the burning of graphite in air
…………………………………………………………..
23.give the word equation for the burning of magnesium
……………………………………………………..
24.
halogen
colour
Physical state
chlorine
fluorine
bromine
Periodic table
1. Periodic table is the arrangement of elements according to…………………………
2. Periodic table contains 3/4…………… and 1/4………………………………
3. Define semimetals
……………………………………………………………………………
4. Give examples……………………………….
5. Give uses of metalloids …………………………………………
6. Name a non metal in the same period as lithium……………………………
7. Elements in the same group have …………….. Properties
1. Same 2. Similar
8. Group one elements are having similar properties. Why?
………………………………………………………………………………
9. The vertical columns are called ………………, there are ……….no of groups
10. Horizontal columns are called ………………, there are ……….no of periods
10 give the uses of noble gases ……………………………………….
11. Why noble gases are uncreative……………………………………………………………………………………………………………………………………………….
12. Name the group I elements ………………………………………………………..
13. In group I the reactivity increases down the group …………………………………………………………………………………………………………………………………………………
14why potassium is more reactive than sodium …………………………………………………………………………………………………………………………………………………
15. Name the gas produced when sodium is kept in water. Give effect of litmus in the remaining solution ……………………………………………………………..
16. Write down three observations that you see when sodium is kept in water
……………………………………………………………………………………………
17. Give the reaction of group I metals( sodium, potassium,lithum) in water. Write the balances equation
…………………………………………………………………………………….
…………………………………………………………………………………….
17. Give the reaction of group 2 metals (magnesium, calcium) in water. Write the balances equation ………………………………………………………………………………………………………………………………………………………………
Work sheet
1.Which of the following is ionic compound? Why?
Element
Melting point
A
6754 C
B
31 C
ANS …………………………………………
WHY? ………………………………………………………………………………………………
…………………………………………………………………
2.Draw and explain the structure of diamond
3.Why magnesium chloride is having a high melting point than sodium chloride
4.Molten sodium chloride conducts electricity. Why?
5 .Molten sodium chloride conducts electricity. Why?
6. Graphite is a good conductor of electricity .why?
7. Why sodium chloride do not conduct electricity in solid state
8. Graphite is having a high sublimation point .why?
9. Give the structure of silicon dioxide
……………………………………………………………………………..
10. Give the structure of solid carbon dioxide
……………………………………………………………………………..
11. Give the structure of diamond
……………………………………………………………………………..
12. Give the structure of copper
……………………………………………………………………………..
13. Give the structure of solid methane
……………………………………………………………………………..
14. Give the structure of sodium chloride
……………………………………………………………………………..
15. Give the structure of graphite
……………………………………………………………………………..
16. Why diamond and graphite is having similar properties
……………………………………………………………………………………….
17. Why diamond and graphite is having different properties
……………………………………………………………………………………….
18. Draw the dot and cross diagram
Calcium chloride
Hydrogen
Chlorine molecule
Magnesium oxide
19.Draw the dot and cross diagram of the following. (USE THE PERIODIC TABLE)
N N
O C O
20.A small piece of potassium per magnate is kept in a beaker of water. What will be the observation after some minutes?
………………………………………………………………………………………………
……………………………………………………………………………
Name the phenomenon
………………………………………………………………………………
If you using hot water in place of cold what is the observation
………………………………………………
……………………………………….
Reason ………………………………………………………………………………………………
…………………………………………………………………………………….
21.Draw the particle arrange in gas and liquid
22.Explain in terms of kinetic theory what happen if the temperature of a solid is increased
………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………………….
A pollen grain is kept under beaker of water what do you observe through the microscope
Name………………………………………………………..
Draw the path
1. Draw the bonding in metals
2. Why metals are malleable
3. Why metals conduct electricity
4. Sodium chloride is having a high melting point. why?
5. Draw and explain the structure of diamond
6. Diamond is having a high melting point .why?
7. Why sodium chloride do not conduct electricity in solid state
8. Diamond is having a high sublimation point .why?
9. Why is diamond used for cutting?
10. Draw the structure of sodium chloride and magnesium chloride
11Why magnesium chloride is having a high melting point than sodium chloride
12.Molten sodium chloride conducts electricity. Why?
13.Define alloys. Why they are hard. Draw the structure of alloys
1. In the periodic table the elements are arranged according to the ………………
……………. Number
2. In the periodic table the elements are arranged according to the ………………
……………. No of particle in the nucleus
3. Give any two property of metal
………………………………………………………………………………….
4. Give any two properties of non metal
……………………………………...
5. NAME THE FIRST GROUP ELEMENTS
……………………………………………..
6. Name the second group elements
…………………………………………………….
7. Name the seventh group elements
………………………………….
8. Group 1 element are always has a …………………. Charge
9. Group 2 element are always has a …………………. Charge
9. Group 7 element are always has a …………………. Charge
10. Group 6 element are always has a …………………. Charge
11. Element A put in water to produce a gas B
And an alkali C It will give lilac color to the flame
Name the element A ……………………
Name the gas B ………………………………………….
NAME THE ALKALI C ………………………………………
Write three things that you see in this experiment
………………………………………………………………………………………………
……………………………………………………………………
What is color change shown by the universal indicator in the solution?
………………………………………….
12. Write down the word equation for the above reaction
…………………………………………………..
13. Write down the balanced equation for the above reaction
…………………………………………………..
14. What is the effect of blue litmus red litmus in the beaker
…………………………………………
15. How will you test the gas evolved………………………………
16. Name a non metal for the same period as lithium ………………………..
17. Name another metal in the same period as lithium……………………………..
18. What is similar in the arrangement of electron in alkali metals?
…………………………………………….
19. Why diamond and graphite have similar properties
………………………………………….
20. Why diamond and graphite have different properties
………………………………………….
21why graphite conducts electricity but diamond not. Why?
………………………………………………………………………………………..
22. Give the equation for the burning of graphite in air
…………………………………………………………..
23.give the word equation for the burning of magnesium
……………………………………………………..
24.
halogen
colour
Physical state
chlorine
fluorine
bromine
iodine
25. How will you convert the hydrogen chloride gas into hydrochloric acid?
……………………………………………..
26. A solution in which the hydrochloric acid is dissolved is tested with blue litmus. What do you observe?
…………………………………………………………………………………
27. Hydro choric acid is dissolved in methyl benzene. What do the change in the blue litmus?
……………………………………………………….
28. Describe the laboratory preparation of chlorine from hydro choric acid ( important)
………………………………………………………………………………………
………………………………………………………………………………..
29. Name the oxidizing agent in this reaction……………………………………
30. Give the test for chlorine…………………………………………………………….
31. Chlorine and bromine have similar properties why?
……………………………………………………………………………….
32. Chlorine is passed through a solution of potassium bromide
Name the type of reaction………………………………………….
The solution became red in color. Why?
…………………………………………………………………………..
Explain why no displacement reaction when we add bromine to potassium chloride …………………………………………………………………………………………
Write the ionic equation for the reaction……………………………………………….
Chlorine is bubbled through the potassium iodide solution. What will be the color of the
Final solution ……………………………………………..
Write down the ionic half equation for the reaction …………………………………
What is the color change during the below reaction
……………………………………………………
H2 + Cl 2___________ 2HCl
Name the process of the industrial extraction of oxygen
………………………………………………..
complete the table
78 %
21%
.9%
What is oxidation?
……………………………………….
What is reduction?
…………………………
Give example for oxidizing agent
…………………………………………….
Give example for reducing agent
………………………………………
Complete these reactions ( balance )
S+ O2__________
Fe + O2__________
Mg + O2 _____________
Cu + O 2_________
C +O 2 ___________
CH4 + O2_____________
Describe an experiment to determine the % of oxygen in air (copper turning )
Give the reaction of sulphur dioxide with water
…………………………………………….
Give two equations for the reactions in contact process ( impor)
…………………………………………………………….
……………………………………………………….
Which one is reversible reaction?
……………………………………………
Describe the laboratory preparation of carbon dioxide
Give the equation for reaction
……………………………………………………………………………….
Give the uses of carbon dioxide
………………………………………………………………
What is dry ice?
……………………………
Give two uses of dry ice
………………………………………..
Give the reaction of nitrogen with oxygen ( 2 equations)
…………………………………………………….
…………………………………………………..
Give the conditions needed for the rusting of iron
……………………………………………………………………………….
Give method for the prevention of rusting
……………………………………………………………………………..
How this method work in preventing rusting of iron
……………………………………………………………………………
Explain sacrificial protection. Give example.
……………………………………………………………………………….
Sodium can be prepared by electrolysis only were as iron fro blast furnace give your answer (in terms of reactivity)
………………………………………………………………………………………
Give the allotropes of sulphur
……………………………………………………………………………
Name the process by which nitrogen is industrially extracted
………………………………………………………………..
Give one use of nitrogen
……………………………………….
Why nitrogen is used for protecting dry foods
……………………………………………….
Describe the laboratory preparation of ammonia (page-244 text book)
Give the equation for the preparation
………………………………………………………………..
Give physical properties of ammonia
………………………………………………………………………………………
ammonia dissolve in water forms …………………………………………
give the equation…………………………………………
Give the uses of ammonia
to prepare nitric acid
to prepare fertilizer
Give the reaction between sulphuric acid and ammonia ( important)
NH 3 + H2SO4 ----------- (NH4)2 SO4 ( AMMONIUM SULPHATE)
Give the reaction between NITRIC acid and ammonia ( important)
NH3+ HNO3 ------- NH 4 NO3( AMMONIUM NITRATE)
What are the elements present in NPK FERTILIZER?
………………………………………………………..
Describe the chemical test for water
……………………………………………………………..
Describe the physical test for water
………………………………………………………………
Describe the laboratory preparation of hydrogen
Name the gas produced when a metal react with dilute acid
…………………………………………………….
Give the test for this gas
…………………………………………………….
Give word equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give balanced equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give one example for transition metal which is used as catalyst
…………………………………………………….
Iron react with chlorine to form a yellow( brown) color solid
Name it…………………………….
Give equation for the reaction
…………………………………………………..
Give the color of ion with sodium hydroxide
Fe2+ …………………
Fe3+…………………
Cu2+ ………………
Name the compounds……………………………………………………..
Write down the action of iron over steam. Give the balanced equation ( notes)
…………………………………………………………………..
Write down the action of iron hydrogen chloride gas Give the balanced equation (notes)
………………………………………………………………………….
Color
Copper oxide-red
Copper hydroxide – blue
Copper nitrate –green
Copper sulfate- blue
Copper carbonate-black
Copper chloride- green
Write down the reaction of ammonia with copper (II) ions ( important)
………………………………………………………………………………
Give the properties of transitional metals( answer given below)
1) variable valency
2) colored compounds
3) catalytic property
Give example for Fe shows variable valency
…………………………………………………………………….
What is reactivity series? Give the order
…………………………………………………………………………………
What do you know about the reactivity of and displacement reaction?
………………………………………………………………………….
A copper dipped in silver nitrate solution .What is observation that you see
After 2hours…………………………………………………….
.
Iron nail dipped in copper sulphate ……………………………………
Silver nail dipped ii copper sulphate……………………………………
Copper nail in iron sulphate……………………………..
(Important portion)
Cat ion
Flame color
Lithium(Li+)
Sodium (Na+)
Potassium( k+)
Calcium( ca2+)
gas
test
result
oxygen
carbon dioxide
hydrogen
chlorine
Sulpher dioxide
ammonia
Test for anions
TEST
result
Chloride(Cl-)
Bromide(Br -
Iodide (I-)
Sulphate(so42+)
Sulphite( S2_)
Carbonate(CO32_)
Test for cat ion
TEST
result
Ammonia (NH4+)
Copper (Cu2+)
Iron( Fe2+)
Iron ( Fe3+)
1. Name the three particles in the atom
2. ……………………………….. is the negative particle in the atom
3. …………………………………..is the positive particle in an atom
4 …………………………………… is the neutral particle in the atom
5. In an atom the number of ……………………………and
…………………………. are same
6. Name the particle out side the nucleus
7. Name the negligible particle in an atom
8. The mass of the atom is concentrated at …………………
9. Define atomic number
………………………………………………………………………………………………
…………………………………………………………………………………………….
10. Define mass number
………………………………………………………………………………………………
………………………………………………………………………………………………
11. find the 1) mass number2) atomic number 3) no. of neutron 4) no. of protons 5) no. of electron
12. Define isotope
………………………………………………………………………………………………
………………………………………………………………………………………………
13. Which one pair of the following is isotope?
Ans…………………………………
14. The following pair represents some gases
Which is 1) mono atomic …………………………..
2) Water molecule …………………………………
3) Mixture of gases ……………………………….
4) Elements ……………………………………….
5) Compounds………………………………………
6) Diatomic molecule ………………………………
15. Define element. Give example
……………………………………………………………………………………………
16. Carbon dioxide is a compound. What is a compound?
…………………………………………………………………………………………
17. Which one of the following belongs to the second group?
Ans : ……………………………………….
Give your answer
……………………………………………………………………………………………
Name the element ……………………………………………….(use periodic table)
18. find the outer electrons in the following atoms
1) bromine 2) carbon 3) boron
19) How many outer electrons are their in radium, iodine and radon
19) The below fig shows the nucleus of an atom
Find the atomic no……………………………
Mass no …………………………………
No of electrons ………………………………….
This atom belongs to ……………………………..group of the periodic table.
Give your answer for the above statement ………………………………………………………………………………………………
…………………………………………………………….
Work sheet 2
States of matter
define matter
complete the fig
A ……………………………
B……………………………..
c………………………………
D……………………………
E………………………………
complete the arrangement of solid particle
Explain what happens when a solid changes to liquid. give your answer using kinetic theory.
define sublimation
Define diffusion. Give example.
A small piece of potassium permanganate is kept under a beaker of water for one hour. What do you see after one hour?
A small piece of copper sulphate ( blue color) is kept under a beaker of water for one hour. What do you see after one hour?
Bromine vapor is kept in a jar and the lid is removed. What do you see after 10
min
A pollen grain is kept on the surface of water and viewed through the micro scope.
What do you see ……………………………………………………………………
Draw the path of the pollen grain
What is dry ice? How is it prepared? Give the use.
A smoke particle in air is viewed through the micro scope.
What do you see ……………………………………………………………………
Draw the path of the smoke particle
1. How will you find the purity of a solid?
…………………………………………………………………………………………
…………………………………………………………………………………………
………………………………………………………………………………… 2
2. Table below shows the melting and boiling points of some substance
SUBSTANCE
MELTING POINT
BOILING POINT
A
-123
-56
B
0
100
C
45
111
D
-23
23
E
-11
1
1) which one of the following is representing water
………………………………………………..
2) which one is a solid at room temperature
………………………………………………….
3) which substances are liquids at -15 c
………………………………………………….
4) which substance a liquid at short range of temperature
……………………………………………………….
3) The below graph represent the meting of a solid
What is the physical state at?
A ……………………………………………
B ………………………………………………..
C …………………………………………….
D ……………………………………………….
WHAT IS THE MELTING POINT
…………………………………
What is the boiling point?
……………………………………….
Some salt is added to the solid what will be the observation in boiling and melting point
……………………………………………………………………..
Reason ………………………………………………………………………………………………
A small piece of potassium per magnate is kept in a beaker of water. What will be the observation after some minutes?
………………………………………………………………………………………………
……………………………………………………………………………
Name the phenomenon
………………………………………………………………………………
If you using hot water in place of cold what is the observation
………………………………………………………………………………………………
……………………………………….
Reason ………………………………………………………………………………………………
…………………………………………………………………………………….
Draw the particle arrange in gas and liquid
Explain in terms of kinetic theory what happen if the temperature of a solid is increased
………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………………….
In the below diagram shows piece of potassium per magnate kept under water
(The substance is purple in color)
What do you see after 1hour?
Observation …………………………………………………………………………………………
Name it and define
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………….
Explain what happen if use hot water
………………………………………………………………………………
Reason ………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………
The below apparatus shows diffusion
1) Which one diffuse fast? Why?..........................................................................................................................
...................................................
2) what happens if you hot solution
………………………………………………………………………………………
…………………………
A pollen grain is kept under beaker of water what do you observe through the microscope
Name………………………………………………………..
Draw the path
Atomic structure
1) Which one is mono atomic gas ………………….
2) Which one are an element ……………………………………
3) Which one are compounds ……………………………………….
4) Which one represent solid…………………………….
5) Which one represent liquid……………………..
6) Which one represent water molecule ……………………
7) Which one represent oxygen molecule ………………………..
8) Which one represent argon …………………………………
1. How will you find the purity of a solid?
…………………………………………………………………………………………
…………………………………………………………………………………………
………………………………………………………………………………… 2
2. Table below shows the melting and boiling points of some substance
SUBSTANCE
MELTING POINT
BOILING POINT
A
-123
-56
B
0
100
C
45
111
D
-23
23
E
-11
1
5) which one of the following is representing water
………………………………………………..
6) which one is a solid at room temperature
………………………………………………….
7) which substances are liquids at -15 c
………………………………………………….
8) which substance a liquid at short range of temperature
……………………………………………………….
3) The below graph represent the meting of a solid
What is the physical state at?
A ……………………………………………
B ………………………………………………..
C …………………………………………….
D ……………………………………………….
WHAT IS THE MELTING POINT
…………………………………
What is the boiling point?
……………………………………….
Some salt is added to the solid what will be the observation in boiling and melting point
……………………………………………………………………..
Reason ………………………………………………………………………………………………
A small piece of potassium per magnate is kept in a beaker of water. What will be the observation after some minutes?
………………………………………………………………………………………………
……………………………………………………………………………
Name the phenomenon
………………………………………………………………………………
If you using hot water in place of cold what is the observation
………………………………………………………………………………………………
……………………………………….
Reason ………………………………………………………………………………………………
…………………………………………………………………………………….
Draw the particle arrange in gas and liquid
Explain in terms of kinetic theory what happen if the temperature of a solid is increased
………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………………….
In the below diagram shows piece of potassium per magnate kept under water
(The substance is purple in color)
What do you see after 1hour?
Observation …………………………………………………………………………………………
Name it and define
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………….
Explain what happen if use hot water
………………………………………………………………………………
Reason ………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………
The below apparatus shows diffusion
3) Which one diffuse fast? Why?..........................................................................................................................
...................................................
4) what happens if you hot solution
………………………………………………………………………………………
…………………………
A pollen grain is kept under beaker of water what do you observe through the microscope
Name………………………………………………………..
Draw the path
Atomic structure
9) Which one is mono atomic gas ………………….
10) Which one are an element ……………………………………
11) Which one are compounds ……………………………………….
12) Which one represent solid…………………………….
13) Which one represent liquid……………………..
14) Which one represent water molecule ……………………
15) Which one represent oxygen molecule ………………………..
16) Which one represent argon …………………………………
define ionic bonding
………………………………………………………………………………………………
…………………………………………………………………..
draw the dot and cross diagram
Lithium fluoride
Calcium chloride
Magnesium oxide
Calcium chloride
Potassium chloride
Magnesium oxide is having a high melting point than sodium chloride why?
Which one of the following is an ion?
Draw the electronic configuration of chloride ion
Draw the electronic configuration of following
Define covalent bonding
Draw the cross and dot diagram for the following
Hydrogen
Water
Methane
Hydrogen chloride
Nitrogen
Ehane
Carbondixide
iodine
25. How will you convert the hydrogen chloride gas into hydrochloric acid?
……………………………………………..
26. A solution in which the hydrochloric acid is dissolved is tested with blue litmus. What do you observe?
…………………………………………………………………………………
27. Hydro choric acid is dissolved in methyl benzene. What do the change in the blue litmus?
……………………………………………………….
28. Describe the laboratory preparation of chlorine from hydro choric acid ( important)
………………………………………………………………………………………
………………………………………………………………………………..
29. Name the oxidizing agent in this reaction……………………………………
30. Give the test for chlorine…………………………………………………………….
31. Chlorine and bromine have similar properties why?
……………………………………………………………………………….
32. Chlorine is passed through a solution of potassium bromide
Name the type of reaction………………………………………….
The solution became red in color. Why?
…………………………………………………………………………..
Explain why no displacement reaction when we add bromine to potassium chloride …………………………………………………………………………………………
Write the ionic equation for the reaction……………………………………………….
Chlorine is bubbled through the potassium iodide solution. What will be the color of the
Final solution ……………………………………………..
Write down the ionic half equation for the reaction …………………………………
What is the color change during the below reaction
……………………………………………………
H2 + Cl 2___________ 2HCl
Name the process of the industrial extraction of oxygen
………………………………………………..
complete the table
78 %
21%
.9%
What is oxidation?
……………………………………….
What is reduction?
…………………………
Give example for oxidizing agent
…………………………………………….
Give example for reducing agent
………………………………………
Complete these reactions ( balance )
S+ O2__________
Fe + O2__________
Mg + O2 _____________
Cu + O 2_________
C +O 2 ___________
CH4 + O2_____________
Describe an experiment to determine the % of oxygen in air (copper turning )
Give the reaction of sulphur dioxide with water
…………………………………………….
Give two equations for the reactions in contact process ( impor)
…………………………………………………………….
……………………………………………………….
Which one is reversible reaction?
……………………………………………
Describe the laboratory preparation of carbon dioxide
Give the equation for reaction
……………………………………………………………………………….
Give the uses of carbon dioxide
………………………………………………………………
What is dry ice?
……………………………
Give two uses of dry ice
………………………………………..
Give the reaction of nitrogen with oxygen ( 2 equations)
…………………………………………………….
…………………………………………………..
Give the conditions needed for the rusting of iron
……………………………………………………………………………….
Give method for the prevention of rusting
……………………………………………………………………………..
How this method work in preventing rusting of iron
……………………………………………………………………………
Explain sacrificial protection. Give example.
……………………………………………………………………………….
Sodium can be prepared by electrolysis only were as iron fro blast furnace give your answer (in terms of reactivity)
………………………………………………………………………………………
Give the allotropes of sulphur
……………………………………………………………………………
Name the process by which nitrogen is industrially extracted
………………………………………………………………..
Give one use of nitrogen
……………………………………….
Why nitrogen is used for protecting dry foods
……………………………………………….
Describe the laboratory preparation of ammonia (page-244 text book)
Give the equation for the preparation
………………………………………………………………..
Give physical properties of ammonia
………………………………………………………………………………………
ammonia dissolve in water forms …………………………………………
give the equation…………………………………………
Give the uses of ammonia
to prepare nitric acid
to prepare fertilizer
Give the reaction between sulphuric acid and ammonia ( important)
NH 3 + H2SO4 ----------- (NH4)2 SO4 ( AMMONIUM SULPHATE)
Give the reaction between NITRIC acid and ammonia ( important)
NH3+ HNO3 ------- NH 4 NO3( AMMONIUM NITRATE)
What are the elements present in NPK FERTILIZER?
………………………………………………………..
Describe the chemical test for water
……………………………………………………………..
Describe the physical test for water
………………………………………………………………
Describe the laboratory preparation of hydrogen
Name the gas produced when a metal react with dilute acid
…………………………………………………….
Give the test for this gas
…………………………………………………….
Give word equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give balanced equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give one example for transition metal which is used as catalyst
…………………………………………………….
Iron react with chlorine to form a yellow( brown) color solid
Name it…………………………….
Give equation for the reaction
…………………………………………………..
Give the color of ion with sodium hydroxide
Fe2+ …………………
Fe3+…………………
Cu2+ ………………
Name the compounds……………………………………………………..
Write down the action of iron over steam. Give the balanced equation ( notes)
…………………………………………………………………..
Write down the action of iron hydrogen chloride gas Give the balanced equation (notes)
………………………………………………………………………….
Color
Copper oxide-red
Copper hydroxide – blue
Copper nitrate –green
Copper sulfate- blue
Copper carbonate-black
Copper chloride- green
Write down the reaction of ammonia with copper (II) ions ( important)
………………………………………………………………………………
Give the properties of transitional metals( answer given below)
1) variable valency
2) colored compounds
3) catalytic property
Give example for Fe shows variable valency
…………………………………………………………………….
What is reactivity series? Give the order
…………………………………………………………………………………
What do you know about the reactivity of and displacement reaction?
………………………………………………………………………….
A copper dipped in silver nitrate solution .What is observation that you see
After 2hours…………………………………………………….
.
Iron nail dipped in copper sulphate ……………………………………
Silver nail dipped ii copper sulphate……………………………………
Copper nail in iron sulphate……………………………..
(Important portion)
Cat ion
Flame color
Lithium(Li+)
Sodium (Na+)
Potassium( k+)
Calcium( ca2+)
gas
test
result
oxygen
carbon dioxide
hydrogen
chlorine
Sulpher dioxide
ammonia
Test for anions
TEST
result
Chloride(Cl-)
Bromide(Br -
Iodide (I-)
Sulphate(so42+)
Sulphite( S2_)
Carbonate(CO32_)
Test for cat ion
TEST
result
Ammonia (NH4+)
Copper (Cu2+)
Iron( Fe2+)
Iron ( Fe3+)
Chemistry of elements
1. In the periodic table the elements are arranged according to the ………………
……………. Number
2. In the periodic table the elements are arranged according to the ………………
……………. No of particle in the nucleus
3. Give any two property of metal
………………………………………………………………………………….
4. Give any two properties of non metal
……………………………………...
5. NAME THE FIRST GROUP ELEMENTS
……………………………………………..
6. Name the second group elements
…………………………………………………….
7. Name the seventh group elements
………………………………….
8. Group 1 element are always has a …………………. Charge
9. Group 2 element are always has a …………………. Charge
9. Group 7 element are always has a …………………. Charge
10. Group 6 element are always has a …………………. Charge
11. Element A put in water to produce a gas B
And an alkali C It will give lilac color to the flame
Name the element A ……………………
Name the gas B ………………………………………….
NAME THE ALKALI C ………………………………………
Write three things that you see in this experiment
………………………………………………………………………………………………
……………………………………………………………………
What is color change shown by the universal indicator in the solution?
………………………………………….
12. Write down the word equation for the above reaction
…………………………………………………..
13. Write down the balanced equation for the above reaction
…………………………………………………..
14. What is the effect of blue litmus red litmus in the beaker
…………………………………………
15. How will you test the gas evolved………………………………
16. Name a non metal for the same period as lithium ………………………..
17. Name another metal in the same period as lithium……………………………..
18. What is similar in the arrangement of electron in alkali metals?
…………………………………………….
19. Why diamond and graphite have similar properties
………………………………………….
20. Why diamond and graphite have different properties
………………………………………….
21why graphite conducts electricity but diamond not. Why?
………………………………………………………………………………………..
22. Give the equation for the burning of graphite in air
…………………………………………………………..
23.give the word equation for the burning of magnesium
……………………………………………………..
24.
halogen
colour
Physical state
chlorine
fluorine
bromine
Periodic table
1. Periodic table is the arrangement of elements according to…………………………
2. Periodic table contains 3/4…………… and 1/4………………………………
3. Define semimetals
……………………………………………………………………………
4. Give examples……………………………….
5. Give uses of metalloids …………………………………………
6. Name a non metal in the same period as lithium……………………………
7. Elements in the same group have …………….. Properties
1. Same 2. Similar
8. Group one elements are having similar properties. Why?
………………………………………………………………………………
9. The vertical columns are called ………………, there are ……….no of groups
10. Horizontal columns are called ………………, there are ……….no of periods
10 give the uses of noble gases ……………………………………….
11. Why noble gases are uncreative……………………………………………………………………………………………………………………………………………….
12. Name the group I elements ………………………………………………………..
13. In group I the reactivity increases down the group …………………………………………………………………………………………………………………………………………………
14why potassium is more reactive than sodium …………………………………………………………………………………………………………………………………………………
15. Name the gas produced when sodium is kept in water. Give effect of litmus in the remaining solution ……………………………………………………………..
16. Write down three observations that you see when sodium is kept in water
……………………………………………………………………………………………
17. Give the reaction of group I metals( sodium, potassium,lithum) in water. Write the balances equation
…………………………………………………………………………………….
…………………………………………………………………………………….
17. Give the reaction of group 2 metals (magnesium, calcium) in water. Write the balances equation ………………………………………………………………………………………………………………………………………………………………
Work sheet
1.Which of the following is ionic compound? Why?
Element
Melting point
A
6754 C
B
31 C
ANS …………………………………………
WHY? ………………………………………………………………………………………………
…………………………………………………………………
2.Draw and explain the structure of diamond
3.Why magnesium chloride is having a high melting point than sodium chloride
4.Molten sodium chloride conducts electricity. Why?
5 .Molten sodium chloride conducts electricity. Why?
6. Graphite is a good conductor of electricity .why?
7. Why sodium chloride do not conduct electricity in solid state
8. Graphite is having a high sublimation point .why?
9. Give the structure of silicon dioxide
……………………………………………………………………………..
10. Give the structure of solid carbon dioxide
……………………………………………………………………………..
11. Give the structure of diamond
……………………………………………………………………………..
12. Give the structure of copper
……………………………………………………………………………..
13. Give the structure of solid methane
……………………………………………………………………………..
14. Give the structure of sodium chloride
……………………………………………………………………………..
15. Give the structure of graphite
……………………………………………………………………………..
16. Why diamond and graphite is having similar properties
……………………………………………………………………………………….
17. Why diamond and graphite is having different properties
……………………………………………………………………………………….
18. Draw the dot and cross diagram
Calcium chloride
Hydrogen
Chlorine molecule
Magnesium oxide
19.Draw the dot and cross diagram of the following. (USE THE PERIODIC TABLE)
N N
O C O
20.A small piece of potassium per magnate is kept in a beaker of water. What will be the observation after some minutes?
………………………………………………………………………………………………
……………………………………………………………………………
Name the phenomenon
………………………………………………………………………………
If you using hot water in place of cold what is the observation
………………………………………………
……………………………………….
Reason ………………………………………………………………………………………………
…………………………………………………………………………………….
21.Draw the particle arrange in gas and liquid
22.Explain in terms of kinetic theory what happen if the temperature of a solid is increased
………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………………….
A pollen grain is kept under beaker of water what do you observe through the microscope
Name………………………………………………………..
Draw the path
1. Draw the bonding in metals
2. Why metals are malleable
3. Why metals conduct electricity
4. Sodium chloride is having a high melting point. why?
5. Draw and explain the structure of diamond
6. Diamond is having a high melting point .why?
7. Why sodium chloride do not conduct electricity in solid state
8. Diamond is having a high sublimation point .why?
9. Why is diamond used for cutting?
10. Draw the structure of sodium chloride and magnesium chloride
11Why magnesium chloride is having a high melting point than sodium chloride
12.Molten sodium chloride conducts electricity. Why?
13.Define alloys. Why they are hard. Draw the structure of alloys
1. In the periodic table the elements are arranged according to the ………………
……………. Number
2. In the periodic table the elements are arranged according to the ………………
……………. No of particle in the nucleus
3. Give any two property of metal
………………………………………………………………………………….
4. Give any two properties of non metal
……………………………………...
5. NAME THE FIRST GROUP ELEMENTS
……………………………………………..
6. Name the second group elements
…………………………………………………….
7. Name the seventh group elements
………………………………….
8. Group 1 element are always has a …………………. Charge
9. Group 2 element are always has a …………………. Charge
9. Group 7 element are always has a …………………. Charge
10. Group 6 element are always has a …………………. Charge
11. Element A put in water to produce a gas B
And an alkali C It will give lilac color to the flame
Name the element A ……………………
Name the gas B ………………………………………….
NAME THE ALKALI C ………………………………………
Write three things that you see in this experiment
………………………………………………………………………………………………
……………………………………………………………………
What is color change shown by the universal indicator in the solution?
………………………………………….
12. Write down the word equation for the above reaction
…………………………………………………..
13. Write down the balanced equation for the above reaction
…………………………………………………..
14. What is the effect of blue litmus red litmus in the beaker
…………………………………………
15. How will you test the gas evolved………………………………
16. Name a non metal for the same period as lithium ………………………..
17. Name another metal in the same period as lithium……………………………..
18. What is similar in the arrangement of electron in alkali metals?
…………………………………………….
19. Why diamond and graphite have similar properties
………………………………………….
20. Why diamond and graphite have different properties
………………………………………….
21why graphite conducts electricity but diamond not. Why?
………………………………………………………………………………………..
22. Give the equation for the burning of graphite in air
…………………………………………………………..
23.give the word equation for the burning of magnesium
……………………………………………………..
24.
halogen
colour
Physical state
chlorine
fluorine
bromine
iodine
25. How will you convert the hydrogen chloride gas into hydrochloric acid?
……………………………………………..
26. A solution in which the hydrochloric acid is dissolved is tested with blue litmus. What do you observe?
…………………………………………………………………………………
27. Hydro choric acid is dissolved in methyl benzene. What do the change in the blue litmus?
……………………………………………………….
28. Describe the laboratory preparation of chlorine from hydro choric acid ( important)
………………………………………………………………………………………
………………………………………………………………………………..
29. Name the oxidizing agent in this reaction……………………………………
30. Give the test for chlorine…………………………………………………………….
31. Chlorine and bromine have similar properties why?
……………………………………………………………………………….
32. Chlorine is passed through a solution of potassium bromide
Name the type of reaction………………………………………….
The solution became red in color. Why?
…………………………………………………………………………..
Explain why no displacement reaction when we add bromine to potassium chloride …………………………………………………………………………………………
Write the ionic equation for the reaction……………………………………………….
Chlorine is bubbled through the potassium iodide solution. What will be the color of the
Final solution ……………………………………………..
Write down the ionic half equation for the reaction …………………………………
What is the color change during the below reaction
……………………………………………………
H2 + Cl 2___________ 2HCl
Name the process of the industrial extraction of oxygen
………………………………………………..
complete the table
78 %
21%
.9%
What is oxidation?
……………………………………….
What is reduction?
…………………………
Give example for oxidizing agent
…………………………………………….
Give example for reducing agent
………………………………………
Complete these reactions ( balance )
S+ O2__________
Fe + O2__________
Mg + O2 _____________
Cu + O 2_________
C +O 2 ___________
CH4 + O2_____________
Describe an experiment to determine the % of oxygen in air (copper turning )
Give the reaction of sulphur dioxide with water
…………………………………………….
Give two equations for the reactions in contact process ( impor)
…………………………………………………………….
……………………………………………………….
Which one is reversible reaction?
……………………………………………
Describe the laboratory preparation of carbon dioxide
Give the equation for reaction
……………………………………………………………………………….
Give the uses of carbon dioxide
………………………………………………………………
What is dry ice?
……………………………
Give two uses of dry ice
………………………………………..
Give the reaction of nitrogen with oxygen ( 2 equations)
…………………………………………………….
…………………………………………………..
Give the conditions needed for the rusting of iron
……………………………………………………………………………….
Give method for the prevention of rusting
……………………………………………………………………………..
How this method work in preventing rusting of iron
……………………………………………………………………………
Explain sacrificial protection. Give example.
……………………………………………………………………………….
Sodium can be prepared by electrolysis only were as iron fro blast furnace give your answer (in terms of reactivity)
………………………………………………………………………………………
Give the allotropes of sulphur
……………………………………………………………………………
Name the process by which nitrogen is industrially extracted
………………………………………………………………..
Give one use of nitrogen
……………………………………….
Why nitrogen is used for protecting dry foods
……………………………………………….
Describe the laboratory preparation of ammonia (page-244 text book)
Give the equation for the preparation
………………………………………………………………..
Give physical properties of ammonia
………………………………………………………………………………………
ammonia dissolve in water forms …………………………………………
give the equation…………………………………………
Give the uses of ammonia
to prepare nitric acid
to prepare fertilizer
Give the reaction between sulphuric acid and ammonia ( important)
NH 3 + H2SO4 ----------- (NH4)2 SO4 ( AMMONIUM SULPHATE)
Give the reaction between NITRIC acid and ammonia ( important)
NH3+ HNO3 ------- NH 4 NO3( AMMONIUM NITRATE)
What are the elements present in NPK FERTILIZER?
………………………………………………………..
Describe the chemical test for water
……………………………………………………………..
Describe the physical test for water
………………………………………………………………
Describe the laboratory preparation of hydrogen
Name the gas produced when a metal react with dilute acid
…………………………………………………….
Give the test for this gas
…………………………………………………….
Give word equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give balanced equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give one example for transition metal which is used as catalyst
…………………………………………………….
Iron react with chlorine to form a yellow( brown) color solid
Name it…………………………….
Give equation for the reaction
…………………………………………………..
Give the color of ion with sodium hydroxide
Fe2+ …………………
Fe3+…………………
Cu2+ ………………
Name the compounds……………………………………………………..
Write down the action of iron over steam. Give the balanced equation ( notes)
…………………………………………………………………..
Write down the action of iron hydrogen chloride gas Give the balanced equation (notes)
………………………………………………………………………….
Color
Copper oxide-red
Copper hydroxide – blue
Copper nitrate –green
Copper sulfate- blue
Copper carbonate-black
Copper chloride- green
Write down the reaction of ammonia with copper (II) ions ( important)
………………………………………………………………………………
Give the properties of transitional metals( answer given below)
1) variable valency
2) colored compounds
3) catalytic property
Give example for Fe shows variable valency
…………………………………………………………………….
What is reactivity series? Give the order
…………………………………………………………………………………
What do you know about the reactivity of and displacement reaction?
………………………………………………………………………….
A copper dipped in silver nitrate solution .What is observation that you see
After 2hours…………………………………………………….
.
Iron nail dipped in copper sulphate ……………………………………
Silver nail dipped ii copper sulphate……………………………………
Copper nail in iron sulphate……………………………..
(Important portion)
Cat ion
Flame color
Lithium(Li+)
Sodium (Na+)
Potassium( k+)
Calcium( ca2+)
gas
test
result
oxygen
carbon dioxide
hydrogen
chlorine
Sulpher dioxide
ammonia
Test for anions
TEST
result
Chloride(Cl-)
Bromide(Br -
Iodide (I-)
Sulphate(so42+)
Sulphite( S2_)
Carbonate(CO32_)
Test for cat ion
TEST
result
Ammonia (NH4+)
Copper (Cu2+)
Iron( Fe2+)
Iron ( Fe3+)
1. Name the three particles in the atom
2. ……………………………….. is the negative particle in the atom
3. …………………………………..is the positive particle in an atom
4 …………………………………… is the neutral particle in the atom
5. In an atom the number of ……………………………and
…………………………. are same
6. Name the particle out side the nucleus
7. Name the negligible particle in an atom
8. The mass of the atom is concentrated at …………………
9. Define atomic number
………………………………………………………………………………………………
…………………………………………………………………………………………….
10. Define mass number
………………………………………………………………………………………………
………………………………………………………………………………………………
11. find the 1) mass number2) atomic number 3) no. of neutron 4) no. of protons 5) no. of electron
12. Define isotope
………………………………………………………………………………………………
………………………………………………………………………………………………
13. Which one pair of the following is isotope?
Ans…………………………………
14. The following pair represents some gases
Which is 1) mono atomic …………………………..
2) Water molecule …………………………………
3) Mixture of gases ……………………………….
4) Elements ……………………………………….
5) Compounds………………………………………
6) Diatomic molecule ………………………………
15. Define element. Give example
……………………………………………………………………………………………
16. Carbon dioxide is a compound. What is a compound?
…………………………………………………………………………………………
17. Which one of the following belongs to the second group?
Ans : ……………………………………….
Give your answer
……………………………………………………………………………………………
Name the element ……………………………………………….(use periodic table)
18. find the outer electrons in the following atoms
1) bromine 2) carbon 3) boron
19) How many outer electrons are their in radium, iodine and radon
19) The below fig shows the nucleus of an atom
Find the atomic no……………………………
Mass no …………………………………
No of electrons ………………………………….
This atom belongs to ……………………………..group of the periodic table.
Give your answer for the above statement ………………………………………………………………………………………………
…………………………………………………………….
Work sheet 2
States of matter
define matter
complete the fig
A ……………………………
B……………………………..
c………………………………
D……………………………
E………………………………
complete the arrangement of solid particle
Explain what happens when a solid changes to liquid. give your answer using kinetic theory.
define sublimation
Define diffusion. Give example.
A small piece of potassium permanganate is kept under a beaker of water for one hour. What do you see after one hour?
A small piece of copper sulphate ( blue color) is kept under a beaker of water for one hour. What do you see after one hour?
Bromine vapor is kept in a jar and the lid is removed. What do you see after 10
min
A pollen grain is kept on the surface of water and viewed through the micro scope.
What do you see ……………………………………………………………………
Draw the path of the pollen grain
What is dry ice? How is it prepared? Give the use.
A smoke particle in air is viewed through the micro scope.
What do you see ……………………………………………………………………
Draw the path of the smoke particle
1. How will you find the purity of a solid?
…………………………………………………………………………………………
…………………………………………………………………………………………
………………………………………………………………………………… 2
2. Table below shows the melting and boiling points of some substance
SUBSTANCE
MELTING POINT
BOILING POINT
A
-123
-56
B
0
100
C
45
111
D
-23
23
E
-11
1
1) which one of the following is representing water
………………………………………………..
2) which one is a solid at room temperature
………………………………………………….
3) which substances are liquids at -15 c
………………………………………………….
4) which substance a liquid at short range of temperature
……………………………………………………….
3) The below graph represent the meting of a solid
What is the physical state at?
A ……………………………………………
B ………………………………………………..
C …………………………………………….
D ……………………………………………….
WHAT IS THE MELTING POINT
…………………………………
What is the boiling point?
……………………………………….
Some salt is added to the solid what will be the observation in boiling and melting point
……………………………………………………………………..
Reason ………………………………………………………………………………………………
A small piece of potassium per magnate is kept in a beaker of water. What will be the observation after some minutes?
………………………………………………………………………………………………
……………………………………………………………………………
Name the phenomenon
………………………………………………………………………………
If you using hot water in place of cold what is the observation
………………………………………………………………………………………………
……………………………………….
Reason ………………………………………………………………………………………………
…………………………………………………………………………………….
Draw the particle arrange in gas and liquid
Explain in terms of kinetic theory what happen if the temperature of a solid is increased
………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………………….
In the below diagram shows piece of potassium per magnate kept under water
(The substance is purple in color)
What do you see after 1hour?
Observation …………………………………………………………………………………………
Name it and define
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………….
Explain what happen if use hot water
………………………………………………………………………………
Reason ………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………
The below apparatus shows diffusion
1) Which one diffuse fast? Why?..........................................................................................................................
...................................................
2) what happens if you hot solution
………………………………………………………………………………………
…………………………
A pollen grain is kept under beaker of water what do you observe through the microscope
Name………………………………………………………..
Draw the path
Atomic structure
1) Which one is mono atomic gas ………………….
2) Which one are an element ……………………………………
3) Which one are compounds ……………………………………….
4) Which one represent solid…………………………….
5) Which one represent liquid……………………..
6) Which one represent water molecule ……………………
7) Which one represent oxygen molecule ………………………..
8) Which one represent argon …………………………………
1. How will you find the purity of a solid?
…………………………………………………………………………………………
…………………………………………………………………………………………
………………………………………………………………………………… 2
2. Table below shows the melting and boiling points of some substance
SUBSTANCE
MELTING POINT
BOILING POINT
A
-123
-56
B
0
100
C
45
111
D
-23
23
E
-11
1
5) which one of the following is representing water
………………………………………………..
6) which one is a solid at room temperature
………………………………………………….
7) which substances are liquids at -15 c
………………………………………………….
8) which substance a liquid at short range of temperature
……………………………………………………….
3) The below graph represent the meting of a solid
What is the physical state at?
A ……………………………………………
B ………………………………………………..
C …………………………………………….
D ……………………………………………….
WHAT IS THE MELTING POINT
…………………………………
What is the boiling point?
……………………………………….
Some salt is added to the solid what will be the observation in boiling and melting point
……………………………………………………………………..
Reason ………………………………………………………………………………………………
A small piece of potassium per magnate is kept in a beaker of water. What will be the observation after some minutes?
………………………………………………………………………………………………
……………………………………………………………………………
Name the phenomenon
………………………………………………………………………………
If you using hot water in place of cold what is the observation
………………………………………………………………………………………………
……………………………………….
Reason ………………………………………………………………………………………………
…………………………………………………………………………………….
Draw the particle arrange in gas and liquid
Explain in terms of kinetic theory what happen if the temperature of a solid is increased
………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………………….
In the below diagram shows piece of potassium per magnate kept under water
(The substance is purple in color)
What do you see after 1hour?
Observation …………………………………………………………………………………………
Name it and define
………………………………………………………………………………………………
………………………………………………………………………………………………
……………………………………………….
Explain what happen if use hot water
………………………………………………………………………………
Reason ………………………………………………………………………………………………
………………………………………………………………………………………………
………………………………………………………………
The below apparatus shows diffusion
3) Which one diffuse fast? Why?..........................................................................................................................
...................................................
4) what happens if you hot solution
………………………………………………………………………………………
…………………………
A pollen grain is kept under beaker of water what do you observe through the microscope
Name………………………………………………………..
Draw the path
Atomic structure
9) Which one is mono atomic gas ………………….
10) Which one are an element ……………………………………
11) Which one are compounds ……………………………………….
12) Which one represent solid…………………………….
13) Which one represent liquid……………………..
14) Which one represent water molecule ……………………
15) Which one represent oxygen molecule ………………………..
16) Which one represent argon …………………………………
define ionic bonding
………………………………………………………………………………………………
…………………………………………………………………..
draw the dot and cross diagram
Lithium fluoride
Calcium chloride
Magnesium oxide
Calcium chloride
Potassium chloride
Magnesium oxide is having a high melting point than sodium chloride why?
Which one of the following is an ion?
Draw the electronic configuration of chloride ion
Draw the electronic configuration of following
Define covalent bonding
Draw the cross and dot diagram for the following
Hydrogen
Water
Methane
Hydrogen chloride
Nitrogen
Ehane
Carbondixide
iodine
25. How will you convert the hydrogen chloride gas into hydrochloric acid?
……………………………………………..
26. A solution in which the hydrochloric acid is dissolved is tested with blue litmus. What do you observe?
…………………………………………………………………………………
27. Hydro choric acid is dissolved in methyl benzene. What do the change in the blue litmus?
……………………………………………………….
28. Describe the laboratory preparation of chlorine from hydro choric acid ( important)
………………………………………………………………………………………
………………………………………………………………………………..
29. Name the oxidizing agent in this reaction……………………………………
30. Give the test for chlorine…………………………………………………………….
31. Chlorine and bromine have similar properties why?
……………………………………………………………………………….
32. Chlorine is passed through a solution of potassium bromide
Name the type of reaction………………………………………….
The solution became red in color. Why?
…………………………………………………………………………..
Explain why no displacement reaction when we add bromine to potassium chloride …………………………………………………………………………………………
Write the ionic equation for the reaction……………………………………………….
Chlorine is bubbled through the potassium iodide solution. What will be the color of the
Final solution ……………………………………………..
Write down the ionic half equation for the reaction …………………………………
What is the color change during the below reaction
……………………………………………………
H2 + Cl 2___________ 2HCl
Name the process of the industrial extraction of oxygen
………………………………………………..
complete the table
78 %
21%
.9%
What is oxidation?
……………………………………….
What is reduction?
…………………………
Give example for oxidizing agent
…………………………………………….
Give example for reducing agent
………………………………………
Complete these reactions ( balance )
S+ O2__________
Fe + O2__________
Mg + O2 _____________
Cu + O 2_________
C +O 2 ___________
CH4 + O2_____________
Describe an experiment to determine the % of oxygen in air (copper turning )
Give the reaction of sulphur dioxide with water
…………………………………………….
Give two equations for the reactions in contact process ( impor)
…………………………………………………………….
……………………………………………………….
Which one is reversible reaction?
……………………………………………
Describe the laboratory preparation of carbon dioxide
Give the equation for reaction
……………………………………………………………………………….
Give the uses of carbon dioxide
………………………………………………………………
What is dry ice?
……………………………
Give two uses of dry ice
………………………………………..
Give the reaction of nitrogen with oxygen ( 2 equations)
…………………………………………………….
…………………………………………………..
Give the conditions needed for the rusting of iron
……………………………………………………………………………….
Give method for the prevention of rusting
……………………………………………………………………………..
How this method work in preventing rusting of iron
……………………………………………………………………………
Explain sacrificial protection. Give example.
……………………………………………………………………………….
Sodium can be prepared by electrolysis only were as iron fro blast furnace give your answer (in terms of reactivity)
………………………………………………………………………………………
Give the allotropes of sulphur
……………………………………………………………………………
Name the process by which nitrogen is industrially extracted
………………………………………………………………..
Give one use of nitrogen
……………………………………….
Why nitrogen is used for protecting dry foods
……………………………………………….
Describe the laboratory preparation of ammonia (page-244 text book)
Give the equation for the preparation
………………………………………………………………..
Give physical properties of ammonia
………………………………………………………………………………………
ammonia dissolve in water forms …………………………………………
give the equation…………………………………………
Give the uses of ammonia
to prepare nitric acid
to prepare fertilizer
Give the reaction between sulphuric acid and ammonia ( important)
NH 3 + H2SO4 ----------- (NH4)2 SO4 ( AMMONIUM SULPHATE)
Give the reaction between NITRIC acid and ammonia ( important)
NH3+ HNO3 ------- NH 4 NO3( AMMONIUM NITRATE)
What are the elements present in NPK FERTILIZER?
………………………………………………………..
Describe the chemical test for water
……………………………………………………………..
Describe the physical test for water
………………………………………………………………
Describe the laboratory preparation of hydrogen
Name the gas produced when a metal react with dilute acid
…………………………………………………….
Give the test for this gas
…………………………………………………….
Give word equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give balanced equation for the reaction between hydrogen and oxygen
…………………………………………………..
Give one example for transition metal which is used as catalyst
…………………………………………………….
Iron react with chlorine to form a yellow( brown) color solid
Name it…………………………….
Give equation for the reaction
…………………………………………………..
Give the color of ion with sodium hydroxide
Fe2+ …………………
Fe3+…………………
Cu2+ ………………
Name the compounds……………………………………………………..
Write down the action of iron over steam. Give the balanced equation ( notes)
…………………………………………………………………..
Write down the action of iron hydrogen chloride gas Give the balanced equation (notes)
………………………………………………………………………….
Color
Copper oxide-red
Copper hydroxide – blue
Copper nitrate –green
Copper sulfate- blue
Copper carbonate-black
Copper chloride- green
Write down the reaction of ammonia with copper (II) ions ( important)
………………………………………………………………………………
Give the properties of transitional metals( answer given below)
1) variable valency
2) colored compounds
3) catalytic property
Give example for Fe shows variable valency
…………………………………………………………………….
What is reactivity series? Give the order
…………………………………………………………………………………
What do you know about the reactivity of and displacement reaction?
………………………………………………………………………….
A copper dipped in silver nitrate solution .What is observation that you see
After 2hours…………………………………………………….
.
Iron nail dipped in copper sulphate ……………………………………
Silver nail dipped ii copper sulphate……………………………………
Copper nail in iron sulphate……………………………..
(Important portion)
Cat ion
Flame color
Lithium(Li+)
Sodium (Na+)
Potassium( k+)
Calcium( ca2+)
gas
test
result
oxygen
carbon dioxide
hydrogen
chlorine
Sulpher dioxide
ammonia
Test for anions
TEST
result
Chloride(Cl-)
Bromide(Br -
Iodide (I-)
Sulphate(so42+)
Sulphite( S2_)
Carbonate(CO32_)
Test for cat ion
TEST
result
Ammonia (NH4+)
Copper (Cu2+)
Iron( Fe2+)
Iron ( Fe3+)
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